Notes on Chemistry SS2 – Edudelight.com

SCHEME OF WORK FOR SECOND TERM SS2 CHEMISTRY

WEEKS           TOPICS

1.           RATE OF A CHEMICAL REACTION
2.           TYPES OF CHEMICAL REACTION
3.           CHEMICAL EQUILIBRIUM
4.           HYDROGEN
5.           OXYGEN AND ITS COMPOUNDS
6.           THE HALOGENS
7.            NITROGEN AND ITS  PROPERTIES
8.           COMPOUNDS OF NITROGEN
9.           SULPHUR AND ITS PROPERTIES
10.           COMPOUNDS OF SULPHUR

WEEK 1

Chemical Reaction Rates

 
Chemical reaction rate refers to one or more of the following:

• The rate of disappearance of one of the reactants.
   
• The rate of appearance of one of the products.
   
• The rate of change of concentration in a constant volume process.

The units of reaction rate are (moles / dm³) / second. Different chemical reactions take place at different rates. The study of reaction rates is the study of the forming and breaking of chemical bonds.

The nature of the reactants involved in a chemical reaction will determine the kind of bonding that occurs.

• Reactions with bond rearrangement or electron transfer generally take longer than ionic reactions.
   
• Ionic reactions are almost instantaneous because of the strong attraction between the charged particles.

Because of electron cloud repulsion, most neutral molecules that come in contact with one another bounce off without reacting. For these molecules to react, they must collide with enough kinetic energy to cause changes in the electron clouds of both molecules. When such a change occurs, an activated complex is formed, which allows the reaction to proceed. The energy required to give molecules enough kinetic energy to form an activated complex is known as the activation energy for the reaction.

Other factors affecting reaction rate:

• Concentration – For a chemical reaction to occur, the particles must collide. An increase in the number of particles per unit volume (concentration) increases the chance of their colliding. Increasing the concentration of the reacting particles generally increases the reaction rate.
   
• Temperature – An increase in the speed of molecules increases the number of molecules that have the required activation energy. Increasing the number of molecules with activation energy generally increases the reaction rate. Increasing the temperature of the reacting particles generally increases the reaction rate.
   
• Catalysts – A catalyst is a substance that increases a reaction rate without being permanently changed. The catalyst appears to be chemically unaffected by the reaction. A catalyst changes the reaction mechanism in such a way that the activation energy required is less than that in the uncatalyzed reaction.
   
• Inhibitors – An inhibitor ties up a reactant in a complex so that it will not react.

Stable compounds do not spontaneously decompose in air.

• Thermodynamically stable – At room temperature, the compound does not spontaneously decompose in air. The overall energy change in the decomposition reaction is positive.
   
• Kinetically stable – At room temperature, the compound spontaneously decomposes in air. The overall energy change in the decomposition reaction is negative.

However, the reaction takes place so slowly, at room temperature, that no observable change takes place for years. Since the rate of change is imperceptible, the substance is still considered stable.

Unstable compounds spontaneously decompose in air. The reaction takes place at such a rate that the change is observable.

 
A rate expression is a mathematical equation used to calculate the rate of a chemical reaction.

The rate expression for a reaction is expressed as k times the product of the concentrations of the reactants.

r a t e = k [ a ] [ b ]

• The brackets [ ] in the rate expression indicate concentration in moles/dm3.
   
• k in the rate expression is the specific rate constant.
   
• Any reaction has only one value for the constant, k, at a given temperature.

If a chemical reaction occurs in one step, the concentrations of the reactants in the rate expression have exponents equal to the coefficients of the reactants in the balanced equation for the reaction.

• For the reaction: H₂O₂ + 2 HI → 2 H₂O + I₂
   
• The rate expression is: rate = k [H₂O₂] [HI]²

If a chemical reaction occurs in more than one step, the concentrations of the reactants in the rate determining step (the step that has the slowest reaction rate) are the only ones that are important.

• The only way to be sure of the rate expression is to use experimental data.
   
• Experimental data proves the actual rate expression for the reaction above is:

rate = k [H₂O₂] [HI]

This shows that the reaction occurs in more than one step

Reactions where a single species falls apart in some way are slightly simpler because you won’t be involved in worrying about the orientation of collisions. Reactions involving collisions between more than two species are going to be extremely uncommon (see below). Reactions involving collisions between two species It is pretty obvious that if you have a situation involving two species they can only react together if they come into contact with each other. They first have to collide, and then they may react. Why “may react”? It isn’t enough for the two species to collide – they have to collide the right way around, and they have to collide with enough energy for bonds to break. (The chances of all this happening if your reaction needed a collision involving more than 2 particles are remote. All three (or more) particles would have to arrive at exactly the same point in space at the same time, with everything lined up exactly right, and having enough energy to react. That’s not likely to happen very often!) The orientation of collision Consider a simple reaction involving a collision between two molecules – ethene, CH2=CH2, and hydrogen chloride, HCl, for example. These react to give chloroethane.

As a result of the collision between the two molecules, the double bond between the two carbons is converted into a single bond. A hydrogen atom gets attached to one of the carbons and a chlorine atom to the other.

The reaction can only happen if the hydrogen end of the H-Cl bond approaches the carbon-carbon double bond. Any other collision between the two molecules doesn’t work. The two simply bounce off each other.

Of the collisions shown in the diagram, only collision 1 may possibly lead on to a reaction

Note:  The only difference if the reaction was endothermic would be the relative positions of the reactants and products lines. For an endothermic change, the products would have a higher energy than the reactants, and so the green arrow would be pointing upwards. It makes no difference to the discussion about the activation energy.

If the particles collide with less energy than the activation energy, nothing important happens. They bounce apart. You can think of the activation energy as a barrier to the reaction. Only those collisions which have energies equal to or greater than the activation energy result in a reaction.

Any chemical reaction results in the breaking of some bonds (needing energy) and the making of new ones (releasing energy). Obviously some bonds have to be broken before new ones can be made. Activation energy is involved in breaking some of the original bonds.

Where collisions are relatively gentle, there isn’t enough energy available to start the bond-breaking process, and so the particles don’t react.

 

Factors Influencing Rate of a Reaction

The rate of reactions is influenced by the following factors:

(i) Concentration of the reactants,

(ii) Temperature of the reactants,

(iii) Particle size and nature of the reacting substances,

(iv) Presence of catalyst, and

(v) Exposure to radiations.

1. Dependence of Rate on Concentration

When a chemical reaction occurs, the reactants change over to products. It is observed that with the passage of time the concentrations of reactants decrease while those of products increase. It is graphically shown in Fig. 20.8.

Fig. 20.8. Time dependence of the concentrations of reactants and products in a reaction.

Now, if we assume that other factors are constant then the rate of a chemical reaction decreases with the decrease in concentration of the reactants. For example, we find that a piece of’ wood bums at a much faster rate in oxygen than in air. It is because of higher concentration of O₂ in the fanner.

Cato Guldberg and Peter Waage proposed a qualitative relationship between the rates of reactions and the concentration of the reacting species. This generalisation is known as ”Law of Mass Action”. It states that:

At a given temperature, the rate of a chemical reaction is directly proportional to the product of molar concentrations of reacting species with each concentration term raised to the power equal to numerical coefficient of  that species in the chemical equation.

Thus, for a hypothetical reaction;

A + B à Produces

The rate according to law of mass action is given as Rate of reaction

where [A] and [B) are the molar concentrations of the reactants A and B respectively and k is a constant of proportionality.

Particle Size and Nature of Reactants

It is commonly observed that the nature of the reacting substances has a marked influence on the reaction rates. The effect of the nature of reactants can be described in terms of the following factors:

(a) Physical state of reactants

(b) Surface area and particle size of reactants

(c) Chemical nature of reactants.

(a) Physical State of the Reactants

We know that the chemical reaction takes place as a result of the collisions between the reacting particles. In order to collide with one another, the reacting particles must come close and intermix. Now intermixing of reacting particles cannot take place easily if they are in solid state. This is due to restriction in their molecular motion. However, in liquid state as well as in gaseous state the intermixing is possible and chances of their collision become larger. Thus, the reaction will be faster if the reactants are mixed in liquid phase or solution phase.

This can be illustrated with the help of following reaction between lead nitrate and potassium iodide.

Pb (NO₃)₂ + KI à PbI₂ + KNO₃

(i) In one of the experiment, solid Pb(NO₃)₂ is mixed with solid KI in a china dish with the help of spatula.

(ii) In another experiment, aqueous solution of Pb(NO3)2 is mixed with aqueous solution of KI in a beaker.

It is observed that rate of formation of ( yellow coloured Pbl₂ is much faster in the second experiment because probability of encoumers between the- reacting particles is much larger in liquid phase than that in solid phase.

(b) Surface Area and Particle Size of Reactants

Particle size and surface area play an important role in rates of heterogeneous reactions particularly when one of reacting species is in .solid phase. Smaller the size of the reacting particles present in the solid phase larger will be their surface area. As a result, more will be the probability of the reacting species coming in contact with each other and therefore, more will be the .number of encounters between them. For example,

• combustion of a piece of coal in air is relatively slower, whereas combustion of coal dust in air is much faster. dissolution of sugar crystals in water is slower but
• dissolution of powdered sugar is rapid. Thus, we find that the rate of the process is increased, if the particle size in the solid state is decreased. In fact, decrease in the particle size increases the surface area.

(c) Chemical Nature of the Reactants

The chemical nature of reacting substances also affects the reaction rates significantly. Depending upon the nature of reacting species, the rates of different reactions may differ very widely from one another. For example, the oxidation of ferrous ions (Fe+2) by potassium permanganate in acidic medium is practically instantaneous. Whereas oxidation of oxalate ions (C₂O ₄^{– 2}) by potassium permanganate in acidic solution is comparatively much slower:

In these two reactions, everything is identical except the nature of reducing agents.

3. The Effect of Pressure

The change in pressure also affects the rate of a reaction.

• If one or more .of the reactant§ is a gas then increasing pressure will effectively increase the concentration of the reactant molecules and speed up the reaction.
• Changing the pressure i.e., lower or higher, results in lesser or greater concentration and so slower or faster reaction, all because of the increased chance of a ‘fruitful’ collision.

4. The Effect of Stirring

In doing rate experiments with a solid and solution reactant e.g., marble chips-acid solution or a solid catalyst like manganese(IV) oxide catalysing the decomposition of hydrogen peroxide solution, it is sometimes forgotten that stirring the mixture is ail important rate factor.

• If the reacting mixture is not stirred ‘evenly’, the reactant concentration in solution becomes much less near the solid, which tends to settle out at the bottom of the flask.
• Therefore, at the bottom of the flask the reaction prematurely slows down distorting the overall rate measurement and making the results uneven and therefore inaccurate. The ‘unevenness’ of the results is even more evident by giving the reaction mixture the ‘odd stir’ ! You get jumps in the graph!! !

5. The Effect of Radiations

Reactions carried out in the presence of light radiations are known as photochemical reactions. The presence/absence of radiations greatly affects the rate of a reaction. Some examples are given below to illustrate the idea.

• Halogenation of methane can readily occur in the presence of uv radiations. The reaction does not occur in dark. The radiations supply sufficient energy to break the bonds of Cl2 molecule giving free radicals and hence the reaction.

Photosynthesis in green plants occurs in the presence of sunlight. The green chlorophyll molecules absorb the photon energy packets of sunlight and initiate the chemical changes summarised below:

Water  + carbon dioxide  à glucose + oxygen

6H₂O(l) + 6CO₂(g)  à C6H1₂O₆(aq) + 6O₂(g)

6. Effect of Temperature

The rates of almost all reactions increase with the increase in temperature. For example,

• In decomposition of N₂O ₅, the time taken for half of original amount of material to decompose is 12 min at 50°C, 5 hours at 25°C and 10 days at 0°C. ·

• In a mixture of KMnO ₄ and Oxalic acid (H2Cp ₄), potassium permanganate gets decolourised faster at higher temperature that at lower temperature.

In most of the cases it has been found that the rate of the reaction becomes almost double for every 10° rise of temperature. This is also expressed in the terms of Temperature coefficient which is the ratio of rate constants of the reaction at two temperatures differing by 10°. The two temperatures generally selected are 298 K and 308 K.

Thus,

Temperature coefficient = Rate constant of 308 K / Rate constant at 298 K

The rate constants for the decomposition of N₂O₅ at different temperatures is given in Table 20.1. From the table, it is seen that rate constant at 273 K is 7.87 x w^-1 and at 298 K is 3.56 X w^-5 This shows that for 25° rise of temperature of the rate constant increases by 43 times.

Table 20.1. Rate Constants for the Decomposition of N₂O ₅ at Different Temperatures

Experiment to Show How Changes in Temperature Affect the Rate of a Reaction

Home → Experiment to Show How Changes in Temperature Affect the Rate of a Reaction

Kids, do you think that temperature will have an affect on how fast or slow a reaction might take place? In order for a chemical reaction to occur, the molecules, which are reactants, must physically come into contact with one another. Anything that increases the frequency of these encounters will increases the rate at which products are formed. Your hypothesis can be that the rate of a chemical reaction will be increased by raising the temperature of reactants.

Requirement

Six clear cups, a measuring cup, a thermometer (-20°C to 110°C), three original formula effervescent Alka-Seltzer® tablets, a stopwatch, a mortar and pestle, a source of hot water, ice cubes, and graph paper.

Procedure

Fill a clear glass with exactly 50 cm³ of hot water. Use the thermometer to take the temperature and record it on a data sheet. Drop 1 Alka-Seltzer® tablet into water. Measure the time required for the tablet to fully dissolve. Be prepared to start and stop on time. The reaction will take less than 15 seconds. Record the time. Repeat this experiment using room temperature water.

For the cold water test, the procedure is a little different. Fill a clear glass with 25 crn³ of water and add enough ice to adjust the level to 50 cm³. Stir the ice water for about 15 seconds so the temperature will come to equilibrium. Use the thermometer to take the temperature and record it on your data sheet. (Leave the ice cubes in the water!) Drop 1 Alka Seltzer tablet into the water. Measure and record the time required for the reaction to be completed.

Analysis

Graph your data points (water temperature vs. time to fully dissolve) to show the effect of temperature on rate of reaction.

Analysis and Interpretation of Simple Graphs on Rates of Reaction

Home → Analysis and Interpretation of Simple Graphs on Rates of Reaction

Graph 1 shows the decrease in the amount of a solid reactant with time. The graph is curved, becoming less steep because the reactants are being used up, so the speed decreases. In the first few minutes the graph will (i) decline less steeply for larger ‘lumps’ and (ii) decline more steeply with a fine powder o time (min) i.e., (i) less surface area gives slower reaction and (ii) more surface area a faster reaction.

Graph 2 shows the increase in the amount of a solid product with time. The graph tends towards a maximum amount possible when all the solid reactant is used up and the graph becomes horizontal. This means the speed has become zero as the reaction has stopped.

Graph 3 shows the decrease in reaction time with increase in temperature as the reaction speeds up. The reaction time can represent how long it takes to form a fixed amount of gas in the first few minutes of a metal carbonate-acid reaction, or the time it takes for so much sulphur to form In tin sodium thiosulphate-hydrochloric acid reaction.

Graph 4 shows the increase in speed of a reaction with increase in temperature as the particles have more and more kinetic energy. The rate of reaction is proportional to lit, where tis the reaction time.

Graph 5 shows the increase in the amount of a gas formed in a reaction with time. Again, the graph becomes horizontal as the reaction stops when one of the reactants is all used up!

Graph 6 shows the effect of Increasing concentration, which decreases the reaction time, as the speed increases because the greater the concentration the greater the chance of fruitful collision.

Graph 7 shows the rate/speed of reaction is often proportional to the concentration of one particular reactant. Rate This is due to the chance of a fruitful collision forming products being proportional to the concentration.

Collision Theory of Reaction Rates

Home → Collision Theory of Reaction Rates

Collision theory, was developed by Max Trautz and William Lewis in 1916-18 and, provides a greater insight into the energetic and mechanistic aspects of reactions and their rates. According to this theory, the reactant molecules are assumed to be hard spheres and reaction can occur when molecules collide with each other. The number of collisions per second per unit volume of the reaction mixture is known as collision frequency (Z). The more frequent are collisions, the faster is the reaction.

POSTULATES OF COLLISION THEORY

(i) Reactions occur due to approach and collisions of reactant particles (atoms, molecules or ions) .

(ii) A successful collisions, the one required for changing reactants to products, can occur between two molecules only if they possess a certain minimum amount of energy in excess of the normal energy of molecules.

(iii) Only a small fraction of collisions are successful in producing a reaction. These collisions are called effective collisions.

(iv) The minimum energy which molecules must possess before collision should be equal to or greater than the activation energy.

(v) The rate of reaction is proportional to the frequency of effective collisions per second.

It must be noted here that the collisions between reactant molecules will not lead to reaction even if the energy requirement is satisfied. It is because the colliding molecules should also have proper orientation.

Proper Orientation:

Improper Orientation:

The proper orientation of reactants molecules leads to bond formation whereas improper orientation make them simply bounce back and no products are formed.

A pictorial representation (Fig. 20.11(n) .and 20.11 (b)) is shown, to illustrate the above idea.

An extension of the collision theory is the activated complex or transition state theory. In the activated. complex theory:

1. When two molecules or species approach each other, their electron clouds repel and distort each other. The potential energy of the reactants rises and reaches a maximum value, forming a combined and highly unstable molecule called the activated complex.

2. When the activated complex loses energy it can either form products or break down to form the original reactants. The activated complex is described as being in the transition state.

3. The energy required to from the activated complex is known as the activated energy barrier, and it is only reacting molecules which have greater energy than the activation energy barrier which ‘climb’ over to form the products (Fig. 20.11).

Collision Theory and Reaction Rate

Home → Collision Theory and Reaction Rate

Reactions can happen only when the reactant particles collide. All the collisions which occur are not successful in forming product molecules. The reason is that particles have a wide range of kinetic energy but only a small fraction of particles have enough kinetic energy to break bonds and bring about chemical change. The minimum kinetic energy required for reaction is known as the activation energy.

The small number of high kinetic energy collisions between particles which do produce a chemical change are called ‘fruitful _collisions’. Here the reactant molecules collide with enough kinetic energy to break the original bonds and form new bonds in the product molecules.

It must be noted that the rate-controlling factors described below are to do with the collision frequency (chance of collision) or the energy of reactant particle collision (energy activation energy) which can summed up as the ‘chance of a fruitful collision’ leading to product formation.

THE EFFECT OF CONCENTRATION

If the concentration of any reactant in a solution is increase, the rate of reaction is increase. Increasing the concentration, increases the probability of a collision between reactant particles because there are more of them in the same volume and so increases the chance of a fruitful collision forming products.

For example, Increasing the concentration of acid molecules in a reaction of zinc with HCl, increase the frequency or chance at which they hit the surface of marble chips to dissolve them faster.

THE EFFECT OF TEMPERATURE

When gases or liquids are heated the particles gain kinetic energy and move faster (see diagrams below.) The increased speed increases the chance (frequency) if collision between reactant molecules and the rate increases.

It must be noted that this is not the main reason for the increased reaction speed, so be careful in your theory explanations if investigating the effect of temperature.

THE EFFECT OF SURFACE AREA-PARTICLE SIZE OF A SOLID REACTANT

If a solid reactant or a solid catalyst is broken down into smaller pieces the rate of reaction increases.

