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9.5 Industrial Chemistry: 2. Equilibrium reactions in industrial processes

Syllabus reference (October 2002 version)
2. Many industrial processes involve manipulation of equilibrium reactions	Students learn to: explain the effect of changing the following factors on identified equilibrium reactions pressure volume concentration temperature interpret the equilibrium constant expression (no units required) from the chemical equation of equilibrium reactions identify that temperature is the only factor that changes the value of the equilibrium constant (K) for a given equation	Students: identify data, plan and perform a first-hand investigation to model an equilibrium reaction choose equipment and perform a first-hand investigation to gather information and qualitatively analyse an equilibrium reaction process and present information from secondary sources to calculate K from equilibrium conditions

Extract from Chemistry Stage 6 Syllabus (Amended October 2002). © Board of Studies, NSW.
[Edit: 9 Jul 09]

Prior Learning: Preliminary modules 8.4.4. HSC Modules 9.3.2, 9.4.2

Background: Many reactions are reversible reactions, a forward and a reverse reaction will proceed at the same time. If left undisturbed in a closed system, these two reactions will eventually proceed at the same rate.  The reaction is said to have reached equilibrium.

At equilibrium, the concentration of all reactants and products stays constant. It seems as if nothing is happening, because the macroscopic properties, those you can see (e.g. colour and temperature), do not change. However, at the microscopic level there is continual change.

	When the concentration of products stays constant and the concentration of reactants stays constant, the system is at equilibrium.

At equilibrium, the rates of the forward and reverse reactions will be equal.

	Equilibrium occurs when the rate of the forward reaction is equal to the rate of the reverse reaction.

We can summarise the characteristics of a chemical system in equilibrium as the following:

1. It is a closed system - no matter or energy enters or leaves the system.
2. Macroscopic properties are constant e.g. state, colour, temperature, pressure and concentration.
3. Concentrations of reactants and products stay constant
4. Continual microscopic change occurs between reactants and products.
5. Rate of forward reaction = rate of backward reaction.

Note: constant does not mean equal.

At equilibrium, rates of reaction are equal, concentrations are constant - but not necessarily equal.

Whenever an equilibrium position is disturbed, the system tries to reduce the amount of change. This is called Le Chatelier's Principle.

You should be able to predict the effects of changes in concentration, adding or removing of chemicals, temperature changes and, when gases are involved, the effects of changes in pressure and volume.

 

identify data, plan and perform a first-hand investigation to model an  equilibrium reaction

• The background information given will help you to identify data to perform the investigation.

• An investigation you could perform is given below:

• Can you think of a way to extend this model to cater for changes in temperature for exothermic and endothermic reactions? Remember, when particles are heated they move faster

• Remember that models have limitations. What are the limitations of this model? Can you think of any features of equilibrium that cannot be covered by this model? For example, when people move faster they become tired and slow down, this doesn't happen for particles. Can you think of any other limitations

• Could the teacher act as a catalyst by encouraging students to get up and dance? (Remember that a catalyst changes the speed of reaction in both directions.)

  Video animations of equilibrium situations are located on the NSW Department of Education and Training's LMP site in the module 8.4 Water, Part 4 Salts in water, Aspects of equilibrium to see the importance of a closed system for equilibrium. Then click on NaCl to see an ionic equation for a reversible reaction. Then chemical reaction to see how concentrations of reactants and products become constant at equilibrium.

 

choose equipment and perform a first-hand investigation to gather information and qualitatively analyse an equilibrium reaction

• There are many possible experiments you could investigate. One example is the equilibrium between nitrogen dioxide and dinitrogen tetroxide.

Background information Nitrogen dioxide can be produced by the reaction of concentrated nitric acid on copper. Cu(s)   +   4HNO3(aq)    Cu(NO3)2(aq)   +   2NO2(g)   +   2H2O(l) Some nitrogen dioxide immediately changes to dinitrogen tetroxide and, if placed in a closed flask, an equilibrium is established between these two substances. 2NO2(g)    N2O4 (g) This reaction is exothermic, so you should be able to predict the effect of heating or cooling the equilibrium mixture.

