Chapter 7 Chemical Equilibrium

Question 1

What is Equilibrium ?

Answer 1

•When confined to a closed system, reversible reactions can reach a point where they are said to be in
equilibrium.

•For a chemical reaction to be at equilibrium, two equal but opposing reactions are proceeding at the same rate.

Question 2

How does a reaction reach Equilibrium ?

Answer 2

• The rate of the forward reaction must equal the rate of the reverse reaction.
• Consequently the effects of reaction 1 and reaction 2 cancel each other out.
• So at equilibrium the concentration of both reaction remains constant over time.
• This constancy of concentrations for all the reactants and products is the hallmark of a chemical system that has reached Equilibrium.

Question 3

Why is a chemical system at equilibrium said to be dynamic as opposed to static ?

Answer 3

This is because it is still active as both the forward and reverse reactions despite the static appearance of a system at equilibrium.

Question 4

What are some points to remember about the equilibrium constant K?

Answer 4

• Only gases and aqueous species appear in K. Solids and liquids have a fixed concentration and so they are not included in the expression for K.
• While K has a constant value for all conditions of concentration and pressure its value does change if temperature changes.

Question 5

What does the different sizes of K indicate?

Answer 5

• large values of K imply the equilibrium favours products
• small values of K imply equilibrium favours reactants
• values of K Close to 1 imply significant concentrations of both reactants and products are present at equilibrium.

Question 6

Define Impose Changes.

Answer 6

It is the ability of a chemical equilibrium system to return to a state of equilibrium after alterations.

Question 7

List the imposed changes that will affect the equilibrium:

Answer 7

• Changing the Concentration of any one species in the system by selectively adding it to the system or removing it from the system.
• Changing the Total Pressure in a Gaseous equilibrium by reducing its volume.
• Changing the Temperature of the system by adding or removing Heat from the system.

Question 8

What happens when the concentration in an equilibrium system is altered by a change imposed on the system?

Answer 8

• A new equilibrium will form in such a way that partially counteracts this imposed change.
• In doing this a system will consume some of the added reagent or replace some of the consumed reagent.

Question 9

What is happen during these changes (system being altered by a an imposed change)?

Answer 9

-The system is temporarily out of equilibrium and the rate of the forward and reverse reactions won’t be equal.

Question 10

When does the system partially minimise the original imposed change in concentration?

Answer 10

-At some point the system returns to equilibrium (when the rates of forward and reverse reactions are again equal) but when it does, the concentrations and amounts of all the reagents will have altered in such a way that has partially minimised the original imposed change in concentration.

Question 11

Explain the Concentration Time Graph.

Answer 11

• A concentration time graph shows how an imposed change results in a new equilibrium.
• In this example the Dichromate/Chromate system is at equilibrium up until time (T1).
• At T1 a small amount of concentrated NaOH(aq) is added to the system (the imposed change).
• This imposed change means the (OH-) has been selectively raised.
• The system is now temporarily out of equilibrium as the increased (OH-) means the forward reaction rate, increases while the reverse reaction rate, remains unchanged (remember rate rises with concentration).
• This temporarily favours the formation of products and consumption of reactants.
• Consequently all concentrations change during the interval T1 to T2 until a new equilibrium is established (at T2) as predicted by Le Chatelier’s principle.

Question 12

How does an equilibrium system involving gases achieve an equilibrium?

Answer 12

It must be in a sealed container.

Question 13

What happens if the pressure in the system is altered?

Answer 13

The equilibrium will change in a way that minimises the pressure change.

Question 14

In relation to Le Chateler’s principle explain what happens when the pressure of an equilibrium system is altered.

Answer 14

• Increasing the pressure by reducing volume will favour the side of the equilibrium reaction with fewer moles of gas. Fewer moles of gas mean less pressure within the system thus partially counteracting the imposed change.
• Decreasing the pressure by increasing volume will favour the side of the equilibrium reaction with the greater moles of gas. More moles of gas means more pressure thus partially counteracting the imposed change.

Question 15

List situations in which the pressure has no effect in the equilibrium system.

Answer 15

• If both sides of the equilibrium reaction have the same number of moles of gas, then changing the pressure has no effect on the equilibrium position.
• Increasing pressure by adding an inert gas has no effect on the equilibrium position.
• Also, changing the pressure by adding or removing one of the gaseous reactants or products is equivalent to selectively altering the concentration of that substance only.

