Equilibrium

Non reversible reactions

• Some chemical and physical reactions occur until one or all the reactants are used up

Example 1 Evaporation of water from a beaker (a physical process)

Reaction is complete when all the H2O molecules are in the gas phase

gaseous molecules from evaporation

water in liquid phase Evaporation

continues

Fig 1: Physical process going to completion

Non reversible chemical reactions

Hydrogen gas

Magnesium

Example 2 Magnesium in excess dilute acid solution

Eventually all the magnesium reacts and gas production stops

In examples 1 and 2 the forward reaction continues until the reactant(s) are used up this is written as:

Reactants Products

Fig 2: Chemical process going to completion

Reversible ReactionsIf we warm our beaker of water with glad wrap over the top we get a reversible reaction because the reaction is now in a closed system

Evaporation is occurring but so is the opposing process of condensation

Example 3

Gaseous water molecules from evaporationWater in liquid

phase

The forward reaction is from left to right

H2O (l) H2O (g)

The reverse reaction is

H2O (g) H2O (l)

The overall equilibrium reaction is written as

H2O (l) H2O (g)

This reaction has reached Equilibrium when evaporation and condensation occur at the same rate

Equilibrium

• When reversible reactions reach a point where the rate of forward reaction equals the rate of the backward reaction the system is said to have reached equilibrium

• The equilibrium is said to be dynamic which means the forward and backward reactions do not stop

Important Points so Far

1. Not all reactions are equilibrium reactions (ie some reactions are not reversible)

2. If an equilibrium reaction is to occur it must be in a closed system

3. When a reaction has reached dynamic equilibrium this means that the forward reaction is occurring at the same rate as the reverse reaction and this is written as:

reactants products

Chemical Equilibrium• Many chemical reactions behave as a dynamic equilibrium

• An Example is the Iron(lll) / thiocyanate equilibrium

• Fe3+ (aq) + SCN-(aq)

yellow colourless ion

In the reaction

A. The red colour forms quickly then remains constant

B. The system is closed as the ions Fe3+, SCN- and FeSCN2+ are all contained in the same aqueous solution

FeSCN 2+(aq)

bright red complex

Iron(lll) / thiocyanate equilibriumcontinued

• The reaction hasn’t gone to completion because it can be shown that their are unreacted Fe3+, SCN- ions in the red solution, this shows the reverse reaction is also occurring

FeSCN 2+(aq) Fe3+ (aq) + SCN-

(aq)

• The equilibrium for the reaction is written:

Fe3+ (aq) + SCN-(aq) FeSCN 2+

(aq)

• At equilibrium the rate at which the complex ions (FeSCN2+) form is equal to the rate at which the complex ions (FeSCN2+) break up

Part A• Fe3+ (Iron III) ions (yellow) &

SCN- (thiocyanate) ions (colourless)

• These two ions react to form a reddish complex ion:

Fe3+ + SCN- FeSCN2+

• Make some of this and place on 4 petrie dishes

Control

Add more Fe3+

Add more SCN-

Add F-

Why does a reaction stop?• Until now, you expect it to be

because one or both of the reactants have run out

• In this reaction, neither the Fe3+, nor the SCN- have run out, yet the reaction appears to stop

• Why?

Two processes occur• one to make FeSCN2+

• the other to decompose FeSCN2+

• The reaction stops when these two reactions balance out

Fe3+ + SCN- forming complex FeSCN2+

Fe3+ + SCN- decomposing FeSCN2+

• This is written as:

Fe3+ + SCN- FeSCN2+

Question - Which one’s correct?

Equilibrium occurs only when:

A. The reactants are used up

B. The concentration of reactants is equal to the concentration of products

C. All chemicals stop reacting

D. The products react together to form the reactants at the same rate as reactants form products

Ans - D

• Starter Questions

• What two conditions do you need for an equilibrium reaction to occur?

• The reaction must be a reversible reaction• The reaction must be in a closed system

• What is meant by the term equilibrium?• The forward reaction is occurring at the same rate as

the reverse reaction (rate forward = rate reverse)

• What factors can affect an equilibrium reaction?• Concentration of reactants or products• Temperature• Surface area• Pressure

Equilibrium constant

• Scientists wanted to put a number on the proportions of products and reactants when equilibrium is reached

• They did heaps of experiments and eventually they came upon a method for finding the equilibrium constant (K) for each particular reaction

Equilibrium ConstantFor a general reversible reaction:

aA + bB cC + dD‘a’ moles of ‘b’ moles of ‘c’ moles of ‘d’ moles of substance A substance B substance C substance D

The ratio of products to reactants is expressed as the equilibrium constant symbol K and is written as

remember as

Writing Equilibrium Equations

Writing the following Equilibrium expression for:

CO (g) + H2O (g) CO2 (g) + H2 (g)

Where: [B] is the concentration of B in molL-1 at equilibrium

[A] is the concentration of A in molL-1 at equilibrium

Equilibrium Constants

Equilibrium constants (K) are specific for a given reaction and temperature.

An equilibrium constant is useful for determining how far a reaction has proceeded.

If the equilibrium constant (K) is large

eg > 105 then at equilibrium the reaction has proceeded almost to completion.