The speed increase happens because smaller pieces of the same mass of solid have a greater surface area compared to larger pieces of the solid.

Therefore, there is more chance that a reactant particle will hit the solid surface and react.

Role of Activation Energy in Chemical Reactions

Home → Role of Activation Energy in Chemical Reactions

A reaction can occur when molecules of reactants collide with each other to form an unstable intermediate (Fig. 20.12). The intermediate exists for a very short time and then breaks up to form product molecules. The energy required to form this intermediate, called activated complex (C), is known as activation energy (E _a). The distribution of kinetic energy of different fraction (N_E/N_T) of molecules is shown in Fig. 20.12.

Fig. 20.12. Distribution curve showing energies among gaseous molecules.

Only those collisions result in the formation of products which possess energy equal to or more than the certain minimum energy called threshold energy. Collisions of the molecules possessing energy less than threshold energy do not form products. It means that between reactants and products there is an energy barrier which must be crossed before the reactants are converted into products. The energy required for crossing this energy barrier is supplied by the kinetic energy of the -molecules.

“The minimum extra energy over and above the average potential energy of the reactants which must be supplied to the reactants to enable them to cross over the energy barrier between reactants and products is called Activation energy”. Thus,

Activation energy= (Threshold energy) – (Average energy of the reactants)

or                                 E _a = E _y-E _R

The idea of activation energy and the energy barrier involved in a reaction is given in the Fig. 20.13.

Fig. 20.13. Illustration of activation energy and energy barrier involved in a reaction.

It is important to note that each reaction has a definite value of Ea and this decides the fraction of total collisions which are effective. Obviously, if the activation energy for a reaction is low, large number of molecules can have this energy and the fraction of effective collision, J, will be large. Such a reaction proceeds at high rate. On the other hand, if the activation energy is high, then f will be small and the reaction may be quite slow.

In short,

For fast reactions; activation energies are low.

For slow reactions; activation energies are high.

For example, for the reaction between NO and O₂, Ea is low and hence, reaction is fast whereas for the reaction between CO and O₂, Ea is high and hence, the reaction is slow.

2NO + O₂ à 2NO₂

2CO + O₂ à 2CO₂

The two theories can be successfully used to explain how the factors of surface area, concentration or pressure of gases, temperature and catalysts affect reaction rates.

In the case of surface area, concentration and pressure of gases at constant temperature, any increase in any individual value is equivalent to an increase in the number of molecules colliding per second. An increase in the frequency of effective collisions results in an increase in reaction rate.

ACTIVATED COMPLEX OR TRANSITION STATE THEORY

It has been pointed out earlier that during the chemical reaction certain bonds are broken and certain new bonds are formed. The breaking of bonds requires energy whereas the formation of bonds results in the release of energy. For example, in the reaction of hydrogen with iodine to form hydrogen iodine, when a molecule of hydrogen approaches that of iodine, H-H and I- I bonds start breaking and H- I bonds start forming. In the beginning, breaking of bonds predominates and therefore, energy of the system starts increasing till it reaches a maxima (corresponding to threshold energy). After this, the energy starts decreasing because the process of bond· formation predominates and finally leads to the product hydrogen iodide. The arrangement of atoms corresponding to energy 11UlXima (threshold energy) is called transition state or activated complex. In transition state, the system has partial reactant character and partial product character as shown in Fig. 20.14 and Fig. 20.15.

Fig. 20.14. Formation of activated complex during the reaction of H2 and I₂ from HI.

Fig. 20.15. Transition state or activated complex.

The difference between energy of the transition state and energy of the reactants is equal to activation energy.

E _{transition state} – E _reactants = E _activation

Effect of Temperature and Catalyst on the Rate Constant

Home → Effect of Temperature and Catalyst on the Rate Constant

Let us now consider the effect of increase in temperature on the number of effective collisions.

Fig. 20.16. Energy distribution at different temperatures

On the basis of probability consideration Fig. 20.16 is drawn to give the energy distribution curves at temperatures T1 and T2 (where T 2 = T 1 + 1 0). Now as we know that the rise in temperature increases the kinetic energy of molecules ( ·: K.E. T) therefore, the energy distribution curve gets flattened and shifts towards higher energy region. A close examination of the curves in the graph clearly reveals that the function of molecules possessing higher kinetic energy, i.e., energy greater than threshold energy, as indicated by shaded portion becomes almost double and therefore the rate of reaction almost doubles for 10 rise of temperature. Thus, increase in the rate of reaction with increase in temperature is mainly due to increase in number of collisions which are energetically effective.

Temperature Dependence of Rate Constant

Arrhenius developed the mathematical relation between temperature and the rate constant on the basis of the observations from the large number of experiments. This temperature dependence of the rate constants is expressed algebraically as

K = A _e^{-E a /RT}

Equation (20.1) is called Arrhenius equation. Here A is pre exponential factor and is called frequency factor, Ea is the energy of activation and Tis the temperature in kelvin scale.

The term e -Ea I RT in the above equation is also called as Boltzmann Factor. Both A and E0 are characteristic of the reaction. Another form of the equation which is more useful for calculations is obtained by taking logarithm of Eqn. (20.1), therefore,

Log k₂ / k₁ = Ea / 2.303 R [ 1/T₁ – 1 / T₂]

Here k₂ and k₁ are rate constant at temperature T₁ and T₂ respectively E_a is activation energy and R is universal gas constant.

EFFECT OF CATALYST ON RATE OF REACTION

A catalyst is a substance that alters the rate of chemical reaction without itself being permanently chemically changed. Never state things like “it doesn’t react, just speeds it up”. It must take part in the reaction and it must change chemically, albeit on a temporary basis. A catalyst provides a different ‘pathway’  ormechanism that makes the bond breaking processes (or other electronic changes in the reactants) occur more readily. In general,

• A catalyst speeds up a reaction, but it must be involved ‘chemically’, however temporarity, in some way,_ and is continually changed and reformed as the reaction proceeds.

• Catalysts work by providing an alternative reaction pathway of lower activation energy.

Thus, the function of a catalyst is thus to lower down the activation energy. In sample words, greater the decrease in the activation energy caused by the catalyst, higher will be the reaction rate. In the presence of a catalyst, the reaction follows a path of lower activation energy. Under this condition, a large number of reacting molecules are able to cross-over the energy barrier and thus the rate of reaction increases. The energy profile diagram for the catalysed and uncatalysed reactions are as shown in the Fig. 20.17. Where dotted curve represents the progress of uncatalysed reaction and solid curve represents:- the catalysed reaction.

Fig. 20.17. Potential energy curves for catalysed and uncatalysed reactions.

For a general reaction of the type A + B à AB the course of uncatalysed and catalysed reaction may be represented as:

(a) Uncatalysed reaction:

(b) Catalysed reaction:

Though the catalyst increases the rate of the reaction, yet it does not effect the state of equilibrium in case of reversible reactions. It is because the activation energy for the forward reaction and backward reaction is reduced to the same extent.

EVALUATION

1. Rate of a chemical reaction depends on the following factors except

a. rate at which gas is evolved  b. rate at which products are formed  c. rate at which colour of reaction changes  d. rate at which reactants diminish.

2.If 2g  of zinc granules was reacted with excess dilute HCl to evolve hydrogen gas which came to completion after 5minutes.Calculate the rate of the chemical reaction in g|hr.

3.Which of the following  lowers the  activation energy of a chemical reaction

a.water  b .freezing mixture  c.catalyst d.reducing agent.

4.Which of  the following will not affect a chemical system at equilibrium?

a.Temperature  b. catalyst   c. pressure   d.concentration

5. An increase in the rate of a chemical reaction on heating is due  to increase in a.activation energy  b.enthalpy c.collision rate    d .reactant concentration 6.A  catalyst a.increase the amount of product in an equilibrium mixture b.only affects gaseous reaction c.must be used in small  quantities THEORY 1.State all the factors that  affect the rate of a reaction 2.Explain how any three of these factors affect the rate of a chemical reaction 3.State the collision Theory. 4.Define Activation Energy. WEEK  2 Heat changes EXOTHERMIC and ENDOTHERMIC INTRODUCTION Energy is conserved in chemical reactions. One way of stating the ‘law of Conservation of Energy’ is to say the amount of energy in the universe at the end of a chemical reaction is the same as before the reaction took place. If a reaction transfers energy to the surroundings the product molecules must have less energy than the reactants, by the amount transferred. Conversely, if a reaction absorbs energy from the surroundings, they must have less energy, and the products must have more energy. An exothermic chemical reaction transfers energy to the surroundings, usually given out in the form of heat energy, so raising the temperature of the surroundings. Therefore the products have less energy than the reactants and the surroundings have more energy.

Exothermic reactions include combustion of fuels, many oxidation reactions, acid-alkali neutralisation reactions, reactive metals with water,

moderately reactive metals with strong acids.

Exothermic reactions are used in self-heating cans and hand warmers.

An endothermic chemical reaction absorbs energy from the surrounding, usually in the form of heat energy, so cooling the surroundings, but sometimes the system is heated to provide the heat energy and a high enough temperature to promote the reaction. This means the products have more energy than the reactants and the surroundings have less energy.

Endothermic reactions include thermal decomposition of compounds e.g. carbonates, the reaction between citric acid and sodium hydrogencarbonate, sports injury packs to produce cooling effects

Why is it important to know about energy changes in chemical reactions?

Its important to know how much energy fuels release on combustion i.e. their calorific value.

Its important to know the energy released on burning petrol. diesel, coal or any other fossil fuel and alternative fuels like hydrogen or biofuels (biomass fuels).

The same sort of data is important in knowing how much energy is released on metabolising foods such as fats and carbohydrates.

Accurate energy change data is important in managing chemical processes in industry.

Exothermic reactions may provide their own heat if the process is carried out at high temperatures, energy transfer data provides some of the information needed.

Conversely, excess heat from an exothermic reaction may have to be removed using heat exchangers to avoid ‘overheating’ and excessive reaction rates that could be dangerous. If gases are involved, lack of control could lead to a build up of pressure resulting in an explosion.

Endothermic chemical processes often need a high temperature to promote

the absorption of heat energy, otherwise the reaction rate might economically far too slow. The amount of energy needed can be calculated from energy transfer data.

Heat Changes in Chemical Reactions

When chemical reactions occur, as well as the formation of the products – the chemical change, there is also a heat energy change which can often be detected as a temperature change.

This means the products have a different energy content than the original reactants (see the reaction profile diagrams below).

If the products contain less energy than the reactants, heat is released or given out to the surroundings and the change is called an exothermic reaction (exothermic energy transfer, exothermic energy change of the system).

This is illustrated by the simple energy level diagram above for an exothermic reaction.

The products have less energy than the original reactants (lower energy level) and the difference comes out as heat energy released to the surroundings.

The difference in heights of the energy levels tells you how much energy is released in an exothermic reaction.

The temperature of the system will be observed to rise in an exothermic change.

transfer,  So an exothermic reaction is one which gives out energy to the surroundings, usually in the form of heat energy, hence the rise in temperature.

Examples of exothermic reactions:

The burning or combustion of hydrocarbon fuels e.g. petrol or candle wax, these are very exothermic reactions.

1.The exothermic burning-combustion of fossil fuels is very important source of energy.

2.methane (natural gas) + oxygen ==> carbon dioxide + water (+ heat energy)

CH₄ + 2O₂ ==> CO₂ + 2H₂O

3.The burning of magnesium, reaction of magnesium with acids, or the reaction of sodium with water

2Mg + O₂ ==> 2MgO (+ heat energy)

4.The neutralisation of acids with alkalis

NaOH + HCl ==>NaCl + H₂O (+ heat energy)

Using hydrogen as a fuel in hydrogen-oxygen fuel cells xplosions are caused by VERY fast exothermic reactions producing very fast large expanding volumes of gases.

Other uses of exothermic reactions:

Hand warmers contain chemicals that when mixed together give out heat.

Self-heating cans of coffee, soup or hot chocolate have chemicals contained in the base of the container that when mixed generate enough energy to heat the contents of the can.

If the products contain more energy than the reactants, heat is taken in or absorbed from the surroundings and the change is called an endothermic reaction (endothermic energy endothermic energy change of the system).

This is illustrated by the simple energy level diagram above for an endothermic reaction.

The products have more energy than the original reactants (higher energy level) and the difference comes out as heat energy absorbed from the surroundings.

The difference in heights of the energy levels tells you how much energy is absorbed in an endothermic reaction.

If the change can take place spontaneously, the temperature of the reacting system will fall but, as is more likely, the reactants must be heated to speed up the reaction and provide the absorbed heat.

So an endothermic reaction is one which absorbs energy from the surroundings, usually in the form of heat, hence the observed fall in temperature in some reaction.

Examples of endothermic reactions

The thermal decomposition of limestone

calcium carbonate (limestone) ==> calcium oxide (lime) + carbon dioxide

CaCO₃ (+ heat energy) ==>CaO + CO₂

This only happens at temperatures above 900^oC.

the cracking of oil fractions

e.g. octane (+ heat energy) ==> hexane + ethene

C₈H₁₈ ==> C₆H₁₄ + C₂H₄

Again this needs a very high temperature AND a catalyst too.

These are two very important endothermic reactions used in the chemical industry.

Dissolving ammonium nitrate in water doesn’t need heating, the salt spontaneously dissolves and the temperature of the water/solution immediately falls as energy from the surroundings is absorbed, in fact from the water itself.

Other uses of endothermic reactions:

Some sports injury packs have a mixture of chemicals (or maybe a salt and water) that when mixed undergo an endothermic energy change, thereby absorbing heat from the surroundings, cooling some poor bruised limb! Rather more convenient and less messy than packs of ice!

One of the most important endothermic reactions, for which most of animal life depends is photosynthesis.

The energy from sunlight is absorbed as water and carbon dioxide are converted to glucose and oxygen.

6H₂O + 6CO₂ + sunlight energy ==> C₆H₁₂O₆ + 6O₂

However, on ‘burning’ the glucose/carbohydrates in our bodies, the ‘stored’ sunlight energy is released to keep us warm and drive all the chemical processes in our cells, so the opposite reaction is exothermic!

C₆H₁₂O₆ + 6O₂  ==>  6H₂O_(l) + 6CO₂  + heat/chemical energy

The difference between the energy levels of the reactants and products gives the overall energy change for the reaction

At a more advanced level the heat change is called the enthalpy change is denoted by delta H, ΔH.

ΔH is negative (-ve) for exothermic reactions i.e. heat energy is given out and lost from the system to the surroundings which warm up.

ΔH is positive (+ve) for endothermic reactions i.e. heat energy is gained by the system and taken in from the surroundings which cool down OR, as is more likely, the system is heated to provide the energy needed to effect the change.

Extra NOTE: In some exothermic changes, no heat is not released e.g. in batteries and fuel cells, where the energy is released as electrical energy.

Reversible Reactions and energy changes

If the direction of a reversible reaction is changed, the energy change is also reversed.

For a reversible reaction, the energy released in the exothermic reaction is numerically equal to the heat absorbed in the reverse reaction.

For example: the thermal decomposition of hydrated copper(II) sulphate is a very good example to observe in the school laboratory, even though it is not practical to measure the actual energy changes involved.

blue hydrated copper(II) sulphate + heat    white anhydrous copper(II) sulphate + water

CuSO₄.5H₂O_(s)     CuSO_4(s) + 5H₂O_(g)

On heating the blue solid, hydrated copper(II) sulphate, steam is given off and the white solid of anhydrous copper(II) sulphate is formed and left as the residue.

This is a thermal decomposition and is endothermic as heat is absorbed (taken in) from the surroundings.

The energy is needed to break down the crystal structure and drive off the water.

When the white solid is cooled and a few drops of water added, blue hydrated copper(II) sulphate is reformed and heat energy is given out to the surroundings, the mixture hots up!

The reverse reaction is exothermic as heat is given out.

i.e. on adding water to white anhydrous copper(II) sulphate the mixture heats up as the blue crystals reform.

Water molecules recombine with the copper ion releasing energy when the new bonds are formed.

Types of Chemical Reactions

Chemical reactions are processes in which substances change into other substances.

You know a chemical reaction takes place if one or more of these occur:
 

• Color changes – Different combinations of molecules reflect light differently. A color change indicates a change in molecules.
   
• Heat content changes – In all chemical reactions, the heat content of the reactants and the heat content of the products is never the

same. Sometimes the difference is great and can be easily detected. At other times, the difference is slight and more difficult to detect.
 

• Gas produced – Whenever a gaseous product forms in a liquid solution, bubbles can be seen. A colorless gas produced in a

reaction of solids is much harder to detect.
 

• Precipitate forms – Precipitates are insoluble products formed by a reaction taking place in a liquid solution. This insoluble product will eventually settle to the bottom, but might immediately appear by turning the clear solution cloudy.

 
Most chemical reactions can be placed into one of five basic types:
 
1. Decomposition Reactions

• A compound breaks into parts.
   
• compound → element + element
   
• 2H₂O → 2H₂ + O₂

Some decomposition complications with heat:

• Some acids, when heated, decompose into an acidic oxide and H₂O.

H₂SO₃ → SO₂ + H₂O

• Metallic hydroxides, when heated, decompose into a metallic oxide and H₂O.

Ca(OH)₂ → CaO + H₂O

• Metallic carbonates, when heated, decompose into a metallic oxide and CO₂.

Li₂CO₃ → Li₂O + CO₂

  Metallic chlorates, when heated, decompose into metallic chlorides and O₂.

2KClO3 → 2KCl + 3O2   2. Synthesis Reactions Elements are joined together.   element + element → compound   2H2 + O2 → 2H2O   Compounds are joined together   compound + compound → compound   6CO2 + 6H2O → C6H12O6 + 6O2   3. Single Displacement Reactions A single element replaces an element in a compound.   element + compound → element + compound   Zn + 2HCl → H2 + ZnCl2   4. Double Displacement Reactions An element from each of two compounds switch places.   compound + compound → compound + compound   H2SO4 + 2NaOH → Na2SO4 + 2H2O   5. Combustion Reactions A hydrocarbon (a compound containing only carbon and hydrogen) combines with oxygen.

• The products of combustion are always carbon dioxide and water.
  hydrocarbon + oxygen → carbon dioxide + water
   
• CH₄ + 2O₂ → CO₂ + 2H₂O
   
• When metallic substances combine with oxygen, the result is an oxidation-reduction reaction.

The rusting of iron – 4Fe + 3O₂ → 2Fe₂O₃

 
Chemical reactions can be classified in other ways as well:

Neutralization Reactions

• Special types of double displacement reactions that involve the reaction between an acid and base to form a salt and water.
   
• acid + base → salt + water
   
• Heat is usually given off in neutralization reactions.
   
• A suspension of solid magnesium hydroxide in water is widely used as an antacid to neutralize excess stomach acid:
   

Mg(OH)_{2 (s)} + 2HCl _(aq) → MgCl_{2 (aq)} + 2H₂O _(l)

Oxidation-Reduction Reactions

• Any reaction in which elements experience a change in oxidation number.
   
• one atom gains e^&minus and another atom loosese^&minus
   

S + O₂ → SO₂

• In the reaction above, sulfur and oxygen both have an oxidation number of zero before the reaction. After the reaction, sulfur is +4 and oxygen is −2.