Question 16

What is an inert gas?

Answer 16

It is one that does not react with the substances in the equilibrium system.

Question 17

In relation to Le Chateler’s principle explain what happens when the temperature of an equilibrium system is altered.

Answer 17

• Raised: A new equilibrium is established favouring the endothermic process. By doing this, the system converts some of the added heat to chemical potential energy (bond energy), thus causing its temperature to fall, ie partially counteracting the imposed change.
• Lowered: A new equilibrium is established favouring the exothermic process. By doing this, the system converts some of the chemical potential energy (bond energy) to heat, thus causing its temperature to rise, ie partially counteracting the imposed change.

Question 18

Define Heat.

Answer 18

Energy that naturally flows from a hot system (high temp) to a cooler one.

Question 19

Define Chemical Potential Energy.

Answer 19

Energy stored by the chemical bonds in a substance.

Question 20

Define Enthalpy.

Answer 20

A measure of the total energy in a chemical system. It includes potential energy (bond energy) and particle kinetic energy.

Question 21

What happens in an exothermic reaction?

Answer 21

In an exothermic reaction temperature rises as some chemical potential energy (bond energy) is converted into particle kinetic energy.

• As a result, heat inevitably flows out of the hot reaction mixture and into the cooler surroundings.
• The formation of heat in an exothermic reaction and its eventual loss to the surroundings means the enthalpy of the reacting system must decrease and hence -ve.

Question 22

What happens in an endothermic reaction?

Answer 22

• In an endothermic reaction temperature falls as particle kinetic energy is converted to chemical potential energy.
• As the reaction mixture becomes cool, heat then flows into the mixture from the warmer surroundings and so its enthalpy increases, ie. +ve.

Question 23

What does a Catalyst do?

Answer 23

The presence of a catalyst in an equilibrium system has no effect on the equilibrium position; however, the catalyst will speed up the rate of attainment of equilibrium.

Question 24

Why doesn’t a Catalyst affect an equilibrium system?

Answer 24

This happens as the catalyst increases the rate of both the forwards and reverse reactions equally.

Question 25

List some of the materials produced from ammonia.

Answer 25

• Fertilisers - like NH4NO3, (NH4)2SO4, (NH4)2HPO4 and urea (NH2)2CO
• Polymers - such as nylon and acrylic plastics.
• Explosives - like trinitrotoluene (TNT), nitroglycerine, and NH4NO3
• Cleaning agents for removing dirt and grease.

Question 26

How is Ammonia manufactured?

Answer 26

By the Haber process which utilises the equilibrium reaction of nitrogen with hydrogen:

N2(g) + 3 H2(g) → 2 NH3(g) (ΔH = −92.4 kJ·mol−1)

Question 27

What is the issue of the system at a room temperature?

Answer 27

-While this reaction has a reasonable equilibrium yield of ammonia at room temperature and pressure, the rate of attainment of equilibrium is extremely low.

Question 28

What is the issue with increasing the temperature of the system?

Answer 28

-Higher temperatures will increase the rate of reaction; however, this also reduces the equilibrium yield of NH3.

Question 29

What is the issue with increasing the pressure of the system?

Answer 29

• Higher pressure favours both a high rate of reaction and a high equilibrium yield of NH3.
• However, using high pressure significantly increases construction and operating cost of the ammonia plant.
• For this reason excessively high pressures are uneconomic.

Question 30

What has to be done to overcome all these issues?

Answer 30

As a consequence of these conflicting problems a compromise in conditions must be used in order to produce ammonia economically.

Question 31

How is an economic production of ammonia achieved?

Answer 31

It is achieved by choosing reaction conditions that maximise the rate of attainment of equilibrium while still allowing a satisfactory equilibrium yield of ammonia without incurring excessive cost.

Question 32

How is the Haber process carried out today?

Answer 32

• It is carried out using a 3:1 molar mixture of H2(g) and N2(g) at a temperature of around 350-550*C and a pressure of 15-35 MPa.
• A catalyst consisting of iron-iron oxide fused with MgO, Al2O3 and SiO2 is also used.
• With these conditions a reasonable rate can be maintained with a yield of around 15-30% ammonia (depending upon the actual conditions used.)