Equilibrium Constants (continued)

• If the equilibrium constant (K) is small

eg < 10 -5 then at equilibrium only a small amount of reactant is used up and only a small amount of product has formed then the reaction has effectively not happened

Writing Equilibrium expressions:

Task: Write the equilibrium expression for

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

O]][H[CH]H][CO[

K24

32

Be very careful!

For equilibrium systems involving solid(s) and pure liquids

The solid(s) and liquids are not included in the equilibrium expression

Eg: write the equilibrium expression for

C (s) + H2O (g) CO (g) + H2 (g)

Be very careful!

O]H[][H [CO]

K 2

2

Writing more Equilibrium expressions:

Eg: write the equilibrium expression for

S(s) + O2 (g) SO2 (g)

][0SO

K2

2][

Be very careful!

N2(g) + 3H2(g) 2NH3(g)

Starter : write the equilibrium expression for:

322

23

]H][N[][NH

K

2N2(g) + 6H2(g) 4NH3(g)

write the equilibrium expression for:

62

22

43

]H[]N[][NH

K

Moving Equilibrium Systems

• Le Chatelier’s Principle: When change is applied to a system at equilibrium, the system responds so that the effects of the change are minimised

Example: H2(g) + Cl2(g) 2HCl(g)

What would be the effect on the amount of HCl formed if hydrogen gas was removed ?

The amount of HCl formed would decrease because the system would try to produce more H2 by shifting the equilibrium to the LHS.

In the above reaction what would be the effect on the amount of HCl formed if Cl2 gas was added?

The amount of HCl formed would increase, because the system will try to minimise the change and remove Cl2 by shifting the equilibrium to the RHS producing more HCl.

Turn to page 147 in your lab book – Equilibrium Systems

Lets see how changing concentration and temperature affect the position of the equilibrium

N2O4(l) 2NO2(g) ΔH = +54kJ(colourless) (brown)

Or…N2O4(l) + energy 2NO2(g)

•Forward reaction is endothermic•Backward reaction is exothermic•Cooling favours exothermic•Heating favours endothermic

•So heating this system will move the equilibrium to the right and the system will appear dark brown

•Cool it and move the equilibrium system to the left and the system will appear pale brown

The endothermic (+H) direction is the backward reaction, so an increase in temperature would move the equilibrium in the backward direction to take more energy in and therefore minimise the increase in temperature.

This would decrease the amount of NH3 produced.

Example

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

Consider another form of the N2O4/NO2 system

N2O4(g) 2NO2(g)

(colourless) (brown)Fewer molecules More molecules

Fewer moleculesMore molecules

• If the pressure is decreased, the system opposes this by increasing the number of molecules in the system– Forms more NO2(g)

– colour will go from light brown to a darker brown as new equilibrium established

Fewer moleculesMore molecules

Pressure

N2(g) + 3H2(g) 2NH3(g)

Now what would happen to the amount of NH3 produced if the pressure was decreased?

For the same reaction at equilibrium, an increase in the volume of the container, (or decrease in pressure), will shift the equilibrium to the side with more gas particles.

This means the concentration of NH3 would

decrease.

•Reactant concentration– Increase – reaction moves forwards (to

the right)– Decrease- reaction moves backwards (to

the left)

•Product concentration – Increase – reaction moves backwards– Decrease – reaction moves forwards

•Catalyst No change in equilibrium position, but equilibrium reached more quickly.

Exercise

1. For the equilibrium

N2O4(g) 2NO2(g) rH= -92 kJ mol -1

what would happen to the number of NO2 molecules if the following changes are made to the system at equilibrium

a) an increase in pressure

b) the concentration of N2O4 is increased

c) the pressure of the system is decreased

d) the temperature is decreased

e) a catalyst is added

PCl3 (g) + Cl2(g) PCl5 (g) H= +110 kJ mol -1

what would happen to the amount of PCl5 if the following changes are made to the system at equilibrium

a) an increase in pressure

b) the concentration of Cl2 is increased

c) the volume of the container is increased

d) the temperature is decreased

Starter Equilibrium Changes

Strategies to lower costs include

•Buy raw materials as cheaply as possible

•Use energy efficiently•Recycle unused reactants and by-products, if possible

•Heat exchangers may be used to remove energy from an exothermic process and deliver it to an endothermic process

•Catalysts assist in the rapid establishment of an equilibrium

•A severe shortage of nitrogen-based fertilisers at the start of the 20th century started hunt for a method of using nitrogen in the air–Nitrogen = 71% of air

•Fritz Haber developed the “Haber Process” used today to make ammonia

•Developed “Ecstasy”•Did fundamental

research into the Bunsen flame

•Developed a process for extracting gold from seawater

•Noble prize 1918 for Haber process

•Haber died in exile after standing up to Hitler’s anti-Semitic policies which he felt were morally wrong and were decimating German research institutes

The Equilibrium

N2(g) + 3H2(g) 2NH3(g) -ΔH

• To move the equilibrium to the right means a pressure drop and the production of heat– High temperatures favour the backward reaction but make the reaction rate faster

– High pressures favour the forward reaction and increase the yield of ammonia

% Yield of Ammonia Using The Haber Process

Pressure (atm)

T(0C) 10 50 100 300 1000200 51 74 82 90 98300 15 39 52 71 93400 4 15 25 47 80500 1 6 11 26 57600 0.5 2 5 14 13