Precipitation Reactions

• Aqueous reactions that involve the formation of a precipitate (solid).
   
• soluble compound + soluble compound → insoluble compound
   

2KI _(aq) + Pb(NO₃)_{2 (aq)} → 2KNO_{3 (aq)} + PbI_{2 (s)}

The physical state symbol (aq) says the reaction is taking place in a water solution. The physical state symbol (s) says the lead (II) iodide is a solid – therefore insolubl

                                     FIRST AND SECOND LAW OF THERMODYNAMICS         .

Thermodynamics is a branch of physics which deals with the energy and work of a system. Thermodynamics deals only with the large scale response of a system which we can observe and measure in experiments. In aerodynamics, the thermodynamics of a gas obviously plays an important role in the analysis of propulsion systems but also in the understanding of high speed flows. The first law of thermodynamics defines the relationship between the various forms of energy present in a system (kinetic and potential), the work which the system performs and the transfer of heat. The first law states that energy is conserved in all thermodynamic processes.

The First Law of Thermodynamics – Energy is conserved.

The internal energy of a system is the sum of all the kinetic and potential energies of all its components.

ΔE = E_final − E_initial

The equation above indicates the change in internal energy, ΔE, is the difference between the final energy, E_final, and the initial energy, E_initial.

Thermodynamic quantities like ΔE have three parts: a number, a unit, and a sign.

• The number and unit give the magnitude of the charge.
   
• The sign gives the direction.

• A positive ΔE, when E_final>E_initial, indicates the system gained energy from its surroundsings. (diagram B below)
   
• A negative ΔE, when E_final<E_initial, indicates the system lost energy to its surroundsings. (diagram A below)

In a chemical reaction the initial state of the system refers to the reactants and the final state refers to the products.

2 H_2(g) + O_2(g) → 2 H₂O_(g)

In the reaction above, the system loses energy to the surroundings as heat. Because heat is lost from the system, the internal energy of the products (final state) is less than that of the reactants (initial state), and ΔE is negative.

Another way of saying this – as heat is added to or removed from a system, work is done on or by the system.

Using q to represent the heat added to or removed from the system, and w to represent work, the First Law of Thermodynamics can be represented by the equation:

ΔE = q + w

For q    

• +  means system gains heat    
   
• −  means system loses heat

For w    

• +  means work is done on system    
   
• −  means work is done by system

For ΔE    

• +  means net gain of energy by system    
   
• −  means net loss of energy by system

When a process (such as a chemical reaction) occurs in which the system absorbs heat, the process is called endothermic.

A process in which the system loses heat is called exothermic.

The work involved in the expansion or compression of gases is called pressure-volume work, or P-V work.

When the pressure is constant, the sign and magnitude of the P-V work is given by

      w = − P ΔV

where P is pressure and ΔV is the change in volume of the system

      (ΔV = V_final − V_initial).

Enthalpy, H is the thermodynamic function that accounts for heat flow in processes occurring at constant pressure when no forms of work are performed other than P-V work.

H = E + PV

At constant pressure, a change in enthalpy equals the change in internal energy plus the product of the constant pressure times the change in volume.

ΔH = ΔE + P ΔV When ΔH is positive, the system has gained heat from the surroundings – endothermic.   When ΔH is negative, the system has lost heat to the surroundings – exothermic. Enthalpy of reaction, the enthalpy change for a chemical reaction is expressed by the equation: ΔH = Hproducts − Hreactants Here is an example of a thermochemical equation:   2 H2(g) + O2(g) → 2 H2O(g)       ΔH = − 483.6 kJ       This equation indicates two moles of hydrogen gas burn to form two moles of water at a constant pressure, releasing 483.6 kJ of heat. Enthalpy change during a chemical reaction can also be represented in an enthalpy diagram, showing the reactants at the top and the products at the bottom. Guidelines for using thermochemical equations and enthalpy diagrams: The magnitude of ΔH is directly proportional to the amount of reactant consumed in the process.   The enthalpy change for a reaction is equal in magnitude, but opposite in sign, to ΔH for the reverse reaction.   The enthalpy change for a reaction depends on the state of the reactants and products.

We can imagine thermodynamic processes which conserve energy but which never occur in nature. For example, if we bring a hot object into contact with a cold object, we observe that the hot object cools down and the cold object heats up until an equilibrium is reached. The transfer of heat goes from the hot object to the cold object. We can imagine a system, however, in which the heat is instead transferred from the cold object to the hot object, and such a system does not violate the first law of thermodynamics. The cold object gets colder and the hot object gets hotter, but energy is conserved. Obviously we don’t encounter such a system in nature and to explain this and similar observations, thermodynamicists proposed a second law of thermodynamics. Clasius, Kelvin, and Carnot proposed various forms of the second law to describe the particular physics problem that each was studying. The description of the second law stated on this slide was taken from Halliday and Resnick’s textbook, “Physics”. It begins with the definition of a new state variable

calledentropy. Entropy has a variety of physical interpretations, including the statistical disorder of the system, but for our purposes, let us consider entropy to be just another property of the system, like enthalpy or temperature.

The second law states that there exists a useful state variable called entropy S. The change in entropy delta S is equal to the heat transfer delta Q divided by the temperature T.

• delta S = delta Q / T

For a given physical process, the combined entropy of the system and the environment remains a constant if the process can be reversed. If we denote the initial and final states of the system by “i” and “f”:

Sf = Si (reversible process)

An example of a reversible process is ideally forcing a flow through a constricted pipe. Ideal means no boundary layer losses. As the flow moves through the constriction, the pressure, temperature and velocity change, but these variables return to their original values downstream of the constriction. The state of the gas returns to its original conditions and the change of entropy of the system is zero. Engineers call such a process an isentropic process. Isentropic means constant entropy.

The second law states that if the physical process is irreversible, the combined entropy of the system and the environment must increase. The final entropy must be greater than the initial entropy for an irreversible process:

Sf > Si (irreversible process)

An example of an irreversible process is the problem discussed in the second paragraph. A hot object is put in contact with a cold object. Eventually, they both achieve the same equilibrium temperature. If we then separate the objects they remain at the equilibrium temperature and do not naturally return to their original temperatures. The process of bringing them to the same temperature is irreversible.

EVALUATION

1. Heat content difference between the products and reactants of a chemical reaction is called

a. entropy change  b. endothermic change   c. exothermic change  d. enthalpy change.

2. Substances which dissolves endothermically would dissolve more at   a. lower temperature  b. any temperature provided not below the freezing point.  c. higher temperature  d. any isolated temperature.

3. From,change in G=change in H-T x change in S, these will make change in G negative except

a. decreasing the temperature  b. decreasing the enthalpy change  c. increasing the temperature  d. making change in H AND T X change in S the same.

4.When change in H is negative, the reaction is said to be a. exothermic  b. endothermic 

c. ionic  d. reversible  d. catalytic

section b

1.Draw the energy diagrams for the catalysed and uncatalysed

i Exothermic reaction

ii Endothermic reaction

          2.  Give 5 examples each of Exothermic and Endothermic       Reactions.

Notes on Chemistry SS2 – Edudelight.com

WEEK 3

AN INTRODUCTION TO CHEMICAL EQUILIBRIA

Reversible reactions

A reversible reaction is one which can be made to go in either direction depending on the conditions.

If you pass steam over hot iron the steam reacts with the iron to produce a black, magnetic oxide of iron called triiron tetroxide, Fe₃O₄.

The hydrogen produced in the reaction is swept away by the stream of steam.

Under different conditions, the products of this reaction will also react together. Hydrogen passed over hot triiron tetroxide reduces it to iron. Steam is also produced. This time the steam produced in the reaction is swept away by the stream of hydrogen.       These reactions are reversible, but under the conditions normally used, they become one-way reactions. The products aren’t left in contact with each other, so the reverse reaction can’t happen. Reversible reactions happening in a closed system A closed system is one in which no substances are either added to the system or lost from it. Energy can, however, be transferred in or out at will.

In the example we’ve been looking at, you would have to imagine iron being heated in steam in a closed container. Heat is being added to the system, but none of the substances in the reaction can escape. The system is closed.

As the triiron tetroxide and hydrogen start to be formed, they will also react again to give the original iron and steam. So, if you analysed the mixture after a while, what would you find?

You would find that you had established what is known as a dynamic equilibrium.

A dynamic equilibrium occurs when you have a reversible reaction in a closed system. Nothing can be added to the system or taken away from it apart from energy.

At equilibrium, the quantities of everything present in the mixture remain constant, although the reactions are still continuing. This is because the rates of the forward and the back reactions are equal.

If you change the conditions in a way which changes the relative rates of the forward and back reactions you will change the position of equilibrium – in other words, change the proportions of the various substances present in the equilibrium mixture.

LE CHATELIER’S PRINCIPLE

Using Le Chatelier’s Principle A statement of Le Chatelier’s Principle If a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract the change.   Using Le Chatelier’s Principle with a change of concentration Suppose you have an equilibrium established between four substances A, B, C and D.  What would happen if you changed the conditions by increasing the concentration of A According to Le Chatelier, the position of equilibrium will move in such a way as to counteract the change. That means that the position of equilibrium will move so that the concentration of A decreases again – by reacting it with B and turning it into C + D. The position of equilibrium moves to the right. This is a useful way of converting the maximum possible amount of B into C and D. You might use it if, for example, B was a relatively expensive material whereas A was cheap and plentiful. What would happen if you changed the conditions by decreasing the concentration of A? According to Le Chatelier, the position of equilibrium will move so that the concentration of A increases again. That means that more C and D will react to replace the A that has been removed. The position of equilibrium moves to the left.

This is esssentially what happens if you remove one of the products of the reaction as soon as it is formed. If, for example, you removed C as soon as it was formed, the position of equilibrium would move to the right to replace it. If you kept on removing it, the equilibrium position would keep on moving rightwards – turning this into a one-way reaction.

Using Le Chatelier’s Principle with a change of pressure

This only applies to reactions involving gases:

What would happen if you changed the conditions by increasing the pressure?

According to Le Chatelier, the position of equilibrium will move in such a way as to counteract the change. That means that the position of equilibrium will move so that the pressure is reduced again.

Pressure is caused by gas molecules hitting the sides of their container. The more molecules you have in the container, the higher the pressure will be. The system can reduce the pressure by reacting in such a way as to produce fewer molecules.

In this case, there are 3 molecules on the left-hand side of the equation, but only 2 on the right. By forming more C and D, the system causes the pressure to reduce.

Increasing the pressure on a gas reaction shifts the position of equilibrium towards the side with fewer molecules.

What would happen if you changed the conditions by decreasing the pressure?

The equilibrium will move in such a way that the pressure increases again. It can do that by producing more molecules. In this case, the position of equilibrium will move towards the left-hand side of the reaction.

What happens if there are the same number of molecules on both sides of the equilibrium reaction?

In this case, increasing the pressure has no effect whatsoever on the position of the equilibrium. Because you have the same numbers of molecules on both sides, the equilibrium can’t move in any way that will reduce the pressure again.

Using Le Chatelier’s Principle with a change of temperature

For this, you need to know whether heat is given out or absorbed during the reaction. Assume that our forward reaction is exothermic (heat is evolved):

WHAT CHANGES THE POSITION OF AN EQUILIBRIUM?

By ‘position of an equilibrium’ we mean what are the relative amounts of reactants and products.

AND, importantly, what changes the position of an equilibrium?

In other words, what factors affect the position of an equilibrium?

In a reversible reaction, changing the reaction conditions e.g. concentration, pressure or temperature will change the net direction the reaction goes i.e. more to the right (forward) or more to left (backward) and this must inevitably change the position of the equilibrium.

If you enforce a change on a chemical system at equilibrium, then the system will respond to alter the equilibrium position, BUT the system responds in a way to minimise the enforced change.

The most important factors to consider that strongly influence the position of an equilibrium are temperature, pressure (if gases) and concentration (if solution).

This shows that 250 kJ is evolved (hence the negative sign) when 1 mole of A re acts completely with 2 moles of B. For reversible reactions, the value is always given as if the reaction was one-way in the forward direction.

The back reaction (the conversion of C and D into A and B) would be endothermic by exactly the same amount.

What would happen if you changed the conditions by increasing the temperature?

According to Le Chatelier, the position of equilibrium will move in such a way as to counteract the change. That means that the position of equilibrium will move so that the temperature is reduced again.

Suppose the system is in equilibrium at 300°C, and you increase the temperature to 500°C. How can the reaction counteract the change you have made? How can it cool itself down again?

To cool down, it needs to absorb the extra heat that you have just put in. In the case we are looking at, the back reaction absorbs heat. The position of equilibrium therefore moves to the left. The new equilibrium mixture contains more A and B, and less C and D.

If you were aiming to make as much C and D as possible, increasing the temperature on a reversible reaction where the forward reaction is exothermic isn’t a good idea!

What would happen if you changed the conditions by decreasing the temperature?

The equilibrium will move in such a way that the temperature increases again.

Suppose the system is in equilibrium at 500°C and you reduce the temperature to 400°C. The reaction will tend to heat itself up again to return to the original temperature. It can do that by favouring the exothermic reaction.

The position of equilibrium will move to the right. More A and B are converted into C and D at the lower temperature.

• Increasing the temperature of a system in dynamic equilibrium favours the endothermic reaction. The system counteracts the change you have made by absorbing the extra heat.
• Decreasing the temperature of a system in dynamic equilibrium favours the exothermic reaction. The system counteracts the change you have made by producing more heat.

Chatelier’s Principle and catalysts Adding a catalyst makes absolutely no difference to the position of equilibrium, and Le Chatelier’s Principle doesn’t apply to them.   This is because a catalyst speeds up the forward and back reaction to the same extent. Because adding a catalyst doesn’t affect the relative rates of the two reactions, it can’t affect the position of equilibrium. So why use a catalyst?

• For a dynamic equilibrium to be set up, the rates of the forward reaction and the back reaction have to become equal. This doesn’t happen instantly. For a very slow reaction, it could take years! A catalyst speeds up the rate at which a reaction reaches dynamic equilibrium.

WHAT CHANGES THE POSITION OF AN EQUILIBRIUM?

By ‘position of an equilibrium’ we mean what are the relative amounts of reactants and products.

AND, importantly, what changes the position of an equilibrium?

In other words, what factors affect the position of an equilibrium?

In a reversible reaction, changing the reaction conditions e.g. concentration, pressure or temperature will change the net direction the reaction goes i.e. more to the right (forward) or more to left (backward) and this must inevitably change the position of the equilibrium.

If you enforce a change on a chemical system at equilibrium, then the system will respond to alter the equilibrium position, BUT the system responds in a way to minimise the enforced change.

The most important factors to consider that strongly influence the position of an equilibrium are temperature, pressure (if gases) and concentration (if solution).

What are the RULES GOVERNING THE POSITION OF A CHEMICAL EQUILIBRIUM?

For industrial processes, it is important to maximise the concentration of the desired products and minimise the ‘leftover’ reactants. A set of rules can be used to predict the best reaction conditions to give the highest possible yield of product.

The three rules outlined below are known as Le Chatelier’s Principle. This essentially states that if a change is imposed on a system, the system will change to minimise the enforced change to re-establish equilibrium.

 Le Chatelier’s Principle: Rule 1 The effect of temperature change on the position of an equilibrium

Reminder: If a forward reaction is exothermic, the reverse backward reaction is endothermic and vice versa.

If the temperature of a chemical system at equilibrium is increased then the relative amount of products at equilibrium increases for an endothermic reaction, but the relative amount of products at equilibrium decreases for an exothermic reaction.

If the temperature of a chemical system at equilibrium is decreased: then the relative amount of products at equilibrium decreases for an endothermic reaction, but the relative amount of products at equilibrium increases for an exothermic
reaction.

Rule 1a: If the forward reaction forming the products is endothermic, raising the temperaturefavours its formation increasing the yield of product (lowering the temperature decreases the yield).

So increasing temperature favours the endothermic direction reaction.

The system attempts to absorb the heat and minimise the increase in temperature.

Rule 1b: If the forward reaction forming the products is exothermic, decreasing the temperature favours its formation (increasing temperature decreases the yield).

So decreasing temperature favours the exothermic direction reaction

The system attempts to release heat to minimise the temperature decrease.

Rule 1 examples

The equilibrium between hydrogen gas, gaseous iodine and gaseous hydrogen iodide.

H₂(g) + I_2(g) 2HI(g) (plus 10 kJ of heat energy, exothermic L to R)

Increasing temperature favours the endothermic direction, backward reaction, some hydrogen iodide will decompose.

Decreasing temperature favours the exothermic reaction, so more hydrogen and iodine react to form hydrogen iodide.

Le Chatelier’s Principle: Rule 2 The effect of changing pressure on the position of an equilibrium

You can increase/decrease the pressure by decreasing/increasing the volume of the gases OR increasing/decreasing the concentration of gases in the same volume.

For reactions involving gases at equilibrium, an increase in pressure causes the equilibrium position to shift towards the side with the smallest number of gaseous molecules as indicated by the balanced symbol equation for that reaction.

A decrease in pressure of chemical reaction system involving gases, causes the equilibrium position to shift towards the side with the larger number of gaseous molecules as indicated by the balanced symbol equation for that reaction.

A correctly balanced equation is important, because ALL gaseous molecules shown in the equation must be taken into account.

Rule 2a:Increasing the pressurefavours the side of the equilibrium with the least number of gaseous molecules as shown by the balanced symbol equation.

So increasing pressure favours the reaction direction to reduce the number of gaseous molecules.

The system is changing to minimise the impact of the increase in pressure by removing some gas molecules.

Rule 2b: Decreasing the pressure favours the side of the equilibrium with the most number of gaseous molecules as shown by the balanced symbol equation.

So decreasing pressure favours the reaction direction to produce the most gaseous molecules.

The system is changing to minimise the impact of the decrease in pressure by increasing the number of gas molecules.

Rule 2 examples

(i) N₂(g) + 3H₂(g) 2NH₃(g)

4 gas molecules ==> 2 gas molecules, so to re-establish a dynamic equilibrium …

Increase in pressures favours the forward reaction to reduce the number of gas molecules, so more ammonia formed.

Decrease in pressure encourages the formation of more gas molecules, so some of the ammonia decomposes into nitrogen and hydrogen.

(ii) N₂O₄(g) 2NO₂(g)

1 gas molecule ==> 2 gas molecules, so to re-establish a dynamic equilibrium …

Increase in pressure favours backward direction to reduce the number of gaseous molecules and give more dinitrogen tetroxide.

Decrease in pressure encourages more gas molecules to form, so the forward reaction gives more nitrogen dioxide.

(iii) N₂(g) + O₂(g) 2NO(g)

2 gas molecules ==> 2 gas molecules

Change in pressure has no effect on equilibrium position

Rules 1 above, and rule 3, below, apply to any reaction, BUT rule 2 above, ONLY applies to a reaction with one or gaseous reactants or gaseous products.

Increase in pressure does not influence the concentration of substances in a solution or solid mixture because they are too dense to be significantly compressed i.e. no effective change in concentration.

The situation is quite different in gases where is a lot of space between the molecules to compress them closer together.

If a reaction involves gases BUT there are equal numbers of gaseous molecules on each side of the equation, increasing or decreasing pressure has no effect on the position of the equilibrium.

e.g the equilibrium position of the reaction to form hydrogen iodide from hydrogen and iodine

H₂(g) + I₂(g) 2HI(g)

is unaffected by change in pressure, the are two molecules (or moles) of gas on each side of the equation.

