DP Chemistry: Equilibrium
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Equilibrium

7.1 Equilibrium (4.5 hours)

Pause for thought

Once you have explained the concept of dynamic equilibrium (as opposed to static equilibrium) it is worth challenging your students by asking them to suggest how it could be shown experimentally that either chemical or physical equilibrium is dynamic and not static. There are several ways in which it could be done but one of the most obvious is to use a radioactive isotope. Although there is no need to mention examples of specific isotopes in Topic 2.1:The nuclear atom the dangers posed by 131-iodine in nuclear accidents is a good example of utilization and other specific isotopes are mentioned in several of the options.

Imagine pouring a saturated solution of potassium iodide (made with normal non-radioactive iodide ions) onto some solid potassium iodide containing 131-iodide ions. The mixture is then stirred and left for several hours. If the equilibrium was static then once the solid had been filtered from the solution the solution should not show any radioactivity as none of the iodide ions would have exchanged. However, if the iodide ions in solution are in dynamic equilibrium with the iodide ions in the solid then over time the solution will become increasingly more radioactive as more and more 131-iodide ions are incorporated into the solution.



You could then utilise this explanation to explain why normal iodine is kept as a safeguard at nuclear reactors. This is because if there is an explosion or melt-down one of the radioactive products is 131-iodine. If it escapes, which did happen for real at the Fukushima nuclear reactor in Japan in 2011 (see above), then flooding the body with normal iodine helps to prevent, or at least minimise, 131-iodine being incorporated into humans via the hormone thyroxine.

Students should be able to write equilibrium expressions and have some idea about what the size of the equilibrium constant tells them about the position of equilibrium. It is therefore reasonable to ask whether equilibrium constants have units. Consider the Contact process.

2SO2(g) + O2(g) ⇌ 2SO3(g)

Students should be able to deduce that:

Since they are told that [ ] means 'concentration' then in terms of units Kc is equal to concentration squared divided by concentration cubed which simplifies to one over concentration, or concentration-1. The units of the equilibrium constant, Kc, will therefore be mol-1 dm3. Using similar logic the units for the equilibrium constant for the Haber process will be concentration-2, or mol-2 dm6. In fact until 2007 the units of equilibrium constants was a specific assessment statement on the IB Chemistry Diploma programme and questions involving the units were often asked in the exams. The problem is that equilibrium constants do not in fact have units! The reasons is that the square brackets, [ ], actually mean activity rather than concentration and activity has no units. For this reason the values given for all equilibrium constants in examinations (and for Kw in the IB Data booklet) on the current syllabus do not have units.[1]

This can cause difficulties when comparing values of Kc for different reactions but not really at IB level. For example, the solubility product, Ksp, for silver chloride, AgCl, is 1.8 x 10-10 whereas the solubility product for silver chromate, Ag2CrO4, is 1.3 x 10-12. Although both are very 'insoluble' it would appear that silver chloride is more soluble as it has a higher Ksp value. However the concentration of silver ions in solution will be greater for silver chromate as the expression is [Ag+(aq)]2[CrO42-(aq)], i.e. concentration cubed, compared to the concentration squared expression, [Ag+(aq)][Cl-(aq)] for silver chloride.

Nature of science

The use of isotopic labelling to define equilibrium provides good evidence for a scientific theory. The term dynamic equilibrium is used in other contexts so is an example of the use of a common language across different disciplines although the chemistry definition may not always be born in mind.

Learning outcomes

After studying this topic students should be able to:

Understand

  • In a closed system, equilibrium is reached when the rate of the forward reaction is equal to the rate of the reverse reaction.
  • The equilibrium law describes how the equilibrium constant, Kc relates to the concentrations of the products and the reactants from which its value can be determined for a particular chemical reaction. The extent of a reaction at equilibrium is indicated by the magnitude of the equilibrium constant whose value is also dependent upon temperature.
  • The position of the equilibrium changes with changes in concentration, pressure, and temperature.
  • The relative amount of products and reactants present during a reaction at a particular point in time are given by the reaction quotient, Q. Q is the equilibrium expression using non-equilibrium concentrations.
  • Adding a catalyst has no effect on the position of equilibrium or the equilibrium constant.

Apply their knowledge to:

  • Characterise chemical and physical systems in a state of equilibrium.
  • Deduce the equilibrium constant expression, Kc, from an equation for a homogeneous reaction.
  • Determine the relationship between different equilibrium constants for the same reaction at the same temperature.
  • Apply Le Chatelier’s principle to predict the qualitative effects of changes of temperature, pressure and concentration on the position of equilibrium and on the value of the equilibrium constant.

Clarification notes

Both physical and chemical systems should be covered.
The relationship between Kc values for reactions that are multiples or inverses of one another should be covered.
No specific details of any industrial process are required.

