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Year 11 - Chemical Equilibria

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Chemical Equilibria - Introduction

Many chemical reactions are not as simple as reactants coming together to form the desired products. Many reactions involve the products reacting together and breaking down back to reactants.

Reactions which do so are called reversible reactions, i.e. the reaction proceeds in both the forwards and backwards direction, or are said to be in equilibrium.

Hence the use of an equilibrium arrow, , in chemical equations instead of the normal equation arrow.

Many industrial reactions in particular are equilibria. For example, in the Contact process of making sulphuric acid,

2SO2(g) + O2(g) 2SO3(g)
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Chemical Equilibria - Changing Conditions

How far the reaction proceeds towards the products is known as the position of equilibrium or equilibrium position.

The equilibrium position for a particular reaction can be moved left or right by changing the physical properties of the reaction mixture, by changing the temperature, pressure or concentrations of components of the mixture.

N.B.: The absence or presence of a catalyst has no effect whatsoever on the equilibrium position of a reaction. A catalyst only speeds up a reaction.

(1) Change in concentration :

When the concentration of a component in a reaction mixture is increased the position of equilibrium moves to reduce the increase, by moving to the opposite side of the equation.

For example, in making ethyl ethanoate from ethanol and ethanoic acid, the reaction equation is,


If the concentration of ethanol, CH3CH2OH, is increased the equilibrium position moves to the right. If the concentration of ethyl ethanoate is increased the equilibrium position moves to the left.

The opposite holds if the concentration of a reaction component is decreased. So, in the above equation, if the concentration of ethanoic acid is decreased the position of equilibrium will move to the left.

(2) Change in pressure :

When the pressure of a reaction mixture is increased, the equilibrium position will move to decrease the pressure. It will do this by moving to the side of the reaction that has the fewest gaseous molecules.

If the pressure is decreased, the position of equilibrium will move to the side with the greatest number of gaseous molecules.

For example, in the formation of sulphur trioxide in the Contact process, the equation is,

2SO2(g) + O2(g) 2SO3(g)

If the pressure is increased the equilibrium position moves to the right, as there are 3 gas molecules on the left and only 2 gas molecules on the right.

If the pressure is decreased, the equilibrium position will move to the left, i.e. the side with the most gaseous molecules.

(3) Change in temperature :

When the temperature of a reaction is increased the position of equilibrium will move to reduce the temperature by moving to the side that is endothermic, i.e. the direction of the reaction that takes in energy.

When the temperature of a reaction is decreased the position of equilibrium will move to increase the temperature by moving to the side that is exothermic, i.e. the direction of the reaction that gives out energy.

For example, in the production of sulphur trioxide,

2SO2(g) + O2(g) 2SO3(g)   D H=-288 kJ mol-1

So, the forwards reaction as written above, i.e. the reaction going from left to right, is exothermic and the backwards reaction is endothermic.

If the temperature of the above equilibrium is increased the equilibrium positive moves to the left. If the temperature is decreased the position of equilibrium moves to the right.

N.B.: If the temperature of a reaction is decreased the overall rate of the reaction will also decrease, i.e. the products will form more slowly, whether or not the forwards reaction is exothermic or endothermic.

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Chemical Equilibria - The Haber Process

The Haber process is the industrial production of ammonia, NH3, from hydrogen and nitrogen.

The hydrogen gas required for the reaction is obtained by cracking methane with super-heated steam (i.e. steam at a high pressure and temperature),

CH4(g) + H2O(g) CO(g) + 3H2(g)

The nitrogen gas for this reaction is obtained from air - either by removing the oxygen from it, by burning the carbon monoxide produced above in air; or, more simply, by liquefying air at below -200 oC, allowing it to warm up and collecting the nitrogen gas as it boils off at about -196 oC. This is fractional distillation of liquid air.

The hydrogen and nitrogen are then combined together at about 400 oC and 250 atm pressure in a large tower filled with small pieces of iron, which act as a catalyst.

3H2(g) + N2(g) 2NH3(g)   D H=-92 kJ mol-1

The overall reaction is exothermic and so the temperature should be kept reasonably low to encourage product formation. About 400 oC is chosen to enable a good rate of production, even if it gives a poorer conversion.

The very high pressure is used to drive the reaction towards the product side of the equation, as there are 2 moles of gas on the right hand side and 4 moles of gas on the left.

Even with the adverse pressure conditions used, only about 15% of the hydrogen and nitrogen mixture is converted into ammonia. After the actual reaction the mixture of hydrogen, nitrogen and ammonia is cooled down to about -35 oC, which liquefies the ammonia in the mixture, allowing it to be piped away. The remaining hydrogen and nitrogen are then recycled back into the reaction tower with the catalyst and fresh reactant gases.

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Chemical Equilibria - Ammonium Salts

(1) Heating ammonium salts :

Ammonium compounds are easily decomposed by simply heating them in a test tube with the hot Bunsen burner flame.

They release ammonia gas, for which a positive test result is when damp red litmus paper turns blue. The gas also smells very strongly.

e.g. NH4Cl(s) -----> NH3(g) + HCl(g)
(2) Reaction with hydroxide compounds :

In a very similar reaction to the above process, when ammonium compounds are heated together with a hydroxide, e.g. sodium hydroxide, they decompose to give ammonia gas, water vapour and a salt.

e.g. NH4Cl(s) + NaOH(s) -----> NH3(g) + H2O(g) + NaCl(s)
(3) Tests for ammonia gas :

There are two main ways of testing for ammonia gas :

A piece of damp red litmus paper will turn blue. This is because ammonia is a basic gas.

When concentrated hydrochloric acid is placed near ammonia gas fumes, a large amount of white smoke is given off. This is ammonium chloride, produced when the acidic and basic gases react together.

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Chemical Equilibria - Production of Nitric Acid


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Chemical Equilibria - Nitrate Salts

(1) Heating nitrates :

When solid metal nitrate salts are subjected to strong heating they decompose. The products of this decomposition vary with the metal nitrate used.

Group I metals form stable compounds since they are at the top of the reactivity series. Their nitrates only decompose to give oxygen gas and a metal nitrite salt (see below).

Group II and Transition metals form less stable compounds and their nitrates are decomposed to give oxygen gas a metal oxide and nitrogen dioxide gas (see below).

The table below gives a summary of the effects of heating metal nitrates as well as two examples of balanced chemical equations -

Metals Examples Equation example
Group I Li, Na, K, Rb, Cs 2NaNO3(s) 2NaNO2(s) + O2(g)
Group II and Transition Mg, Ca, Fe, Cu, Pb 2Mg(NO3)2(s) 2MgO(s) + 4NO2(g) + O2(g)

(2) Test for the nitrate ion :

When a small amount of nitrate salt is dissolved in aqueous sodium hydroxide, some powdered aluminium (or Devarda's alloy which is an alloy of aluminium (45%), zinc(5%) and copper(50%)) added and the whole mixture gently heated, ammonia gas is produced. This ammonia gas can be detected for in the normal manner.

The equation for the reaction is very complex,

3NO3- + 8Al + 5-OH + 18H2O 3NH3(g) + 8[Al(OH)4]-

(source : Vogel's Qualitative Inorganic Analysis 6th edt., pg 183)

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written by Dr Richard Clarkson : © Saturday, 1 November 1997

updated : Tuesday, 29th July, 2008

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