Le Chatelier’s Principle: Rule 3 The effect of changing concentration on the position of an equilibrium

If the concentration of any of the reactants or products is changed, the system cannot any longer be at equilibrium.

The concentrations of all the substances will change until equilibrium is reached again.

If the concentration of a reactant is increased, more products will be formed until equilibrium is reached again.

If the concentration of a product is decreased, more reactants will react until equilibrium is reached again.

Rule 3a: If the concentration of a reactant (on the left) is increased, then some of it must change to the products (on the right) to maintain a balanced equilibrium position.

Rule 3b: If the concentration of a reactant (on the left) is decreased, then some of  the products (on the right) must change back to reactants to maintain a balanced equilibrium position.

Rule 3 examples

• e.g. nitrogen + hydrogen ammonia
  • or N_2(g) + 3H_2(g) 2NH_3(g)
  • If the nitrogen or hydrogen concentration was increased, some of this extra gas would change to ammonia.
  • If the nitrogen or hydrogen concentration was decreased, some of ammonia would change back to nitrogen and hydrogen.
  • At AS-A2 advanced level things can get more complicated e.g. can you figure out why in terms of concentration to maintain the equilibrium balance? (and if a gcse student, don’t worry if you can’t) …
  • So in terms of enforced change ==>system response:
  • Increasing nitrogen concentration ==> decreases hydrogen concentration and increases ammonia concentration
  • Increasing hydrogen concentration ==> decreases nitrogen concentration and increases ammonia concentration
  • Increasing ammonia concentration ==> increases both nitrogen and hydrogen concentrations
  • Decreasing ammonia concentration ==> decreases both nitrogen and hydrogen concentration
  • Decreasing nitrogen concentration ==> increases hydrogen concentration and decreases ammonia concentration
  • Decreasing hydrogen concentration ==> increases nitrogen concentration and decreases ammonia concentration

Le Chatelier’s Principle: Rule 4 The effect of using a catalyst on the position of an equilibrium

A catalyst does NOT affect the position of an equilibrium.

You just get to the equilibrium position here faster!

A catalyst usually speeds up both the forward and reverse reaction but there is no way it can influence the final ‘balanced’ concentrations.

However, the importance of a catalyst lies with economics e.g.

(i) bringing about reactions with high activation energies at lower temperatures and so saving the cost on energy,

(ii) and saving time is saving money, i.e. a catalyst increases the efficiency of the chemical process e.g. the Haber synthesis of ammonia.

Rule 4 examples

Iron catalyst in the synthesis of ammonia.

Vanadium pentoxide catalyst in the Contact Process for manufacturing sulfuric acid.

Both of these chemical processes are faster and made economically more efficient by use of a catalyst, but you don’t get a greater % yield in the final reacted mixture.

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Applying the rules 1 to 4 to some chemical processes

(a) The formation of calcium oxide (lime) and carbon dioxide from calcium carbonate (limestone)

CaCO_3(s) CaO_(s) + CO_2(g)

The forward reaction is endothermic, 178kJ of heat energy is absorbed (taken in) for every mole of calcium oxide formed.

One mole of gas is formed in the process, so there is a net increase in the moles of gas in lime formation, since there are no gaseous reactants.

From rule 1: increasing the temperature will increase the yield of calcium oxide or lime, CaO which is endothermically formed.

From rule 2: decreasing the pressure will favour the formation of more gas molecules if possible, so more carbon dioxide formed, and hence more lime.

Lime is made commercially by heating limestone to a high temperature (e.g. 1000^oC) in a limekiln that is well ventilated (this reduces the carbon dioxide pressure and so reduces the un-desired backward reaction).

---

(b) The formation of hydrogen chloride from hydrogen and chlorine

H_2(g) + Cl_2(g) 2HCl_(g)

The forward reaction is very exothermic, 184kJ of heat energy is given out in forming hydrogen bromide according to the above equation (184/2 = 92kJ per mole of HCl formed).

There is no net change in the moles of gas (2 moles reactants 2 moles of product)

From rule 1: decreasing the temperature favours the exothermic formation of hydrogen chloride, so the equilibrium moves proportionately to the right-hand side (more HCl, less H₂ or Cl₂). If hydrogen chloride is heated to a very high temperature, endothermic direction, then more HCl decomposes into H₂ or Cl₂.

From rule 2: since there is no net change in the number of moles of gas on reaction, pressure has no effect on the yield of hydrogen chloride and the proportions of HCl, H₂ or Cl₂ stay the same.

(c) The formation of ammonia

nitrogen + hydrogen ammonia

N_2(g) + 3H_2(g) 2NH_3(g)     (plus 92 kJ heat energy, exothermic)

Ammonia formation is favouredby …

Rule 1. Lowering the temperature, because it is an exothermic reaction, but this may make it too slow, compromise required.

Rule 2. Increasing pressure, because there is a reduction in the molecules of gas, but the higher the pressure, the more costly the engineering.

Rule 4. An iron catalyst speeds up the reaction, but has no effect on the % ammonia in the reacted mixture exiting the reactor chamber in the chemical plant of a Haber synthesis process.

In the exam you may have to explain more about the rule and how they are used, which is what this page is all about, plus some more examples below

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(d) The manufacture of sulfur trioxide

The process of making sulfur trioxide from sulfur dioxide is one stage in the manufacture of sulfuric acid by the Contact Process.

The equilibrium equation is given below.

2SO_2(g) + O_2(g) 2SO_3(g)     (plus 95 kJ heat energy, exothermic)

3 gas molecules ==> 2 gas molecules.

Rule 1. The exothermic reaction is favoured by a lower temperature, but this may be too slow, so a compromise temperature of around 450oC is used, which gives a fast economic rate of sulphur trioxide production.

Rule 2. The reaction is favoured by high pressure (pressure equilibrium rule, 3 => 2 gas molecules, LHS ==> RHS), but only a small increase in pressure is used to give high yields of sulphur trioxide, because the formation of SO₃ on the right hand side is so energetically favourable (approx. 99% yield, i.e. only about 1% SO₂ unreacted).

Rule 4. The use of the V₂O₅ catalyst ensures a fast reaction without having to use too a higher temperature which would favour the left hand side and reduce the yield BUT it does not change the % of sulphur trioxide formed, you simply get there faster.

EVALUATION

1.         How do the equilibrium positions of the following  reaction change with pressure?What is the equilibrium constant and the units?

a.N₂(g)+3H₂(g)=2NH₃(g)    b.H₂(g)+I₂(g)=2HI(g)

2. State Le Chatellier’s Principle

3. Consider the equation for the equilibrium reaction

N₂(g)+3H₂-2NH₃(g) )   ∆H=-92KJmol^-1.

Write the equilibrium constant of the reaction.

4. Consider the following equilibrium reaction.

X₂(g) +2Y(g) –XY₂(g) )   ∆H=-52KJmol^-1

i.State what happens to the yield of XY2 when the temperature is increased.

ii.Explain the effect of decrease in pressure on the equilibrium position

iii.State the effect of a catalyst position of equilibrium.

Notes on Chemistry SS2 – Edudelight.com

WEEK 4

History of Hydrogen

 Hydrogen comes from Greek meaning “water producer” (“hydro” =water and “gennao”=to make). First isolated and identified as an element by Cavendish in 1766, hydrogen was believed to be many different things. Cavendish himself thought that it was “inflammable air from metals”, owing to its production by the action of acids on metals. Before that, Robert Boyle and Paracelsus both used reactions of iron and acids to produce hydrogen gas and Antoine Lavoisier gave hydrogen its name because it produced water when ignited in air. Others thought it was pure phlogiston because of its flammability. Hydrogen is among the ten most abundant elements on the planet, but very little is found in elemental form due to its low density and reactivity. Much of the terrestrial hydrogen is locked up in water molecules and organic compounds like hydrocarbons.

Naturally Occurring Hydrogen

Hydrogen is the fuel for reactions of the Sun and other stars (fusion reactions). Hydrogen is the lightest and most abundant element in the universe. About 70%- 75% of the universe is composed of hydrogen by mass. All stars are essentially large masses of hydrogen gas that produce enormous amounts of energy through the fusion of hydrogen atoms at their dense cores. In smaller stars, hydrogen atoms collided and fused to form helium and other light elements like nitrogen and carbon(essential for life).  In the larger stars, fusion produces the lighter and heavier elements like calcium, oxygen, and silicon.

On Earth, hydrogen is mostly found in association with oxygen; its most abundant form being water (H₂O). Hydrogen is only .9% by mass and 15% by volume abundant on the earth, despite water covering about 70% of the planet.  Because hydrogen is so light, there is only 0.5 ppm (parts per million) in the atmosphere, which is a good thing considering it is EXTREMELY flammable.  

Other Sources of Hydrogen

Hydrogen gas can be prepared by reacting a dilute strong acid like hydrochloric acids with an active metal. The metal becomes oxides, while the H⁺ (from the acid) is reduced to hydrogen gas. This method is only practical for producing small amounts of hydrogen in the lab, but is much too costly for industrial production:  

Zn (s) +2H + (aq) →Zn 2+ (aq) +H 2(g)  

The purest form of H₂(g) can come from electrolysis of H₂O(l), the most common hydrogen compound on this plant. This method is also not commercially viable because it requires a significant amount of energy (ΔH=572kJ  ): 

2H 2 O (l) →2H 2(g) +O 2(g)  

\(H_{_2}O\) is the most abundant form of hydrogen on the planet, so it seems logical to try to extract hydrogen from water without electrolysis of water. To do so, we must reduce the hydrogen with +1 oxidation state to hydrogen with 0 oxidation state (in hydrogen gas). Three commonly used reducing agents are carbon (in coke or coal), carbon monoxide, and methane. These react with water vapor form H₂(g):

C (s) +2H 2 O (g) →CO(g)+H 2(g)  

CO (g) +2H 2 O (g) →CO2+H 2(g)  

Reforming of Methane:

CH 4(g) +H 2 O (g) →CO(g)+3H 2(g)  

These three methods are most industrially feasible (cost effective) methods of producing H₂(g). 

Isotopes

There are two important isotopes of hydrogen. Deuterium (²H) has an abundance of 0.015% of terrestrial hydrogen and the nucleus of the isotope contains one neutron.

 Three Hydrogen Isotopes

Protium (¹H) is the most common isotope, consisting of 99.98% of naturally occurring hydrogen. It is a nucleus containing a single proton.

Deuterium (²H) is another an isotope containing a proton and neutron, consisting of only 0.0156% of the naturally occurring hydrogen. Commonly indicated with symbol D and sometimes called heavy hydrogen, deuterium is separated by the fractional distillation of liquid hydrogen but it can also be produced by the prolonged electrolysis of ordinary water. Approximately 100,000 gallons of water will produce a single gallon of D₂O, “heavy water”. This special kind of water has a higher density, melting point, and boiling point than regular water and used as a moderator in some fission power reactors. Deuterium fuel is used in experimental fusion reactors. Replacing protium with deuterium has important uses for exploring reaction mechanisms via the kinetic isotope effect.

Tritium (³H) contains two neutrons in its nucleus and is radioactive with a 12.3-year half-life, which is continuously formed in the upper atmosphere due to cosmic rays.  It is can also be made in a lab from Lithium-6 in a nuclear reactor. Tritium is also used in hydrogen bombs. It is very rare (about 1 in every 1,018 atoms) and is formed in the environment by cosmic ray bombardment. Most tritium is manufactured by bombarding Li with neutrons. Tritium is used in thermonuclear weapons and experimental fusion reactors

Hydrogen is prepared in the laboratory:

1. by the action of a dilute strong acid on metals, such as zinc:

Zn + H₂SO₄ → ZnSO₄ + H₂­;

2. by reaction amphoteryc metals with a strong base, such as sodium hydroxide:

2Al + 6NaOH + 6H₂O → 2Na₃[Al(OH)₆]₃ + 3H₂­,

Zn  + 2NaOH + 2H₂O → Na₂[Zn(OH)₄] + H_2­­;

3. by electrolysis of water:

Industrially, hydrogen is prepared from water and hydrocarbons. Until re­cently the water-gas reaction was an important way of hydrogen preparing. The water-gas reaction is an industrial process in which steam is passed over red-hot coke giving a gaseous mixture of carbon monoxide and hydrogen:

C + H₂O(g)→ CO + H₂.

Now two general processes are used, both starting with natural gas or petroleum hydrocarbons. The steam-reforming process is an industrial preparation of hy­drogen and carbon monoxide mixtures by the reaction of steam and hydrocarbons at high temperature and pressure over a nickel catalyst. For example,

CH₄ + H₂O(g) → CO + 3H₂

C₂H₈ + 3H₂O(g) → 3CO + 7H₂ .

The second process involves the partial oxidation of hydrocarbons. Natural gas, for example, is mixed with a limited supply of oxygen and burned at elevated pres­sures:

2CH₄  +  O₂→ 2CO + 4H₂.

                  natural gas

As natural gas and petroleum become more expensive, the water-gas reaction may become widely used again.

The gaseous mixture of carbon monoxide and hydrogen from these reactions (called synthesis gas) is used in the preparation of methanol. Hydrogen is obtained from this mixture free of carbon monoxide by means of the water-gas shift reaction,an industrial process in which carbon monoxide reacts with steam in the presence of a catalyst producing carbon dioxide and hydrogen:

CO + H₂O → CO₂ + H₂.

Carbon dioxide is removed by dissolving the gas in a basic solution to give carbon­ate ion.

1. Hydrogen reacts with many nonmetals. In these reactions it derives  hydrogen-cation, H⁺. Hydrogen combines with nitrogen in the presence of a catalyst forming ammonia:

3H₂ + N₂→ 2NH₃;

with sulfur forming hydrogen sulfide:

H₂ + S → H₂S;

with chlorine forming hydrogen chloride:

H₂ + Cl₂→ 2HCl;

and with oxygen forming water:

2H_2­ + O₂→ 2H₂O.

The reaction of oxygen and hydrogen takes place at room temperature in the presence of a catalyst such as finely divided platinum only. When hydrogen is mixed with the air or oxygen and ignited, the mixture explodes.

2. Hydrogen also combines with some metals. Because of the small electron affinity of the hydrogen atom (73 kJ/rnol, com­pared with 349 kJ/mol for Cl), ionic compounds containing the hydride ion, H^–, are formed with metals of the lowest ionization energy only: the alkali metals and the alkaline earth metals. Thus, sodium and calcium react with hydrogen gas at mod­erate temperatures giving the hydrides:

2Na + H₂→ 2NaH;

Ca + H₂→ CaH₂.

These ionic hydrides conduct electricity when molten, indicating the presence of ions. Hydrogen liberates at the electrode connected to the positive terminal of the battery, according to the electrode reaction:

2H^–→ H₂ + 2e^–.

Ionic hydrides react with water, giving hydrogen:

NaH + H₂O → NaOH + H₂­;

CaH₂ + 2H₂O→ Ca(OH)₂ + 2H₂­.

According to the last reaction, calcium hydride is a convenient, portable source of small quantities of hydrogen.

3. It acts as a reducing agent on metallic oxides, such as copper oxide, removing the oxygen and leaving the metal in a free state:

CuO + H₂→ Cu + H₂O.

4. Hydrogen reacts with unsaturated organic compounds forming corresponding saturated compounds:

CH₂=CH₂ + H₂→ CH₃–CH₃ .

General Test for Hydrogen

1)  Hydrogen gas, H_2(g) has no colour or smell.

2)  Hydrogen gas has no effect on moistlitmus paper
or moist universal indicator paper – it is neutral.

3)  Hydrogen gas burns with a characteristic ‘pop’.

 Specific Test for Hydrogen Gas?

Hydrogen gas is recognised by the ‘pop’ when it burns.
The ‘pop’ is the sound of a small explosion.
Hydrogen gas is highly flammable!

hydrogen  +  oxygen    water  (hydrogen oxide).
2H_2(g)    +    O_2(g)           2H₂O_(l)                

This highly exothermic reaction is an example of oxidation.

Collection of Hydrogen

Hydrogen gas is less dense than air and can be collected
by upward delivery, over water or by using a gas syringe

Uses & Application

The vast majority of hydrogen produced industrially today is made either from treatment of methane gas with steam or in the production of “water gas” from the reaction of coal with steam. Most of this hydrogen is used in the Haber process to manufacture ammonia.

Hydrogen is also used for hydrogenation vegetable oils, turning them into margarine and shortening, and some is used for liquid rocket fuel. Liquid hydrogen (combined with liquid oxygen) is a major component of rocket fuel (as mentioned above combination of hydrogen and oxygen relapses a huge amount of energy). Because hydrogen is a good reducing agent, it is used to produce metals like iron, copper, nickel, and cobalt from their ores. 

Because one cubic feet of hydrogen can lift about 0.07 lbs, hydrogen lifted airships or Zeppelins became very common in the early 1900s.However, the use of hydrogen for this purpose was largely discontinued around World War II after the explosion of The Hindenburg; this prompted greater use of inert helium, rather than flammable hydrogen for air travel.

Recently, due to the fear of fossil fuels running out, extensive research is being done on hydrogen as a source of energy.Because of their moderately high energy densities liquid hydrogen and compressed hydrogen gas are possible fuels for the future.A huge advantage in using them is that their combustion only produces water (it burns “clean”). However, it is very costly, and not economically feasible with current technology.

Combustion of fuel produces energy that can be converted into electrical energy when energy in the steam turns a turbine to drive a generator.  However, this is not very efficient because a great deal of energy is lost as heat.  The production of electricity using voltaic cell can yield more electricity (a form of usable energy).  Voltaic cells that transform chemical energy in fuels (like H₂ and CH₄) are called fuel cells.   These are not self-contained and so are not considered batteries.  The hydrogen cell is a type of fuel cell involving the reaction between H₂(g) with O₂(g) to form liquid water; this cell is twice as efficient as the best internal combustion engine.  In the cell (in basic conditions), the oxygen is reduced at the cathode, while the hydrogen is oxidized at the anode.

Reduction: O₂(g)+2H₂O(l)+4e^– → 4OH^–(aq) 

 Oxidation: H₂(g) + 2OH^–(aq) → 2H₂O(l) + 2e-

Overall: 2H₂(g) + O₂(g) → 2H₂O(l)                    

E°cell= Reduction- Oxidation= E_O2/OH^– – E_H2O/H2 = 0.401V – (-0.828V) = +1.23

However, this technology is far from being used in everyday life due to its great costs.    

 Image of A Hydrogen Fuel Cell.

EVALUATION

1.Using relevant chemical equations,state the three general methods of preparing Hydrogen in the Air.

2. Mention 3 chemical properties of hydrogen giving balanced chemical equation.

3. What are the general test for Hydrogen?

4. What are the specific test for Hydrogen?

5. How is Hydrogen collected in the Laboratory?

WEEK 5 

OXYGEN AND ITS COMPOUNDS
Oxygen is the most abundant element on earth. It is found almost in everything in nature and also in free state. Oxygen is a gas under ordinary conditions of temperature and pressure and it is a gas that is colorless, odorless, tasteless and slightly heavier than air, all these physical properties are characteristic of this element.