International-mindedness

The Haber process has huge global significance as it has revolutionized world food production due to artificial fertilizers. However, because the oxidation of ammonia forms nitric acid, the precursor for many explosives, it has also had a large impact on weaponry in many world conflicts.

Teaching tips

To illustrate static equilibrium I often balance a ruler on a pencil. Once equilibrium has been reached there is no more movement until a force is added to one of the sides. In fact, although demonstrating dynamic equilibrium is more spectacular, it is still not obvious that the reactions are still continuing once the position of equilibrium has been reached.

You will need to explain the concept of closed and open systems and the fact that at equilibrium the concentrations of reactants and products are constant because the rate of the forward reaction is equal to the rate of the reverse reaction.

Suitable experiments for demonstrating dynamic equilibrium are the formation of iron(III) thiocyanate, [FeSCN]2+, a gas syringe full of nitrogen (IV) oxide or by adding ammonia and/or hydrochloric acid to aqueous copper(II) ions. Better still is to get your students to do the experiments themselves - the iron(III) thiocyanate formation is the easiest and safest.

You should also discuss dynamic physical equilibrium. Some teachers use the analogy of pouring water into a tub with a hole in the bottom. Equilibrium will be reached when the rate of water entering the tub is the same as the rate of water leaving through the hole. It is also worth asking students to explain what constitutes a good 'drying day' for washed clothes and why wet clothes do not dry in a closed space.

I tend to put much more emphasis on getting students to be able to write correct equilibrium expressions than on quoting Le Chatelier's principle. The principle is useful for predicting what will happen when a change is made to an equilibrium system but the equilibrium expression can be used to actually explain it.

For example, in the expression for the Contact process above adding more SO2(g) looks as if it will lower the value of Kc, but Kc is a constant so to keep it constant the concentration of SO3(g) must increase.

For temperature heat can be thought of as a reactant for endothermic reactions, so putting in more heat will move the position of equilibrium to the product side. Similarly, for exothermic reactions think of heat as a product so taking heat away (i.e. lowering the temperature) will cause more product to form.

Kc is of course only a constant at a stated temperature. By applying the equilibrium law it is easy to see that Kc must increase at higher temperatures for endothermic reactions and decrease at higher temperatures for exothermic reactions.

It is also easy for students to deduce how the reaction quotient, Q, will change as the system proceeds towards equilibrium provided they know the initial concentrations of all the reactants (and products, if any). You will also need to give students practice at manipulating equilibrium constants if the reactions are written differently.

Catalysts work by providing an alternative pathway with a lower activation energy. They affect the rate of forward and reverse reactions equally so do not affect the position of equilibrium. However equilibrium will be reached quicker so they tend to be used in industry for exothermic processes where higher temperatures result in a lower yield.

Study guide

Pages 54 & 55

Questions

For ten 'quiz' multiple choice questions with the answers explained see MC test: Equilibrium.

For short-answer questions which can be set as an assignment for a test, homework or given for self study together with model answers see Equilibrium questions.

Vocabulary list

Dynamic
Static
Steady state
Closed system
Macroscopic properties
Homogeneous equilibria
Heterogeneous equilibria
Le Chatelier's principle
Equilibrium law
Equilibrium constant, Kc
Reaction quotient, Q

IM, TOK, Utilization etc.

See separate page which covers all of Topic 7

Practical work

Le Chatelier's principle

Teaching slides

Teachers may wish to share these slides with students for learning or for reviewing key concepts.

  

Other resources

1. A video produced by the University of Surrey for revision purposes. It has quite a good illustration of the liquid water - water vapour equilibrium in open and closed systems.

Dynamic equilibrium (1)

2. A good video to illustrate Le Chatelier's principle. It shows how colourless SCN-(aq) ions and a very pale yellow solution of Fe3+(aq) ions react to form the red complex [Fe(H2O)5SCN]2+ ion using petri dishes and the effect of changing concentrations on the position of equilibrium.

Le Chatelier's principle (1)  

3. A University of Surrey video which applies Le Chatelier's principle to the Haber process.

Le Chatelier's principle (2)  

4. A good example of both concentration effect and le Chatelier's principle using a gas syringe filled with a mixture of NO2(g) and N2O4(g).

Le Chatelier's principle (3)  

5. A really good and informative article on the Fritz Haber story which you could give your students to read. Also see the information given in 'Incorporating IM, TOK, 'Utilization etc. for Topics 7 & 17.

Footnotes

  1. ^ Actually in the current version of the IB data booklet (4th. edition dated JIB Docs (2) Teamary 2017) in Section 2 the ionic product for water, Kw is given units whereas the solubility product constants, Ksp given in Section 32 do not have units! 
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