• When choosing the equipment needed make sure safety issues are considered, such as protective goggles as an acid solution of extremely low pH is used.  Also work near a tap so that a supply of running water is available. If acid is splashed on the skin or eyes, you must wash it off immediately and continue washing for 15 minutes. To choose equipment:

identify and set up the most appropriate equipment or combination of equipment needed to undertake the investigation.	Choose appropriate equipment, e.g. a flask for which you have stoppers. It should be a convenient size to fit in your water bath. Plan a way to keep your water baths at constant temperature long enough to ensure that the temperature of gases in the flask is stable - for at least a few minutes.
carry out a risk assessment of intended experimental procedures and identify and address potential hazards, including disposal procedures.	NO2 and N2O4  are toxic gases so you should check the Chemical Safety in Schools Package for precautions necessary for their use. Concentrated HNO3 is highly corrosive, so check precautions for its use also.
identify technology that could be used and determine its suitability and effectiveness for its potential role in the procedure or investigations.	No hi-tech equipment is needed here, but consider the use of a thermometer and a fume cupboard.

• A method of performing the investigation would be:

1. A small piece of copper is placed in a reaction flask.

2. Add 10 mL of concentrated nitric acid (Caution!)

3. Seal the flask loosely with a stopper.

After the copper has dissolved, the reaction flask contains a mixture of NO₂ and N₂O₄.  The opening of another flask held upside down against the reaction flask opening can collect some of the gas mixture.

4. Show the effect of temperature on the equilibrium by putting the flask into water baths at different temperatures such as 0°C, room temperature, and 50°C.

Predicting results

Before you start the investigation, predict the effect different temperatures will have on equilibrium position and thus the colour of the mixture expected at each temperature. Write this up in the form of a table such as the one shown below.

Then carry out your investigation and record all observations.

Safety

When you carry out your investigation, make sure that you:

• identify and use safe work practices
• minimise hazards and wastage of resources
• dispose carefully and safely of any waste materials produced.

Analyse the results

Tabulating predictions and observations makes it easier to compare and analyse results.

Use this information to:

• identify and explain how data supports or refutes your initial predictions.
• identify trends, patterns and relationships as well as contradictions in data and information.
• make a generalisation regarding the effect of temperature changes on equilibrium. You can justify the generalisation by seeking evidence for similar relationships in other equilibrium reactions e.g. the Haber process.

• Another equilibrium reaction you could investigate is the equilibrium between chromate and dichromate ions in aqueous solution.

  If you start with about 20 mL of 0.1 molL^­1 potassium chromate solution in a container, it forms the following equilibrium:

  2CrO₄^2­(aq)   +   2H⁺(aq)    Cr₂O₇ ^2­(aq)   +   H₂O(l)
  Chromate ions are yellow, dichromate ions are orange.

  Adding a few drops of 2 molL^­1 sulfuric acid or 2 molL^­1 sodium hydroxide will affect the concentration of the hydrogen ions and thus affect the equilibrium.

  Choose your equipment, write up your method and analyse your results as described above.

  Do not forget the risk assessment - potassium chromate and dichromate can irritate the eyes, skin and respiratory system. They are also carcinogenic (cancer producing) and can sensitise skin, producing an allergic reaction.

• 

   

explain the effect of changing the following factors on identified equilibrium reactions
­ concentration
­ pressure
­ volume
­ temperature

• In the following examples, the equilibrium is moving to counter the change in the factors, as predicted by Le Chatelier's Principle.

  

Change	Effect
Increase concentration  of a reactant e.g.[CH4] increases.	Shifts right to use methane, increasing yield of CO and H2.
Decrease concentration of a product e.g. remove CO.	Shifts right to make more.
Increase pressure (decrease volume).	Shifts left (fewer particles) to drop pressure again.
Increase temperature.	Endothermic reaction ( is positive) so equilibrium shifts right to absorb added heat.

• For the effect of temperature changes, it may help if you re-write the equation as follows and just think of heat as another "substance" in the equation.

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

• Adding a chemical that reacts with a reactant or product removes the substance it reacts with and so affects equilibrium.