Laboratory Preparation of Oxygen

Oxygen can be prepared in a number of ways in the laboratory. These include:

(a). Heating a mixture of potassium trioxochlorate(V), KClO₃ and maganese(IV) oxide, MnO₂.

The MnO₂ acts as a catalyst. The reaction is actually the decomposition of KClO₃.

2KClO_3(s) → 2KCl_(s) + 3O_2(g)

The laboratory setup for preparation of oxygen is shown below.

(b). The decomposition of hydrogen peroxide, H₂O₂ using manganese(IV) oxide – this does not require heating.

Hydrogen peroxide is added drop wisely unto MnO₂.

2H₂O_2(aq) → 2H₂O_(l) + O_2(g)

Or by the drop wise addition of hydrogen peroxide on acidified KMnO₄

5H₂O_2(aq) + 2KMnO_4(aq) + 3H₂SO_4(aq)  → K₂SO_4(aq) + 2MnSO_4(aq) + 8H₂O_(l) + 5O_2(g)

(c). The reaction between water and sodium peroxide.

Hydrogen peroxide is formed, and it immediately decomposes by the catalytic effect of the OH^– ions in solution.

Na₂O_2(s) + 2H₂O → H₂O₂ + 2Na⁺ + 2OH^–

Then, 2H₂O₂ → 2H₂O + O₂

(d). Application of heat on trioxonitrate(V) salts of metals.

Trioxonitrate(V) salts of metals, e.g, sodium trioxonitrate(V), NaNO₃ give off a part of their oxygen upon being heated. NaNO₃ loses one third of its oxygen.

2NaNO_3(s) → 2NaNO_2(s) + O_2(g)

(e). Application of heat on certain oxides of the least active metals, e.g., mercury oxide, HgO and silver oxide, Ag₂O. These oxides undergo complete dissociation when heated.

2HgO_(s) → 2Hg_(s) + O_2(g)

2Ag₂O_(s) → 4Ag_(s) + O_2(g)

(f). Application of heat on oxides of certain metals with more than one oxidation state, e.g., lead(IV) oxide, PbO₂ and mangan- ese(IV) oxide, MnO₂.

These oxides give off only a part of their oxygen when they are heated. Such reactions usually require very high temperatures.

2PbO_2(s) → 2PbO_(s) + O_2(g)

3MnO_2(s) → Mn₃O_4(s) + O_2(g)

(g). The electrolysis of water.

Oxygen is produced at the anode.

Test for Oxygen

Oxygen can be distinguished from all other gases except dinitrogen oxide, N₂O by its rekindling of a glowing splint of wood.

It is however distinguished from N₂O by the following observations:

(1). Oxygen does not have smell, while N₂O has a sweet, sickly smell.

(2). Oxygen produces brown fumes of nitrogendioxide, NO₂ with nitrogen monoxide,

2NO_(g) + O_2(g) → 2NO_2(g) (brown fumes),

while N₂O does not.

Properties of Oxygen

Physical Properties:

It is colourless, has no odour and is neutral.
 It is slightly soluble in water.

Density – The density increases down the group with the increase in atomic number.

Melting point and boiling point – It shows a regular increase down the group.

Ionization energy – These elements possess large values on ionization energies which decreases gradually from O to Po. The decrease in ionization energy from oxygen to polonium is due to increase in size of the atoms and increase in screening effect of the electrons belonging to lower shells.

Metallic and non-metallic character – Metallic character depends on ionization energy. Lesser the ionization energy greater will be the metallic character. As we move down the group the ionization energy decreases and therefore the metallic character increases.

Catenation – In this group only S has a strong tendency of catenation. Oxygen also shows this tendency to a limited extent.

Oxidation state – Oxygen being first member in this group shows -2 oxidation state in its compounds owing to its high electronegativity. It also exhibit -1 in H₂O₂ and zero in O₂. Unlike oxygen the other elements have a tendency to show +2, +4 and +6 oxidation states.

Multiple bond formation – The tendency of these elements to form multiple bonds to carbon and nitrogen decreases as we move down the group from sulfur to tellurium.

Electron affinity – These elements have high electron affinity and on moving down the group the size of the atom increases and distance of valence shell from nucleus increases. The additional electron feels lesser attraction as we move down the group and hence electron affinity decreases.

Electronegativity – Oxygen is strongly electronegative in character. The value of electronegativity decreases with increase in atomic number down the group. This is due to increase in size of the atoms and decrease in effective nuclear charge as we move down the group.

Elemental state – Oxygen exist in a diatomic gaseous state at room temperature while other elements are solids.

Allotropy – All elements of the group exhibit allotropy. Oxygen exists in diatomic and triatomic states. Sulfur is found as rhombic, plastic and colloidal sulfur. Selenium has rhombic, mono clinic and grey forms. The grey form is the most stable form and consists of regularly arranged spirals of selenium atoms
It is very active and reacts with many metals and non-metals to form basic and acidic oxides respectively.

Basic oxide – examples include MgO, Na₂O and CaO;
Acidic oxide – examples are CO₂, SO₂ and P₄O₁₀ (acidic oxides of non-metal are also called acid anhydrides).

Summary of Reactivity of Oxygen with Metals

K, Na, Ca, Mg, Al, Zn, Fe, Pb, Cu – – show decreasing readiness to form oxides when heated in air, with Cu the least reactive with oxygen.

Hg, Ag, Au – – these metals show the least readiness to form oxides. Their oxides are easily decomposed to the
metal and oxygen.

K, Na, Ca, Mg, Al, Zn – – the oxides of these metals are not reduced to the metals by heating in a stream of hydrogen,
carbon or carbon(II) oxide.

Fe, Pb, Cu – – the oxides of these metals are reduced to the metals by heating in a stream of hydrogen, carbon or carbon(II) oxide.

Note: the more readily a metal combines with oxygen to form an oxide, the less readily it will be reduced to the metal by either heating in a stream of hydrogen or CO.

Therefore, the oxides of K to Zn are not reduced, while those of Fe and below are reduced.

Chemical Properties:

(1). Reaction with compounds – most hydrocarbons and compounds of carbon, hydrogen and oxygen burn in oxygen to form carbon (IV) oxide and water.

CH_4(g) + 2O_2(g) → CO_2(g) + 2H₂O_(g)

C₂H₅OH_(l) + 3O_2(g) → 2CO_2(g) + 3H₂O_(g)

4NH_3(g) + 5O_2(g) → 4NO_(g) + 6H₂O_(l)

Formation of oxides

The elements of this group form a number of oxides. Some of the oxides formed by these elements are SO, TeO, PoO, SO₂, SeO₂, TeO₂, PoO₂, SO₃, SeO₃ and TeO₃. All elements form monoxide. All elements form dioxide with formula MO₂. SO₂ is a gas, SeO₂ is a volatile solid while TeO₂ and PoO₂ are non-volatile crystalline solids.

Selenium and tellurium dioxides are solids having polymeric chain. The acidic character of the dioxides decreases as we move down the group.

Formation of oxyacids

Theoxyacids of sulfur are more numerous and important than those of Se and Te. The list of oxyacids of sulfur are H₂SO₄, H₂SO₃, H₂SO₅ (Caros acid), H₂S₂O₈ (Marshells acid), H₂S₂O₃, H₂S₂O₆ and H₂S₂O₇. Selenium form selenous acid (H₂SeO₃) and Selenic acid (H₂SeO₄) and Tellurium form tellurous acid (H₂TeO₃) and telluric acid (H₂TeO₄).

Oxygen Family Elements

Elements listed below belongs to Group 16 (VI A) of the long form of periodic table. Except Po all other elements are also called chalogens, which means ore or mineral forming elements. Out of all these elements oxygen, sulfur and tellurium are non metals while polonium is a radioactive metal.

Atoms of these elements have the outer electronic configuration as ns²np⁴, where n varies from 2 to 6.

S.No	Element	Atomic number	Electronic configuration
1	Oxygen	8	[He]2s22p4
2	Sulfur	16	[Ne]3s23p4
3	Selenium	34	[Ar]3d104s24p4
4	Tellurium	52	[Kr]4d105s25p4
5	Polonium	84	[Xe]4f145d106s26p4

There are two half filled p-orbitals which are used for bonding with other elements. From the similar outer electronic configuration of these elements it is expected that they will show similar physical and chemical properties.

Oxygen Family Facts

1. Oxygen is the most abundant element (46.5%) in the earth crust.
2. Joseph Priestley obtained oxygen by focusing the sun light on mercuric oxide in 1774.
3. About 21% of the components of dry air consist of O₂ molecules.
4. In combined state it is also present in plants and animals tissues.
5. Ozone is the second form of oxygen and considered as its allotrope.
6. It has three isotopes of atomic weight 16, 17 and 18 in the ratio of about 10,000:1:8.
7. Oxygen always exist as diatomic (O₂) molecule.
8. With metallic elements oxygen forms ionic compounds.
9. Oxygen reacts with non metals to form a large number of molecular compounds.
10. Only the noble gases and few inactive metals such as Pt and Au do not react directly with oxygen.

Oxygen Family Name

1. The elements oxygen, sulfur, selenium, tellurium and polonium constitute of group VI A of periodic table.
2. In the long form of periodic table they find place in group 16 and p-block elements.
3. First four members of this group (O, S, Se, Te) are also known as chalcogens from Greek words “chalkos” for copper and “genes” for “born”. Most copper minerals are either oxygen or sulfur compounds.
4. Polonium is a radioactive element and is the most metallic in the group.

Uses of Oxygen

(1). Used as an aid for breathing where problem of breathing arises. Example in high altitude flying or climbing, and when a patient is under anesthetics.

(2). In the oxyacetylene (i.e. oxygen-ethyne) flame – used in welding and cutting steel plate – due to the very high temperature of the flame (about 2200^oC).

(3). Used in the L-D process for making steel.

Commercial Production of Oxygen from Liquid Air

The process involves air firstly liquefied by compression, cooling, expansion and successive cooling. Then by fractional distillation of the liquid air – oxygen is separated.

Liquid air contains mainly oxygen and nitrogen. Nitrogen evolves first at 77 K while oxygen evolves later at 90 K, 760 mm Hg.
The laboratory setup for preparation of oxygen is shown below.

Compounds of Oxygen

Oxygen in organic compound can occurs in a variety of forms in nature. The list of oxygen compounds are alcohols, ethers, phosphates, aldehydes, ketones, amide, esters, oxides, carbonyl compounds etc. Oxygen compounds are given in the below table with their general formula.

S.No	Compounds	General formula
1	Alcohols	R-OH
2	Ethers	R–O–R
3	Phosphates	(HPO3)x
4	Aldehydes	RCHO
5	Ketones	R-CO-R
6	Amide	RC(O)NHR’
7	Esters	RCOOR’
8	Oxides	M2O

EVALUATION

1.Describe the properties of oxygen.

2.State all the allotropes of oxygen.

3.Draw a well labeled diagram of the laboratory preparation of oxygen

4.Mention the only allotropes of Oxygen.How is it prepared?

5.What are the uses of oxygen

Notes on Chemistry SS2 – Edudelight.com

WEEK 6

THE HALOGENS

There are six elements in Group VIIA, the next-to-last column of the periodic table. As expected, these elements have certain properties in common. They all form diatomic molecules (H₂, F₂, Cl₂, Br₂, I₂, and At₂), for example, and they all form negatively charged ions (H^–, F^–, Cl^–, Br^–, I^–, and At^–).

When the chemistry of these elements is discussed, hydrogen is separated from the others and astatine is ignored because it is radioactive. (The most stable isotopes of astatine have half-lives of less than a minute. As a result, the largest samples of astatine compounds studied to date have been less than 50 ng.) Discussions of the chemistry of the elements in Group VIIA therefore focus on four elements: fluorine, chlorine, bromine, and iodine. These elements are called the halogens (from the Greek hals, “salt,” and gennan, “to form or generate”) because they are literally the salt formers.

None of the halogens can be found in nature in their elemental form. They are invariably found as salts of the halide ions (F^–, Cl^–, Br^–, and I^–). Fluoride ions are found in minerals such as fluorite (CaF₂) and cryolite (Na₃AlF₆). Chloride ions are found in rock salt (NaCl), the oceans, which are roughly 2% Cl^– ion by weight, and in lakes that have a high salt content, such as the Great Salt Lake in Utah, which is 9% Cl^– ion by weight. Both bromide and iodide ions are found at low concentrations in the oceans, as well as in brine wells in Louisiana, California, and Michigan.

The Halogens in their Elemental Form

Fluorine (F₂), a highly toxic, colorless gas, is the most reactive element known so reactive that asbestos, water, and silicon burst into flame in its presence. It is so reactive it even forms compounds with Kr, Xe, and Rn, elements that were once thought to be inert. Fluorine is such a powerful oxidizing agent that it can coax other elements into unusually high oxidation numbers, as in AgF₂, PtF₆, and IF₇.

Fluorine is so reactive that it is difficult to find a container in which it can be stored. F₂ attacks both glass and quartz, for example, and causes most metals to burst into flame. Fluorine is handled in equipment built out of certain alloys of copper and nickel. It still reacts with these alloys, but it forms a layer of a fluoride on the surface that protects the metal from further reaction.

Fluorine is used in the manufacture of Teflon or poly(tetrafluoroethylene), (C₂F₄)_n which is used for everything from linings for pots and pans to gaskets that are inert to chemical reactions. Large amounts of fluorine are also consumed each year to make the freons (such as CCl₂F₂) used in refrigerators.

Chlorine (Cl₂) is a highly toxic gas with a pale yellow-green color. Chlorine is a very strong oxidizing agent, which is used commercially as a bleaching agent and as a disinfectant. It is strong enough to oxidize the dyes that give wood pulp its yellow or brown color, for example, thereby bleaching out this color, and strong enough to destroy bacteria and thereby act as a germicide. Large quantities of chlorine are used each year to make solvents such as carbon tetrachloride (CCl₄), chloroform (CHCl₃), dichloroethylene (C₂H₂Cl₂), and trichloroethylene (C₂HCl₃).

Bromine (Br₂) is a reddish-orange liquid with an unpleasant, choking odor. The name of the element, in fact, comes from the Greek stem bromos, “stench.” Bromine is used to prepare flame retardants, fire-extinguishing agents, sedatives, antiknock agents for gasoline, and insecticides.

Iodine is an intensely colored solid with an almost metallic luster. This solid is relatively volatile, and it sublimes when heated to form a violet-colored gas. Iodine has been used for many years as a disinfectant in “tincture of iodine.” Iodine compounds are used as catalysts, drugs, and dyes. Silver iodide (AgI) plays an important role in the photographic process and in attempts to make rain by seeding clouds. Iodide is also added to salt to protect against goiter, an iodine deficiency disease characterized by a swelling of the thyroid gland.

Some of the chemical and physical properties of the halogens are summarized in the table below. There is a regular increase in many of the properties of the halogens as we proceed down the column from fluorine to iodine, including the melting point, boiling point, intensity of the color of the halogen, the radius of the corresponding halide ion, and the density of the element. On the other hand, there is a regular decrease in the first ionization energy as we go down this column. As a result, there is a regular decrease in the oxidizing strength of the halogens from fluorine to iodine.

F2> Cl2> Br2> I2

oxidizing strength

This trend is mirrored by an increase in the reducing strength of the corresponding halides.

I–> Br–> Cl–> F–

reducing strength

Some Properties of F₂, Cl₂, Br₂, and I₂

		Melting Point (C)		Boiling Point (C)		Color		Natural Abundance (ppm)		1st Ionization Energy (kJ/mol)		Electron Affinity (kJ/mol)		Ionic Radius (nm)		Density (g/cm3)
F2		-218.6		-188.1		colorless		544		1680.6		322.6		0.133		1.513
Cl2		-101.0		-34.0		pale green		126		1255.7		348.5		0.184		1.655
Br2		-7.3		59.5		dark red-brown		2.5		1142.7		324.7		0.196		3.187
I2		113.6		185.2		very dark violet almost black		0.46		1008.7		295.5		0.220		3.960

Methods of Preparing the Halogens from their Halides

The halogens can be made by reacting a solution of the halide ion with any substance that is a stronger oxidizing agent. Iodine, for example, can be made by reacting the iodide ion with either bromine or chlorine.

2 I–(aq)

+

Br2(aq)

I2(aq)

+

2 Br–(aq)

Bromine was first prepared by A. J. Balard in 1826 by reacting bromide ions with a solution of Cl₂ dissolved in water.

2 Br–(aq)

+

Cl2(aq)

Br2(aq)

+

2 Cl–(aq)

To prepare Cl₂, we need a particularly strong oxidizing agent, such as manganese dioxide (MnO₂).

2 Cl–(aq)

+

MnO2(aq)

+

4 H+(aq)

Cl2(aq)

+

Mn2+(aq)

+

2 H2O(l)

The synthesis of fluorine escaped the efforts of chemists for almost 100 years. Part of the problem was finding an oxidizing agent strong enough to oxidize the F^– ion to F₂. The task of preparing fluorine was made even more difficult by the extraordinary toxicity of both F₂ and the hydrogen fluoride (HF) used to make it.

The best way of producing a strong reducing agent is to pass an electric current through a salt of the metal. Sodium, for example, can be prepared by the electrolysis of molten sodium chloride.

	electrolysis
2 NaCl(l)		2 Na(s)	+	Cl2(g)

In theory, the same process can be used to generate strong oxidizing agents, such as F₂.

Attempts to prepare fluorine by electrolysis, however, were initially unsuccessful. Humphry Davy, who prepared potassium, sodium, barium, strontium, calcium, and magnesium by electrolysis repeatedly tried to prepare F₂ by the electrolysis of fluorite (CaF₂), and succeeded only in ruining his health. Joseph Louis Gay-Lussac and Louis Jacques Thenard, who prepared elemental boron for the first time, also tried to prepare fluorine and suffered from very painful exposures to hydrogen fluoride. George and Thomas Knox were badly poisoned during their attempts to make fluorine, and both Paulin Louyet and Jerome Nickles died from fluorine poisoning.

Finally, in 1886 Henri Moissan successfully isolated F₂ gas from the electrolysis of a mixed salt of KF and HF and noted that crystals of silicon burst into flame when mixed with this gas. Electrolysis of KHF₂ is still used to prepare fluorine today, as shown in the figure below.

	electrolysis
2 KHF2(s)		H2(g)	+	F2(g)	+	2 KF(s)

Common Oxidation Numbers for the Halogens

Fluorine is the most electronegative element in the periodic table. As a result, it has an oxidation number of -1 in all its compounds. Because chlorine, bromine, and iodine are less electronegative, it is possible to prepare compounds in which these elements have oxidation numbers of +1, +3, +5, and +7, as shown in the table below.

Common Oxidation Numbers for the Halogens

Oxidation Number		Examples
-1		CaF2, HCl, NaBr, AgI
0		F2, Cl2, Br2, I2
+1		HClO, ClF
+3		HClO2, ClF3
+5		HClO3, BrF5, BrF6–, IF5
+7		HClO4, BrF6+, IF7

General Trends in Halogen Chemistry

There are several patterns in the chemistry of the halogens.

1. Neither double nor triple bonds are needed to explain the chemistry of the halogens.

2. The chemistry of fluorine is simplified by the fact it is the most electronegative element in the periodic table and by the fact that it has no d orbitals in its valence shell, so it can’t expand its valence shell.

3. Chlorine, bromine, and iodine have valence shell d orbitals and can expand their valence shells to hold as many as 14 valence electrons.

4. The chemistry of the halogens is dominated by oxidation-reduction reactions.

The Hydrogen Halides (HX)

The hydrogen halides are compounds that contain hydrogen attached to one of the halogens (HF, HCl, HBr, and HI). These compounds are all colorless gases, which are soluble in water. Up to 512 mL of HCl gas can dissolve in a single mL of water at 0^oC and 1 atm, for example. Each of the hydrogen halides ionizes to at least some extent when it dissolves in water.

	H2O
HCl(g)		H+(aq)	+	Cl–(aq)

Several of the hydrogen halides can be prepared directly from the elements. Mixtures of H₂ and Cl₂, for example, react with explosive violence in the presence of light to form HCl.

H2(g)

+

Cl2(g)

2 HCl(g)

Because chemists are usually more interested in aqueous solutions of these compounds than the pure gases, these compounds are usually synthesized in water. Aqueous solutions of the hydrogen halides are often called mineral acids because they are literally acids prepared from minerals. Hydrochloric acid is prepared by reacting table salt with sulfuric acid, for example, and hydrofluoric acid is prepared from fluorite and sulfuric acid.

2 NaCl(s)	+	H2SO4(aq)		2 HCl(aq)	+	Na2SO4(aq)
CaF2(s)	+	H2SO4(aq)		2 HF(aq)	+	CaSO4(aq)

These acids are purified by taking advantage of the ease with which HF and HCl gas boil out of these solutions. The gas given off when one of these solutions is heated is collected and then redissolved in water to give relatively pure samples of the mineral acid.

Preparation of halogens

Laboratory Preparation of Chlorine

Any of the methods listed below, can be used to prepare chlorine in the laboratory.

From concentrated hydrochloric acid by oxidation.

Chlorine can be prepared by removing the hydrogen from hydrochloric acid using an oxidizing agent. Any oxidising agent such as manganese dioxide, lead dioxide, trilead tetroxide, potassium permanganate or potassium dichromate can be used. Firstly, the oxidising agents are taken in the round bottomed flask. Concentrated hydrochloric acid is then added through a thistle funnel. This mixture is then heated. The oxygen of the oxidizing agents combines with the hydrogen of the hydrochloric acid leaving behind chlorine i.e. hydrogen is removed from hydrochloric acid. The metallic ions of the oxidising agents combine with part of chlorine to form the respective chlorides.

Remember:-

No heating is required in when potassium permanganate is used as an oxidizing agent in the above method of preparing chlorine.

• From sodium chloride, manganese dioxide and concentrated sulphuric acid.

Chlorine can be prepared by the action of hot concentrated sulphuric acid on a mixture of a chloride salt and an oxidising agent. Sodium chloride, being the cheapest and the most easily available chloride salt, is used along with manganese dioxide as the oxidising agent.

A mixture of almost equal quantity of sodium chloride and manganese dioxide is taken in a round bottomed flask. Concentrated sulphuric acid is then poured through the thistle funnel. The reaction takes place in two stages. In the first stage sulphuric acid reacts with the chloride to form hydrochloric acid. In the second stage the hydrochloric acid so formed combines with the oxidising agent to liberate chlorine.

(1)

(2) Manganese (II) chloride, formed during the reaction, reacts with sulphuric acid to form manganese (II) sulphate and hydrochloric acid as under:

From Bleaching Powder [Ca(OCl)Cl]

The apparatus is set up as shown in figure 14.5.

Bleaching powder is taken in a round bottom flask. Any dilute mineral acid is poured through the thistle funnel. Chlorine can be prepared by dropping any acid on bleaching powder.

The chorine produced from hydrochloric acid by the above method, is passed through two wash bottles. The first wash bottle contains water, to remove traces of hydrogen chloride gas from chlorine. The second wash bottle contains concentrated sulphuric acid to dry the gas.

Industrial Preparation of Chlorine

Chlorine is mostly obtained as a by-product during the manufacture of caustic soda, by the electrolysis of brine or molten sodium chloride. Hence chlorine is rather prepared by cheap methods. During this electrolysis, chlorine is liberated at the anode.

Remember:-

As chlorine is denser than air it is collected by the upward displacement of air.

The Interhalogen Compounds

Interhalogen compounds are formed by reactions between different halogens. All possible interhalogen compounds of the type XY are known. Bromine reacts with chlorine, for example, to give BrCl, which is a gas at room temperature.

Br2(l)

+

Cl2(g)

2 BrCl(g)

Interhalogen compounds with the general formulas XY₃, XY₅, and even XY₇ are formed when pairs of halogens react. Chlorine reacts with fluorine, for example, to form chlorine trifluoride.

Cl2(g)

+

3 F2(g)

2 ClF3(g)

These compounds are easiest to form when Y is fluorine. Iodine is the only halogen that forms an XY₇interhalogen compound, and it does so only with fluorine.

ClF₃ and BrF₅ are extremely reactive compounds. ClF₃ is so reactive that wood, asbestos, and even water spontaneously burn in its presence. These compounds are excellent fluorinating agents, which tend to react with each other to form positive ions such as ClF₂⁺ and BrF₄⁺ and negative ions such as IF₂^– and BrF₆^–.

2 BrF5(l)

[BrF4+][BrF6–](s)

Neutral Oxides of the Halogens

Under certain conditions, it is possible to isolate neutral oxides of the halogens, such as Cl₂O, Cl₂O₃, ClO₂, Cl₂O₄, Cl₂O₆, and Cl₂O₇. Cl₂O₇, for example, can be obtained by dehydrating perchloric acid, HClO₄. These oxides are notoriously unstable compounds that explode when subjected to either thermal or physical shock. Some are so unstable they detonate when warmed to temperatures above -40^oC.

Oxyacids of the Halogens and Their Salts

Chlorine reacts with the OH^– ion to form chloride ions and hypochlorite (OCl^–) ions.

Cl2(aq)

+

2 OH–(aq)

Cl–(aq)

+

OCl–(aq)

+

H2O(l)

This is a disproportionation reaction in which one-half of the chlorine atoms are oxidized to hypochlorite ions and the other half are reduced to chloride ions.

When the solution is hot, this reaction gives a mixture of the chloride and chlorate (ClO₃^–) ions.

3 Cl2(aq)

+

6 OH–(aq)

5 Cl–(aq)

+

ClO3–(aq)

+

3 H2O(l)

Under carefully controlled conditions, it is possible to convert a mixture of the chlorate and hypochlorite ions into a solution that contains the chlorite (ClO₂^–) ion.

ClO3–(aq)

+

ClO–(aq)

2 ClO2– (aq)

The last member of this class of compounds, the perchlorate ion (ClO₄^–), is made by electrolyzing solutions of the chlorate ion.

The names of the oxyanions of the halogens use the endings –ite and –ate to indicate low and high oxidation numbers and the prefixes hypo– and per– to indicate the very lowest and very highest oxidation numbers, as shown in the table below. Each of these ions can be converted into an oxyacid, which is named by replacing the –ite ending with –ous and the –ate ending with –ic.

Oxyanions and Oxyacids of Chlorine

		Oxyanions			Oxyacids
Oxidation State of the Chlorine		Compound		Name			Compound		Name
+1		ClO–		hypochlorite			HClO		hypochlorous acid
+3		ClO2–		chlorite			HOClO		chlorous acid
+5		ClO3–		chlorate			HOClO2		chloric acid
+7		ClO4–		perchlorate			HOClO3		perchloric acid

EVALUATION

1.Arrange   Flourine ,Chlorine,Bromine and Iodine inorder of  increasing oxidizing ability

2.Why is the  preparation of  Halogens carried out in the Fume cupboard?

3.Itemize using relevant chemical equation where necessary how you would test for chlorine in the laboratory.

4.Explain what happen when chlorine water is exposed to sunlight.

Notes on Chemistry SS2 – Edudelight.com

WEEK  7

NITROGEN

History

Nitrogen, which makes up about 78% of our atmosphere, is a colorless, odorless, tasteless and chemically unreactive gas at room temperature. It is named from the Greek nitron + genes for soda forming.  For many years during the 1500’s and 1600’s scientists hinted that there was another gas in the atmosphere besides carbon dioxide and oxygen. It was not until the 1700’s that scientists could prove there was in fact another gas that took up mass in the atmosphere of the Earth.

Daniel Rutherford, a Scottish Physician, discovered Nitrogen in 1772. But it was Lavoisier, who proved that nitrogen is an element and that it was not a supporter of combustion and respiration. He called it ‘azote’ meaning ‘no life’. The name ‘Nitrogen’ was given to it by Jean Antoine Chaptal (1756 – 1832), in the year 1790.

Nitrogen is found to have either 3 or 5 valence electrons and lies at the top of Group 15 on the periodic table. It can have either 3 or 5 valence electrons because it can bond in the outer 2p and 2s orbitals. Nitrogen is not reactive at standard temperature and pressure. Nitrogen is a colorless, and odorless gas that is usually found in its molecular form of (N₂). For the most part, Nitrogen is inert.

A Bohr diagram of the nitrogen atom.

Nitrogen is a non-metal element that occurs most abundantly in the atmosphere, nitrogen gas (N₂) comprises 78.1% of the volume of the Earth’s air. It only appears in 0.002% of the earth’s crust by mass. Compounds of nitrogen are found in foods, explosives, poisons, and fertilizers. Nitrogen makes up DNA in the form of nitrogenous bases as well as in neurotransmitters. It is one of the largest industrial gases, and is produced commercially as a gas and a liquid.

Table 1: General Properties of Nitrogen
Name and Symbol	Nitrogen, N
Category	non-metal
Atomic Weight	14.0067
Group	15
Electron Configuration	1s2 2s2 2p3
Valence Electrons	2, 5
Phase	Gas

Occurrence of Nitrogen

1. Nitrogen exists in the free state and in combined form in the atmosphere of the earth.
2. In the free state it occurs as N₂ and occupies about 80% by volume of air.
3. In the combined form, it is present as ammonia and all ammonium compounds, in nitric acid and in all nitrates, nitrites and in all living things in the form of proteins.

Isotopes

Nitrogen has two naturally occurring isotopes, nitrogen-14 and nitrogen-15, which can be separated with chemical exchanges or thermal diffusion. Nitrogen also has isotopes with 12, 13, 16, 17 masses, but they are radioactive.

• Nitrogen 14 is the most abundant form of nitrogen and makes up more than 99% of all nitrogen found on Earth. It is a stable compound and is non-radioactive. Nitrogen-14 has the most practical uses, and is found in agricultural practices, food preservation, biochemicals, and biomedical research. Nitrogen-14 is found in abundance in the atmosphere and among many living organisms. It has 5 valence electrons and is not a good electrical conductor.
• Nitrogen-15 is the other stable form of nitrogen. It is often used in medical research and preservation. The element is non-radioactive and therefore can also be sometimes used in agricultural practices. Nitrogen-15 is also used in brain research, specifically nuclear magnetic resonance spectroscopy (NMR), because unlike nitrogen-14 (nuclear spin of 1), it has a nuclear spin of 1/2 which has benefits when it comes to observing MRI research and NMR observations. Lastly, nitrogen-15 can be used as label or in some proteins in biology. Scientists mainly use this compound for research purposes and have not yet seen its full potential for uses in brain research.

Preparation of Nitrogen from Air

The chief source of free nitrogen is atmospheric air and nitrogen is usually prepared from it. Air free from dust, water vapour and carbon dioxide is compressed in a compression chamber for liquefaction.

Firstly, the pressure on the air is increased to about 200 atmospheres. It is then released through a spiral into a low-pressure area, where intense cooling of the air takes place.

Preparation of Nitrogen From Chemical Compounds

Ammonia and Ammonium Compounds

By treating excess ammonia with chlorine, ammonium chloride and nitrogen are formed.

Laboratory Preparation of Nitrogen

In the laboratory, nitrogen is prepared by heating a mixture of ammonium chloride and sodium nitrite and a small quantity of water. If ammonium nitrite is heated by itself, it decomposes to produce nitrogen gas. However, this reaction is very fast and may prove to be explosive.

Physical Properties of Nitrogen

• Color: Nitrogen is a colorless gas.
• Odor: It has no odor.
• Taste: It is a tasteless gas.
• It provides an inter atmosphere.
• Liquid nitrogen is colorless and odorless but resembles water in appearance.
• Nitrogen exists in two allotropic forms a and b, with a transition temperature of -237° C.
• Density of nitrogen is 1.25 x 10^-3 g.cm^-3 at 20°C
• Melting point of nitrogen is -210 °C
• Boiling point of nitrogen is -195.8 °C

Chemical Properties of Nitrogen

Combination with elements. Nitrogen does not easily combine with other elements under ordinary conditions. A molecule of nitrogen is diatomic. These two atoms have combined by mutually sharing three pairs of electrons.

• Nitrogen forms compounds through biological activity, at high or moderate temperature in the presence of catalysts.
• It forms nitric oxide and nitrogen dioxide with oxygen.
• Its reacts with hydrogen to give ammonia.
• It gives nitrogen sulfide when treated with sulfur.
• At very high temperatures nitrogen can combine with certain active metals, such as lithium, magnesium and titanium to form metal nitrides.
• It is not a supporter of combustion.
• Energy of first ionization: 1402 kJ.mol ^-1
• Energy of second ionization: 2856 kJ.mol ^-1
• Energy of third ionization: 4577 kJ.mol ^-1
• Electronegativity according to Pauling: 3.0

Uses of Nitrogen

1. Nitrogen is used in high temperature thermometers where mercury cannot be used.
2. This is because mercury boils at 356.7^oC and hence cannot be used in such thermometers.
3. A volume of nitrogen is enclosed in a vessel and introduced into the region of high temperature.
4. Depending upon the temperature, expansion of the nitrogen takes place.
5. Then applying the gas equation, the temperature is calculated

.The Nitrogen Cyclefor the gaseous element N_2(g)

• Nitrogen is an extremely important element for all plant or animal life! It is found in important molecules such as amino acids, which are combined to form proteins. Protein is used everywhere in living organisms from muscle structure in animals to enzymes in plants/animals.
• Nitrogen from the atmosphere:
  • Action of nitrifying bacteria, e.g. they function in the root nodules of certain plants like peas/beans (the legumes), can directly convert atmospheric nitrogen into nitrogen compounds in plants e.g. nitrogen => ammonia => nitrates which plants can absorb.
    • However, most plants can’t do this conversion from nitrogen => ammonia, though they can all absorb nitrates, so the ‘conversion’ or ‘fixing’ ability might be introduced into other plant species by genetic engineering.
  • The nitrogen from air is converted into ammonia in the chemical industry, and from this artificial fertilisers are manufactured to add to nutrient deficient soils. However, some of the fertiliser is washed out of the soil and can cause pollution.
  • The energy of lightning causes nitrogen and oxygen to combine and form nitrogen oxides which dissolve in rain that falls on the soil adding to its nitrogen content.
    • N_2(g) + O_2(g)==> 2NO_(g), then 
    • then 2NO_(g) + O_2(g)==> 2NO_2(g) 
    • NO_2(g) + water ==> nitrates_(aq) in rain/soil
    • Incidentally, reactions 1. and 2. can also happen in a car engine, and NO₂ is acidic and adds to the polluting acidity of rain as well as providing nutrients for plants!
• :
• :

• More ‘biological detail’ of the NITROGEN CYCLE with reference to the above diagram so you can show an understanding of how nitrogen is recycled.
  • a) Nitrogen gas in the air (78%, ~⁴/₅th) cannot be used directly by most plants and all animals.
    • No animals and only a few specialised plants can directly use the very unreactive nitrogen from air, but all plants nitrogen in some form to synthesise amino acids and proteins for growth and maintenance and for DNA in cell reproduction.
    • However, nitrogen can be changed into nitrogen compounds like nitrates which the plants can use.
    • Animals rely on plants or other animals in the food chain for their source of nitrogen compounds e.g. protein in grass, crops or other animals.
  • b) Action of Nitrogen-fixing bacteria living in root nodules of plants or in the soil, their function is to fix nitrogen gas from the atmosphere into a chemical form the plant can metabolise.
    • Leguminous plants like peas, lentils, clover and beans can absorb nitrogen from the air via their root nodules (swellings on the root surface) which contain enzymes capable of converting (‘fixing’) atmospheric nitrogen into soluble nitrate – a nutrient essential for amino acids, proteins and therefore plant growth.
      • Legumes and their root nodule bacteria are an example of mutualism (see section 3.19 b) because the plant root supplies the bacteria with carbohydrate food and minerals and the bacteria supplies the plant in the form of the nitrate ion.
      • The process of converting nitrogen in air into nitrogen compounds is sometimes called ‘nitrogen fixation‘.
  • c) The action of lightning can convert nitrogen gas into nitrates.
    • The very high electrical energy discharges from lightning activates nitrogen and oxygen molecules to react and form nitrogen oxides. These dissolve in rain to form nitrates which end up in the soil when rainwater trickles into the soil.
  • d) Action of decomposers: Their function is to break down dead animals and plants
    • Decomposers, e.g. various organisms like bacteria, fungi or worms can break down dead animals or plants. They break down proteins to amino acids.
  • e) Action of soil bacteria: Their function is to convert proteins and urea into ammonia or ammonium ions.
    • Decomposer bacteria in the soil can change proteins from dead plants/animals and urea in animal urine/droppings into ammonia/ammonium ion compounds.
    • d) plus e) is sometimes called putrefaction by putrefying bacteria.
  • f) Action of nitrifying bacteria: Their function is to convert this ammonia to nitrates – the process of nitrification
    • Nitrifying bacteria oxidise ammonia/ammonium ions from the decayed material to form nitrates, the nitrate ion can be absorbed by plants through their root systems.
  • g) Plants absorb nitrates from the soil.
    • Plants absorb nitrates (soluble in water) in the moisture that the roots absorb from the surrounding soil.
    • Plants can use the nitrate ion in forming amino acids from which the plant can make its proteins.
  • h) Nitrates are needed by plants to make proteins for growth.
    • Nitrates are an essential nutrient for plants to synthesis amino acids and hence proteins.
  • i) Nitrogen compounds pass along a food chain or web of food chains.
    • All food chains involve the passing of carbon compounds e.g. sugars, carbohydrates, fats and proteins up to the next trophic level i.e. the consecutive eating along a food chain (and waste produced on the way).
      • e.g. grass ==> cow ==> human
      • Plants make their own protein from nitrates, but animals must obtain it from plants or other animals. In fact the protein is broken down in digestion to amino acids and each animal makes its own proteins from these amino acid residues.
  • j) Action of denitrifying bacteria: Their function is to convert nitrates to nitrogen gas.
    • Particular bacterial organisms can remove the oxygen from nitrate compounds to form the element nitrogen gas.
    • These denitrifying bacteria live in anaerobic conditions like waterlogged soils and use the nitrate ion to respire.
    • This is the opposite function of the nitrogen-fixing bacteria .

EVALUATION

1.Describe the physical and chemical properties of Nitrogen.

2.With the aid of a well labelled diagram,describe the preparation of Nitrogen from air

3.Describe the Laboratory preparation of a named Hydride.