• CH₃COO^-(aq)   +   H₂O(l)    CH₃COOH(aq)  +   OH^-(aq)

• Adding hydrochloric acid provides hydrogen ions that react with the hydroxide ions (OH^-), removing them as water. This makes the equilibrium move to the right to make more OH^- .

• Adding sodium acetate will also affect the equilibrium. Sodium acetate will ionise, forming acetate ions (CH₃COO^-), so the equilibrium moves to the right to use up these extra ions.

process and present information from secondary sources to calculate K from equilibrium conditions

• For any equilibrium reaction we can calculate a constant, called the equilibrium constant (K).

      We can write a general equilibrium equation as:

      Where a, b, c, d are the number of moles of substances A, B, C and D. Then:

where [ ] means the concentration (in mol L^-1), at equilibrium, of each substance.

We can start with A and B, or C and D, and any concentration we like of these reactants and products, and this relationship will still hold. K will not change as long as we do not change the temperature.

The system is at equilibrium in diagrams 4, 5 and 6.

If we start with N₂O₄, at the same temperature, we can show equilibrium being reached as:

The system is at equilibrium in diagrams 10, 11 and 12.

  Calculations of K

The simplest type of question gives you the concentrations of reactants and products at equilibrium, as follows:

An equilibrium mixture for the reaction,

H₂(g)  +   I₂(g)    2HI(g)

contains the following concentrations of gases:

[H₂] = 2.9070 x 10^­3 molL^-1,  [I₂] = 1.7069 x 10^­3 molL^-1, 
[HI] = 16.482 x 10^­3 molL^-1

1. a) Write the equilibrium expression for this equation.

b) Determine the value of the equilibrium constant for the reaction at this temperature.

Click here to see the answer.

A slightly harder problem might ask you to work out the concentrations at equilibrium so you can calculate K.

Consider the   H₂(g)   +   I₂(g)    2HI(g)   equilibrium.

2. 5 mol of HI is placed in an empty 1.0 L container and allowed to come to equilibrium. At equilibrium there is 1 mol of HI present. Calculate the value of K.

Click here to see the answer.

You may also be required to find the concentrations at equilibrium from a graph.

2 mol each of hydrogen and 3 mol of iodine gas were placed in a 1 L container and kept at a constant temperature. The graph below shows changes that occurred in the concentration of hydrogen and hydrogen iodide as equilibrium was reached.

H₂(g)  +   I₂(g)    2HI(g) 

3.a) What do graphs A and B show?

b) When does equilibrium occur?

c) Complete the following reaction table:

d) Calculate K.

Click here to see the answer.

interpret the equilibrium constant expression (no units required) from the chemical equation of equilibrium reactions

• If K is large (e.g. >10³) the equilibrium lies to the right so the concentration of products is high. This means the reaction goes almost to completion.

• If K is small (e.g. <10^­3) the equilibrium lies to the left so the concentration of reactants is high.  This means there is very little reaction.

• The size of K values can indicate the relative strength of acids or the solubility of substances.
  
  For example, the following table shows K values for 3 acids, at the same temperature.

acetic acid	phosphoric acid	boric acid
1.8 x 105	7.1 x 103	7.3 x 1016

• From this information we can tell that the strongest of these three acids is phosphoric acid as it has the largest K value. This means that, at the same temperature, more phosphoric acid will ionise than either acetic or boric acid.

identify that temperature is the only factor that changes the value of the equilibrium constant (K) for a given equation

• The value of K for a given equation is not affected by:
  ­ changes in concentration or pressure
  ­ the addition of a catalyst. Adding a catalyst only reduces the time taken to reach equilibrium.

• The value of K for a given equation is affected by temperature. When temperature changes, the effect on K depends on whether the reaction is endothermic or exothermic.

• For example, the Haber process for producing ammonia is exothermic, and the equilibrium reaction can be written as:

  N₂(g) + 3H₂(g) 2NH₃(g) + heat

  K for this given equation =

  Heating for the reaction vessel pushes the equilibrium left to absorb the extra heat. This will decrease [NH₃] and increase [N₂] and [H₂]. These changes will make K smaller.

• K decreases when exothermic reactions are heated and endothermic reactions are cooled.

• K increases when exothermic reactions are cooled and endothermic reactions are heated.