4.Give an example of a reaction in which ammonia behaves as a i.reducing agent  ii.base iii.precipitating agent

 WEEK 8

COMPOUNDS OF NITROGEN

Nitrogen has 5 electrons in its valence shell. It has a valency of 3 with respect to hydrogen and a valency up to 5 with respect to oxygen. So, it can combine with various elements to form many compounds.
The well known compounds of nitrogen are

1. Hydrazine
2. Hydrazoic acid
3. Hydroxylamine
4. Nitrous oxide
5. Nitric oxide
6. Di nitrogen trioxide
7. Nitrogen dioxide

Hydrazine

1. This compound can be regarded as being formed by replacing one H atom of NH₃ molecule by NH₂ group. This was first prepared by Curtius in 1889.
2. Free hydrazine is a colorless liquid, very hygroscopic and soluble in alcohol and water. It is not very stable.
3. On exposure to air, it readily absorbs carbon dioxide and moisture.
4. It dissolves sulfur, arsenic, selenium and phosphorus. Its boiling point is 113.5°C and melting point is 1.4°C.

Hydrazoic acid

1. This is also called as azoemide.
2. The formula of hydrazoic acid is N₃H.
3. This is a colorless volatile liquid. Its boiling point is 37°C.
4. It is extremely poisonous and dissociates into N₂ and H₂ with the evolution of a large amount of heat.

Hydroxylamine

• It is a hydroxy derivative of NH₃ and was discovered by Lossen in 1865.
• As a rule only the salts of hydroxylamine are prepared because the preparation of free hydroxylamine is difficult and not without hazards under some conditions.
• It is a needle-like white substance.
• Melting point is 33°C and boiling point is 58°C at 22 mm pressure.
• It decomposes at steam temperature.
• It is readily soluble in water but not much in alcohol.
• It can be crystallized out from its ethereal solution.
• Its solution in water is a weak base

• It explodes with halogen and also with permanganate and dichromate.
• It is very unstable and the solid decomposes slowly above 15°C.

Nitric Oxide

• This oxide is also called as nitrous oxide and is produced when a mixture of nitrogen and oxygen is passed through an electric ar

• It is a colorless gas, heavier than air and very sparingly soluble in water.
• Under pressure, it changes to a colorless liquid under -151°C.
• The gas is neutral to litmus solution.
• It combines with O₂ to form brown fumes and nitrogen peroxide(NO₂).
• It reacts with burning charcoal, S, P to give oxides.

Nitrous Oxide

• The formula of this oxide is N₂O.
• When this oxide is inhaled in small quantities, it produces hysterical laughter and for this reason this oxide is also called as laughing gas.
• It is prepared by heating a mixture of NaNO₃ and (NH₄)₂SO₄

• Nitrous oxide is a colorless gas having a faint, sweetish smell and a sweet taste.
• When inhaled for a long time, it produces insensibility.
• Large quantities may prove fatal.
• This is soluble in water and the solution has a sweetish taste.
• That is why it is collected over hot water.
• It is heavier than air. It can easily be liquified at 0°C and at a pressure of 30 atm.
• The liquid boils at -89.5°C.
• The gas is neutral to litmus.

Di nitrogen Trioxide

• The formula is N₂O₃.
• This oxide is also called nitrogen sesquioxide.
• It is also called nitrous anhydride. This means it is the anhydride of HNO₂, since it gives HNO₂ when heated with H₂O.

• Nitrogen trioxide is prepared by the reduction of nitric acid with arseniousoxid

• It is a red colored gas.
• On condensing, it gives a dark blue liquid.
• It is acidic in nature and reacts with sodium hydroxide to form sodium nitrite

The Synthesis of ammonia – The Haber Process of Nitrogen Fixation

• Ammonia gas is synthesised in the chemical industry by reacting nitrogen gas with hydrogen gas in what is known as the Haber-Bosch Process, named after two highly inventive and subsequently famous chemists.
  • The Haber synthesis of ammonia is important for agriculture because nitrogen is an important element for plant growth.
  • But, it is a very stable molecule and only a few plants like legumes (peas, beans etc.) can directly ‘fix’ nitrogen the from air and incorporate it into protein molecules.
  • The Haber synthesis allows the efficient mass production of ‘artificial fertilisers’.
• If the chemical feedstocks for the Haber Process are nitrogen and hydrogen, where do we get these materials from?
• The nitrogen was once obtained from the fractional distillation of liquified air (80% N₂).
  • the air is filtered to remove dust and then compressed under high pressure.
  • The filtered air is cooled and water condensed out, and then carbon dioxide freezes out at -78^oC.
  • The air further cooled to form liquid air (liquefaction) at around -200^oC.
  • The liquid air is fractionally distilled at low temperature to separate oxygen (used in welding, hospitals etc.), nitrogen (for making ammonia), Noble Gases e.g. argon for light bulbs, helium for balloons).
  • Oxygen and argon are very close in boiling point and initially come out in the same fraction so a further fractional distillation is needed to separate them.
  • However, nitrogen is also produced by ‘deoxygenating’ air by combustion with methane.
• The hydrogen is made by (i) reacting methane (natural gas) and water or (ii) from cracking hydrocarbons (both reactions are done at high temperature with a catalyst).
  • (i) methane + water (steam) ==> hydrogen + carbon monoxide
    • CH₄ + H₂O ==> 3H₂ + CO
    • The ‘deadly’ carbon monoxide can be reacted with water in a 2nd stage to make more hydrogen.
      • CO + H₂O ==> CO₂ + H₂
        • and the carbon dioxide is removed to give the desired hydrogen gas.
        • These are called ‘reforming’ reactions.
  • or (ii) from cracking an alkane hydrocarbon from crude oil
    • e.g. C₈H₁₈==> C₈H₁₆ + H₂
  • AND, but just in passing and convenient ….
  • Other uses of hydrogen
    • Hydrogen-oxygen fuel cells to make electricity on small-scale.
    • Hydrogenation vegetable oils to make margarine.
    • Reducing metal oxides to free the metal.
    • An atomic hydrogen-oxygen welding torch.
    • Inflating weather balloons.
• The balanced equation for the Haber Synthesis reversible reaction is …
  • N_2(g) + 3H_2(g) 2NH_3(g)   (plus 92 kJ of heat energy given out, exothermic reaction)
• .. which means an equilibrium will form, so there is no chance of 100% yield even if you use, as you actually do, the theoretical reactant ratio of nitrogen : hydrogen of 1 : 3 !
• In forming ammonia 92kJ of heat energy is given out (i.e. exothermic, 46kJ of heat released per mole of ammonia formed).
• Also, four moles of ‘reactant’ gas form two moles of ‘product’ gas, so there is a net decrease in gas molecules on forming ammonia.
•  
• So applying the equilibrium rules, the formation of ammonia should be favoured by  …
  • (a) Using high pressure because you are going from 4 to 2 gas molecules, so high pressure favours the forward reaction to give fewer gas molecules.
    • The high pressure also speeds up the reaction because it effectively increases the concentration of the gas molecules,
    • but, the higher pressure means more dangerous and more costly engineering, so a compromise needed.
    • Does this prediction match the graph?
  • (b) Carrying out the reaction at a low temperature, favouring the forward reaction,
    • because it is an exothermic reaction favoured by lowering the temperature,
    • but, this may produce too slow a rate of reaction, so a compromise needed.
    • Does this prediction match the graph?
  • Therefore, the idea is to use a set of optimum conditions to get the most efficient yield of ammonia and this involves getting a low % yield (e.g. 8% – 15% conversion) but fast.
  • Described below are the conditions to give the most economic production of ammonia.
  • These arguments make the point that the yield of an equilibrium reaction depends on the conditions used.
    • The word ‘yield‘ means how much product you get compared to the theoretical maximum possible if the forward reaction goes 100% that way.
    • .
• In industry pressures of 200 – 300 times normal atmospheric pressure are used in line with the theory (200-300 atm).
• Theoretically a low temperature would give a high yield of ammonia BUT …
  • Nitrogen is very stable molecule and not very reactive i.e. chemically inert, so the rate of reaction is too slow at low temperatures.
  • To speed up the reaction an iron catalyst is used as well as a higher temperature (e.g. 400-450^oC).
  • The higher temperature is an economic compromise, i.e. it is more economic to get a low yield fast, than a high yield slowly!
  • Note: a catalyst does NOT affect the yield of a reaction, i.e. the equilibrium position BUT you do get there faster!
• With reference to the HABER SYNTHESIS chemical plant DIAGRAM
  • Hydrogen and nitrogen gases are mixed in the ratio 3:1 (to fit in with the molecular equation mole ratio) and the gaseous mixture fed into the top of the reaction chamber.
  • The gases are pumped down through the reaction chamber filled with lots of beds (‘shelves’) coated in the iron catalyst.
  • The hydrogen and nitrogen gases react on the surface of the iron catalyst to form ammonia.
    • N_2(g) + 3H_2(g) 2NH_3(g)
    • The atom economy is 100%.
    • The initial yield is 6%-8%, but unreacted gases are recycled to raise this to nearer 100% eventually … read on …
  • At the end of the process, when the gases emerge from the bottom of the iron catalyst reaction chamber, the gas mixture is cooled under high pressure, when only the ammonia liquefies and is so can be removed, tapped off from the cooled compression chamber and stored in cylinders for use e.g. making fertilisers.
  • Because the reaction is reversible, not all the nitrogen and hydrogen are converted to ammonia.
  • Any unreacted nitrogen and hydrogen (NOT liquified), is recycled back through the reactor chamber, very little is wasted!
    • Nitrogen (-196^oC) and hydrogen (-252^oC) have much lower boiling points than ammonia (-33^oC) and stay as gases.
    • Boiling points increase with pressure, but these normal atmospheric pressure values offer a fair comparison and the higher the boiling point of the liquid, the higher condensation point of the gas.
    • The temperature in the lower chamber is never low enough to condense out the unreacted hydrogen or nitrogen so only the desired product, ammonia gas condenses out, then the liquid ammonia is drained off at the bottom of the.
    • Since the hydrogen and nitrogen are still gases above the liquid ammonia, they are easily pumped around and mixed with new hydrogen and nitrogen and hence recycled through the reactor.
    • This means non of the original hydrogen and nitrogen reactants is wasted, despite the reaction being an equilibrium.
      • In fact the yield of ammonia can be as little as 6% conversion, but FAST, and the other 94% of reactant gases is recycled FAST.
• To sum up: A low % yield of ammonia is produced quickly at moderately high temperatures and pressure in the presence of an iron catalyst, and is more economic than getting a higher % equilibrium yield of ammonia at a more costly high pressure and a slower lower temperature reaction.
  • Using an effective iron catalyst can the reduce the cost of manufacturing ammonia by increasing the rate of reaction (more efficient) and lowering the energy requirements if the process can be done at lower temperatures (activation energy reduced).
  • Increasing the rate of reaction saves time and operating at a lower temperature saves energy and therefore saves money.
  • However, catalysts can be very specialised and expensive to produce and they get contaminated (‘poisoned’) and become less efficient, in this case sulfur compounds contaminate the iron catalyst.
  • So the iron catalyst might have to be extracted and cleaned up, but if a true catalyst (and it is), this’ refurbishment’ should enable the iron catalyst to be reused.
    • Remember, theoretically catalysts take part in the reaction, but are not consumed in the reaction and can be reused over and over again.

• .

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The summary of the formation of ammonia

nitrogen + hydrogen ammonia

N_2(g) + 3H_2(g) 2NH_3(g)     (plus 92 kJ heat energy, exothermic)

Ammonia formation is favouredby …

Rule 1. Lowering the temperature, because it is an exothermic reaction, but this may make it too slow, compromise required.

Rule 2. Increasing pressure, because there is a reduction in the molecules of gas, but the higher the pressure, the more costly the engineering.

Rule 4. An iron catalyst speeds up the reaction, but has no effect on the % ammonia in the reacted mixture exiting the reactor chamber in the chemical plant of a Haber synthesis process.

The Uses of Ammonia

Ammonia is used to make ammonium salts (mostly artificial fertilisers) and is used in the chemical industry in the manufacture of explosives, nitric acid and nitrates, pharmaceutical products and plastics.

(a) Ammonia is used to manufacture nitric acid

• Ammonia is oxidised with oxygen from air using a hot platinum catalyst to form nitrogen monoxide and water.
  • 4NH_3(g) + 5O_2(g)==> 4NO_(g) + 6H₂O_(g)
• The gas is cooled and reacted with more oxygen to form nitrogen dioxide.
  • 2NO_(g) + O_2(g) ==>2NO_2(g)
• This is reacted with more oxygen and water to form nitric acid.
  • 4NO_2(g)+ O_2(g) + 2H₂O_(l)==> 4HNO_3(aq)
• Nitric acid is used to make nitro-aromatic compounds from which dyes are made.
• It is also used in the manufacture of artificial nitrogenous fertilisers (like ammonium nitrate, see below).

---

(b) Ammonia is used to manufacture ‘artificial’ nitrogenous fertilisers

• Ammonia is a pungent smelling alkaline gas that is very soluble in water.
• The gas or solution turns litmus or universal indicator blue because it is a soluble weak base or weak alkali and is neutralised by acids to form salts.

• Ammonia is a synthetic rich source of artificial nitrogenous fertilisers essential for increased growth of plants e.g. cereal crops.
• Ammonium salts are used as ‘artificial’ or ‘synthesised’ fertilisers i.e. nitrogenous fertilisers ‘man-made’ in a chemical works, and used as an alternative to natural manure or compost etc.
• Ammonia is a base, and fertiliser salts are made by neutralising ammonia solution with the appropriate acid.

• The resulting solution is heated, evaporating the water to crystallise the salt e.g. with correct equations, with and without state symbols …

(i) ammonia + hydrochloric acid ==> ammonium chloride

NH₃ + HCl ==> NH₄Cl

NH_3(aq) + HCl_(aq) ==> NH₄Cl_(aq) 

(ii) ammonia + sulphuric acid ==> ammonium sulphate

2NH₃ + H₂SO₄ ==> (NH₄)₂SO₄

2NH_3(aq) + H₂SO_4(aq)==> (NH₄)₂SO_4(aq) 

(iii) ammonia + nitric acid ==> ammonium nitrate

NH₃ + HNO₃ ==> NH₄NO₃

NH_3(aq) + HNO_3(aq)==> NH₄NO_3(aq) 

Reactions (ii) and (iii) are used in fertiliser production, as is reaction (iv)

ammonia + phosphoric acid ==> ammonium phosphate

EVALUATION

1.Describe the following oxides of Nitrogen.

N₂O_4,  NO₂ , N₂O₅  , NO

2.Which of the oxides of nitrogen is a neutral oxide?

a.N₂O₄  b.NO₂  C.N₂O₅   d.NO

3.Describe the Haber process

4.How is trioxonitrate v acid produced from ammonia?

5.Describe the brown ring experiment.

WEEK  9 

SULPHUR AND ITS COMPOUNDS

Sulfur belongs to the chalcogen family. Other members of the family are oxygen, selenium, tellurium, and polonium. These elements make up Group 16 (VIA) of the periodic table. The periodic table is a chart that shows how chemical elements are related to each other.

The term chalcogen comes from two Greek words meaning “ore forming.” An ore is a naturally occurring mineral used as a source for an element. Many ores are compounds of a metal and oxygen or a metal and sulfur. Compounds that contain two elements, one of which is sulfur, are called sulfides. For example, a beautiful gold-colored mineral is called pyrite, or “fool’s gold,” because it looks so much like real gold. Pyrite is iron sulfide (FeS₂ ).

Sulfur was known to ancient peoples. Its physical and chemical properties are very distinctive. It often occurs as a brilliant yellow powder. When it burns, it produces a clear blue flame and a very strong odor.

SYMBOL
S

ATOMIC NUMBER
16

ATOMIC MASS
32.064

FAMILY
Group 16 (VIA)
Chalcogen

PRONUNCIATION
SUL-fur

Sulfur, also spelled as sulphur, is a very important element in today’s world. Its most important use is in the manufacture of sulfuric acid (H ₂ SO ₄ ). There is more sulfuric acid made than any other chemical in the world. It has an enormous number of important uses.

Discovery and naming

Sulfur must have been well known to ancient peoples. They sometimes referred to it as brimstone. Sulfur sometimes occurs in bright yellow layers on the top of the earth. It has a sharp, offensive odor. When it burns, it gives off a strong, suffocating smell. The odor is like that produced when a match is struck.

The Bible mentions brimstone in a number of places. For example, Sodom and Gomorrah were two towns destroyed by God for the wicked ways of their citizens: “The Lord rained upon Sodom and upon Gomorrah brimstone and fire.”

But ancient people certainly did not think about sulfur the way modern chemists do. In fact, they used the word “element” to talk about anything that was basic. Ancient Greek philosophers, for example, thought that everything consisted of four elements: earth, fire, water, and air. Other philosophers thought there were only two elements: sulfur and mercury.

But early thinkers were often confused as to what they meant by the word “sulfur.” They often were talking about anything that burned and gave off large amounts of smoke. To them, “sulfur” was really a “burning substance.” It took centuries for scientists to identify sulfur as an element.

Physical properties

Sulfur exists in two allotropic forms. Allotropes are forms of an element with different physical and chemical properties. The two forms of sulfur are known as α-form and β-form (the Greek letters alpha and beta, respectively). Both allotropes are yellow, with the α-form a brighter yellow and the β-form a paler, whitish-yellow. The α-form changes to the β-form at about 94.5°C (202°F). The α-form can be melted at 112.8°C (235.0°F) if it is heated quickly. The β-form has a melting point of 119°C (246°F). The boiling point of the α-form is 444.6°C (832.3°F).

The two allotropes have densities of 2.06 grams per cubic centimeter (α-form) and 1.96 grams per cubic centimeter (β-form). Neither allotrope will dissolve in water. Both are soluble in other Liquids, such as benzene (C ₆ H ₆ ), carbon tetrachloride (CCl₄ ), and carbon disulfide (CS ₂ ).

Another allotrope of sulfur is formed when the element is melted. This allotrope has no crystalline shape. It looks like a dark brown, thick, melted plastic.

Chemical properties

Sulfur’s most prominent chemical property is that it burns. When it does so, it gives off a pale blue flame and sulfur dioxide (SO ₂ ) gas. Sulfur dioxide has a very obvious strong, choking odor.

Sulfur sometimes occurs in bright yellow layers on the top of the earth. It has a sharp, offensive odor.

Sulfur also combines with most other elements. Sometimes it combines with them easily at room temperature. In other cases, it must be heated. The reaction between magnesium and sulfur is typical. When the two elements are heated, they combine to form magnesium sulfide (MgS):

Sulfur also combines with hydrogen gas:

The compound formed in this reaction is hydrogen sulfide (H ₂ S). Hydrogen sulfide has one of the best known odors of all compounds. It smells like rotten eggs. Hydrogen sulfide is added to natural gas (methane) used in homes for cooking and heating. Methane is odorless. So the unique smell of hydrogen sulfide makes it easy to know when there is a methane leak.

Occurrence in nature

At one time, sulfur occurred in layers along the Earth’s surface. They were easy for humans to find and take. Deposits like these are more difficult to find today. One place they still occur is in the vicinity of volcanoes. Sulfur is released from volcanoes as a gas. When it reaches the cold air, it changes back to a solid. It forms beautiful yellow deposits along the edge of a volcano.

Large supplies of sulfur still occur underground. They are removed by the Frasch process (see accompanying sidebar).

Sulfur also occurs in a number of important minerals. Some examples are barite, or barium sulfate (BaSO₄ ); celestite, or strontium sulfate (SrSO₄ ); cinnabar, or mercury sulfide (HgS); galena, or lead sulfide (PbS); pyrites, or iron sulfide (FeS₂ ); sphalerite, or zinc sulfide (ZnS); and stibnite, or antimony sulfide (Sb ₂ S ₃ ).

The abundance of sulfur in the Earth’s crust is thought to be about 0.05 percent. It ranks about number 16 among the elements in terms of their abundance in the earth. It is more abundant than carbon, but less abundant than barium or strontium.

The largest producers of sulfur in the world are the United States, Canada, China, Russia, Mexico, and Japan. In 1996, the United States produced about 11,800,000 metric tons of sulfur. It is mined in 30 states, Puerto Rico, and the U.S. Virgin Islands.

Isotopes

There are four naturally occurring isotopes of sulfur: sulfur-32, sulfur-33, sulfur-34, and sulfur-36. Isotopes are two or more forms of an element. Isotopes differ from each other according to their mass number. The number written to the right of the element’s name is the mass number. The mass number represents the number of protons plus neutrons in the nucleus of an atom of the element. The number of protons determines the element, but the number of neutrons in the atom of any one element can vary. Each variation is an isotope.

Sulfur occurs in the vicinity of volcanoes.

Six radioactive isotopes of sulfur are known also. A radioactive isotope is one that breaks apart and gives off some form of radiation. Radioactive isotopes are produced when very small particles are fired at atoms. These particles stick in the atoms and make them radioactive.

One radioactive isotope of sulfur, sulfur-35, is used commercially. In medicine, the isotope is used to study the way fluids occur inside the body. It also has applications in research as a tracer. A tracer is a radioactive isotope whose presence in a system can easily be detected. The isotope is injected into the system at some point. Inside the system, the isotope gives off radiation. That radiation can be followed by means of detectors placed around the system.

As an example, a company that makes rubber tires might want to know what happens to the sulfur added to tires. Sulfur-35 is added to rubber along with non-radioactive sulfur. Researchers follow the radioactive isotope in the tires to see what happens to the sulfur when the tires are used.

The Frasch method of removing sulfur

T heFrasch method is one of the most famous mining systems ever invented. It was developed by German-American chemist Herman Frasch (1851-1914) in 1887.

The Frasch method is based on the low melting point of sulfur. The element melts at a temperature slightly higher than that of boiling water (100°C). Here is how the method works:

A set of three nested pipes (one inside each other) is sunk into the ground. The innermost pipe has a diameter of about an inch. The middle pipe has a diameter of about four inches. And the outer pipe has a diameter of about eight inches.

A stream of superheated water is injected into the outer pipe. Superheated water is water that is hotter than its boiling point, but that has not started to boil. Superheated water can be made by raising the pressure on the water. Its temperature can reach 160°C (320°F).

The superheated water passes down the outer pipe into the underground sulfur, causing it to melt. The molten (melted) sulfur forms a lake at the bottom of the pipe.

At the same time, a stream of hot air under pressure is forced down the innermost (one-inch) pipe. The hot air stirs up the molten sulfur and hot water at the bottom of the pipe. A foamy, soupy mixture of sulfur and water is formed. The mixture is forced upward through the middle pipe. When it reaches the surface, it is collected. The sulfur cools and separates from the water.

Similar applications of sulfur-35 involve studying sulfur in steel when it is made, seeing how sulfur affects the way engines operate, following what happens when proteins (which contain sulfur) are digested, and learning how drugs that contain sulfur are processed in the body.

METHOD
Superheated steam at 170oC under 16 atmp is pumped down the outer most pipe . Since melting point of sulphur is 113oC , it is melted and collects at the bottom. Hot compressed air is blown down the inner most pipe to produce foam of molten sulphur . This foam rises through the annular space between the inner most pipe and the next. The foam of sulphur is collected in wooden tubs. After few hours, sulphur is converted into solid blocks.

EVALUATION

1.Sulphur iv oxide bleaches by a.oxidation  b.reduction c.decomposition d.carboxylation

2.sulphur iv oxide is used for these except a.germicide and fungicide b.refrigerant  c.preszerving liquids like orange juice  d.used for restoring ozone layer.

3.sulphur reacts with soft rubber used to harden it is a. direct linkage b.polymerization c.crosss linkage  d.smoking

4.The bleaching  property of SO₂ is due to

a.Acidic nature

b.basic nature

c.reducing property

d.oxidizing property

5. What is the transition temperature for the two allotropes of sulphur?

 a.25⁰C   b.78⁰C   c.96⁰C   d.100 ⁰C

.Describe the following under Sulphur.

Physical properties, chemical properties ,  occurrence ,Extraction

WEEK  10

COMPOUNDS OF SULPHUR WITH HYDROGEN

Hydrosulphuric acid (H₂S). This substance is a gas having the composition expressed by the formula H₂S and is commonly called hydrogen sulphide. It is found in the vapors issuing from volcanoes, and in solution in the so-called sulphur waters of many springs. It is formed when organic matter containing sulphur undergoes decay, just as ammonia is formed under similar circumstances from nitrogenous matter.

Preparation.Hydrosulphuric acid is prepared in the laboratory by treating a sulphide with an acid. Iron sulphide (FeS) is usually employed:

FeS + 2HCl = FeCl₂ + H₂S.

A convenient apparatus is shown in Fig. 41. A few lumps of iron sulphide are placed in the bottle A, and dilute acid is added in small quantities at a time through the funnel tube B, the gas escaping through the tube C.

Explanation of the reaction. Iron sulphide is a salt of hydrosulphuric acid, and this reaction is therefore similar to the one which takes place when sulphuric acid acts upon a nitrate. In both cases a salt and an acid are brought together, and there is a tendency for the reaction to go on until a state of equilibrium is reached. This equilibrium is constantly disturbed by the escape of the gaseous acid set free, so that the reaction goes on until all of the original salt has been decomposed. The two reactions differ in that the first one is complete at ordinary temperatures, while in the case of sulphuric acid acting upon sodium nitrate, the reacting substances must be heated so as to secure a temperature at which nitric acid is a gas.

Physical properties.Hydrosulphuric acid is a colorless gas, having a weak, disagreeable taste and an exceedingly offensive odor. It is rather sparingly soluble in water at ordinary temperatures, about three volumes dissolving in one of water. In boiling water it is not soluble at all. In pure form it acts as a violent poison, and even when diluted largely with air produces headache, dizziness, and nausea. It is a little heavier than air, having a density of 1.18.

Chemical properties. The most important chemical properties of hydrosulphuric acid are the following:

1. Acid properties.Hydrosulphuric acid is a weak acid. In solution in water it turns blue litmus red and neutralizes bases, forming salts called sulphides.

2. Action on oxygen. The elements composing hydrosulphuric acid have each a strong affinity for oxygen, and are not held together very firmly. Consequently the gas burns readily in oxygen or the air, according to the equation

H₂S + 3O = H₂O + SO₂.

When there is not enough oxygen for both the sulphur and the hydrogen, the latter element combines with the oxygen and the sulphur is set free:

H₂S + O = H₂O + S.

[Pg 148]

3. Reducing action. Owing to the ease with which hydrosulphuric acid decomposes and the strong affinity of both sulphur and hydrogen for oxygen, the substance is a strong reducing agent, taking oxygen away from many substances which contain it.

4. Action on metals.Hydrosulphuric acid acts towards metals in a way very similar to water. Thus, when it is passed over heated iron in a tube, the reaction is represented by the equation

3Fe + 4H₂S = Fe₃S₄ + 8H.

Water in the form of steam, under similar circumstances, acts according to the equation

3Fe + 4H₂O = Fe₃O₄ + 8H. 

Salts of hydrosulphuric acid,—sulphides. The salts of hydrosulphuric acid, called sulphides, form an important class of salts. Many of them are found abundantly in nature, and some of them are important ores. They will be frequently mentioned in connection with the metals.

Most of the sulphides are insoluble in water, and some of them are insoluble in acids. Consequently, when hydrosulphuric acid is passed into a solution of a salt, it often happens that a sulphide is precipitated. With copper chloride the equation is

CuCl₂ + H₂S = CuS + 2HCl.

Because of the fact that some metals are precipitated in this way as sulphides while others are not, hydrosulphuric acid is extensively used in the separation of the metals in the laboratory.

Explanation of the reaction. When hydrosulphuric acid and copper chloride are brought together in solution, both copper and sulphur ions are present, and these will come to an equilibrium, as represented in the equation

Cu⁺ + S^–<–>CuS.

Since copper sulphide is almost insoluble in water, as soon as a very small quantity has formed the solution becomes supersaturated, and the excess keeps precipitating until nearly all the copper or sulphur ions have been removed from the solution. With some other ions, such as iron, the sulphide formed does not saturate the solution, and no precipitate results.

OXIDES OF SULPHUR

Sulphur forms two well-known compounds with oxygen: sulphur dioxide (SO₂), sometimes called sulphurous anhydride; and sulphur trioxide (SO₃), frequently called sulphuric anhydride.

Sulphur dioxide (SO₂). Sulphur dioxide occurs in nature in the gases issuing from volcanoes, and in solution in the water of many springs. It is likely to be found wherever sulphur compounds are undergoing oxidation.

Preparation. Three general ways may be mentioned for the preparation of sulphur dioxide:

1. By the combustion of sulphur. Sulphur dioxide is readily formed by the combustion of sulphur in oxygen or the air:

S + 2O = SO₂.

It is also formed when substances containing sulphur are burned:

ZnS + 3O = ZnO + SO₂.

2. By the reduction of sulphuric acid. When concentrated sulphuric acid is heated with certain metals, such as copper, part of the acid is changed into copper sulphate, and part is reduced to sulphurous acid. The latter then decomposes into sulphur dioxide and water, the complete equation being

Cu + 2H₂SO₄ = CuSO₄ + SO₂ + 2H₂O.

3. By the action of an acid on a sulphite.Sulphites are salts of sulphurous acid (H₂SO₃). When a sulphite is treated with an acid, sulphurous acid is set free, and being very unstable, decomposes into water and sulphur dioxide. These reactions are expressed in the equations

Na₂SO₃ + 2HCl = 2NaCl + H₂SO₃,

H₂SO₃ = H₂O + SO₂. 

Explanation of the reaction. In this case we have two reversible reactions depending on each other. In the first reaction,

(1) Na₂SO₃ + 2HCl <–> 2NaCl + H₂SO₃,

we should expect an equilibrium to result, for none of the four substances in the equation are insoluble or volatile when water is present to hold them in solution. But the quantity of the H₂SO₃ is constantly diminishing, owing to the fact that it decomposes, as represented in the equation

(2) H₂SO₃<–> H₂O + SO₂,

and the sulphur dioxide, being a gas, escapes. No equilibrium can therefore result, since the quantity of the sulphurous acid is constantly being diminished because of the escape of sulphur dioxide.

Physical properties. Sulphur dioxide is a colorless gas, which at ordinary temperatures is 2.2 times as heavy as air. It has a peculiar, irritating odor. The gas is very soluble in water, one volume of water dissolving eighty of the gas under standard conditions. It is easily condensed to a colorless liquid, and can be purchased in this condition stored in strong bottle..

Chemical properties. Sulphur dioxide has a marked tendency to combine with other substances, and is therefore  active substance chemically. It combines with oxygen gas, but not very easily. It can, however, take oxygen away from some other substances, and is therefore a good reducing agent. Its most marked chemical property is its ability to combine with water to form sulphurous acid (H₂SO₃).

Sulphurous acid (H₂SO₃). When sulphur dioxide dissolves in water it combines chemically with it to form sulphurous acid, an unstable substance having the formula H₃SO₃. It is impossible to prepare this acid in pure form, as it breaks down very easily into water and sulphur dioxide. The reaction is therefore reversible, and is expressed by the equation

H₂O + SO₂<–> H₂SO₃.

Solutions of the acid in water have a number of interesting properties.

1. Acid properties. The solution has all the properties typical of an acid. When neutralized by bases, sulphurous acid yields a series of salts called sulphites.

2. Reducing properties. Solutions of sulphurous acid act as good reducing agents. This is due to the fact that sulphurous acid has the power of taking up oxygen from the air, or from substances rich in oxygen, and is changed by this reaction into sulphuric acid:

H₂SO₃ + O = H₂SO₄,

H₂SO₃ + H₂O₂ = H₂S0₄ + H₂O.

3. Bleaching properties.Sulphurous acid has strong bleaching properties, acting upon many colored substances in such a way as to destroy their color. It is on this account used to bleach paper, straw goods, and even such foods as canned corn.

4. Antiseptic properties.Sulphurous acid has marked antiseptic properties, and on this account has the poweR of arresting fermentation. It is therefore used as a preservative.

Salts of sulphurous acid,—sulphites. The sulphites, like sulphurous acid, have the power of taking up oxygen very readily, and are good reducing agents. On account of this tendency, commercial sulphites are often contaminated with sulphates. A great deal of sodium sulphite is used in the bleaching industry, and as a reagent for softening paper pulp.

Sulphur trioxide (SO₃). When sulphur dioxide and oxygen are heated together at a rather high temperature, a small amount of sulphur trioxide (SO₃) is formed, but the reaction is slow and incomplete. If, however, the heating takes place in the presence of very fine platinum dust, the reaction is rapid and nearly complete.

Experimental preparation of sulphur trioxide. The experiment can be performed by the use of the apparatus shown in Fig. 43, the fine platinum being secured by moistening asbestos fiber with a solution of platinum chloride and igniting it in a flame. The fiber, covered with fine platinum, is placed in a tube of hard glass, which is then heated with a burner to about 350°, while sulphur dioxide and air are passed into the tube. Union takes place at once, and the strongly fuming sulphur trioxide escapes from the jet at the end of the tube, and may be condensed by surrounding the receiving tube with a freezing mixture.

Properties of sulphur trioxide. Sulphur trioxide is a colorless liquid, which solidifies at about 15° and boils at 46°. A trace of moisture causes it to solidify into a mass of silky white crystals, somewhat resembling asbestos fiber in appearance. In contact with the air it fumes strongly, and when thrown upon water it dissolves with a hissing sound and the liberation of a great deal of heat. The product of this reaction is sulphuric acid, so that sulphur trioxide is the anhydride of that acid:

SO₃ + H₂O = H₂SO₄. 

Catalysis. It has been found that many chemical reactions, such as the union of sulphur dioxide with oxygen, are much influenced by the presence of substances which do not themselves seem to take a part in the reaction, and are left apparently unchanged after it has ceased. These reactions go on very slowly under ordinary circumstances, but are greatly hastened by the presence of the foreign substance. Substances which hasten very slow reactions in this way are said to act as catalytic agents or catalyzers, and the action is called catalysis. Just how the action is brought about is not well understood.

DEFINITION: A catalyzer is a substance which changes the velocity of a reaction, but does not change its products.

Examples of Catalysis. We have already had several instances of such action. Oxygen and hydrogen combine with each other at ordinary temperatures in the presence of platinum powder, while if no catalytic agent is present they do not combine in appreciable quantities until a rather high temperature is reached. Potassium chlorate, when heated with manganese dioxide, gives up its oxygen at a much lower temperature than when heated alone. Hydrogen dioxide decomposes very rapidly when powdered manganese dioxide is sifted into its concentrated solution.[Pg 154]

On the other hand, the catalytic agent sometimes retards chemical action. For example, a solution of hydrogen dioxide decomposes more slowly when it contains a little phosphoric acid than when perfectly pure. For this reason commercial hydrogen dioxide always contains phosphoric acid.

Many reactions are brought about by the catalytic action of traces of water. For example, phosphorus will not burn in oxygen in the absence of all moisture. Hydrochloric acid will not unite with ammonia if the reagents are perfectly dry. It is probable that many of the chemical transformations in physiological processes, such as digestion, are assisted by certain substances acting as catalytic agents. The principle of catalysis is therefore very important.

Sulphuric acid (oil of vitriol) (H₂SO₄). Sulphuric acid is one of the most important of all manufactured chemicals. Not only is it one of the most common reagents in the laboratory, but enormous quantities of it are used in many of the industries, especially in the refining of petroleum, the manufacture of nitroglycerin, sodium carbonate, and fertilizers.

Manufacture of sulphuric acid. 

1. Contact process. The reactions taking place in this process are represented by the following equations:

SO₂ + O = SO₃,

SO₃ + H₂O = H₂SO₄. 

To bring about the first of these reactions rapidly, a catalyzer is employed, and the process is carried out in the following way: Large iron tubes are packed with some porous material, such as calcium and magnesium sulphates, which contains a suitable catalytic substance scattered through it. The catalyzers most used are platinum powder,[Pg 155] vanadium oxide, and iron oxide. Purified sulphur dioxide and air are passed through the tubes, which are kept at a temperature of about 350°. Sulphur trioxide is formed, and as it issues from the tube it is absorbed in water or dilute sulphuric acid. The process is continued until all the water in the absorbing vessel has been changed into sulphuric acid, so that a very concentrated acid is made in this way. An excess of the trioxide may dissolve in the strong sulphuric acid, forming what is known as fuming sulphuric acid.

2. Chamber process. The method of manufacture exclusively employed until recent years, and still in very extensive use, is much more complicated. The reactions are quite involved, but the conversion of water, sulphur dioxide, and oxygen into sulphuric acid is accomplished by the catalytic action of oxides of nitrogen. The reactions are brought about in large lead-lined chambers, into which oxides of nitrogen, sulphur dioxide, steam, and air are introduced in suitable proportions.

Reactions of the chamber process. In a very general way, the various reactions which take place in the lead chambers may be expressed in two equations. In the first reaction sulphur dioxide, nitrogen peroxide, steam, and oxygen unite, as shown in the equation

(1) 2SO₂ + 2NO₂ + H₂O + O = 2SO₂ (OH) (NO₂).

The product formed in this reaction is called nitrosulphuric acid or “chamber crystals.” It actually separates on the walls of the chambers when the process is not working properly. Under normal conditions, it is decomposed as fast as it is formed by the action of excess of steam, as shown in the equation

(2) 2SO₂ (OH) (NO₂) + H₂O + O = 2H₂SO₄ + 2NO₂.

The nitrogen dioxide formed in this reaction can now enter into combination with a new quantity of sulphur dioxide, steam, and oxygen, and the series of reactions go on indefinitely. Many other reactions occur, but these two illustrate the principle of the process.

The relation between sulphuric acid and nitrosulphuric acid can be seen by comparing their structural formulas:

O= -OH	O= -OH
S	S
O= -OH	O= -NO2

EVALUATION

1.Which of the following  is a colourless gas with a repulsive smell like that of a rotten egg?

a.H₂S   b.HCl  c.SO₂   d.Cl₂

2.Which of the following gases serves as both reducing and bleaching agents?

a.H₂S    b.SO₂  c.SO₂ d.HCl

3.The PH of the solution made by dissolving  hydrogen sulphide in water is

a.1   b.7   c.11   d.4              

4.Which of the following acts both as a reducing and as an oxidizing agent.

a.CO₂  b SO₂   c.H₂S   d.O₂

 _{1.How would you prepare a jar of hydrogen sulphide . Give the equation of reaction}

_{2.How are hydrogen tetraoxosulphate vi formed? Why are they known as acid salts?}