Year 12 Chemistry: Chapter 16 ~ Controlling the Yield of Reactions
Chapter 16.1 A Problem for Industry: Incomplete Reactions
Ammonia is prepared from nitrogen and hydrogen gas.
N2(g) + 3H2(g) ® 2NH3(g)
Suppose you mixed 1 mol of nitrogen and 3 mol of hydrogen gas in a sealed container. You might expect from the equation that 2 mol of ammonia would be produced.
The reaction seems to ‘stop’ when much less than 2 mol of ammonia is present. The stage when the quantities of reactants and products in the reaction remain unchanged is called chemical equilibrium.
The presence of large amounts of unreacted starting materials in reaction mixtures is wasteful and costly. The profitability of an industry is dependent upon the reaction yield – the extent of conversion of reactants to products.
To maximise the yield from such reactions, the following questions are vital:
J Why do some reactions reach equilibrium?
J Hoe can the amount of product from a reaction that reaches equilibrium be increased?
Why are some reactions incomplete?
Reversible reactions
Some physical and chemical changes can be reversed.
J An ice block in a drink melts, forming water. If the drink is placed in a freezer the water freezes.
H2O(s) H2O(l)
J Water melting is presented by the forward equation:
H2O(s) ® H2O(l)
J Water freezing is presented by the reverse equation:
H2O(l) ® H2O(s)
A double arrow ( ) is used to indicate that this is a reversible reaction.
J Nickel-metal hydride (NiMH) batteries (used in digital cameras and calculators). A chemical reaction provides electrical energy when the batteries are in use. During recharging the reverse reaction takes place, regenerating chemical that had been consumed.
NiOOH(s) + MH(s) Ni(OH)2(s) + M(s)
M represents metals (V, Ti, Cr, Co)
J Heating blue-coloured hydrated copper(II) sulfate (CuSO4.5H2O) in a test tube results in a white solid forming (CuSO4). When water is added and mixed with the white powder, the blue compound is re-formed.
The reaction between nitrogen and hydrogen to form ammonia, can also be a reversible reaction, since ammonia decomposes into nitrogen and hydrogen when it is heated.
Equilibrium explained
Chemist have shown that in these equilibrium reactions the forward and reverse reactions occur simultaneously.
N2(g) + 3H2(g) 2NH3(g)
J As the forward reaction proceeds, the concentrations of nitrogen and hydrogen decrease, so the rate of the ammonia production decreases.
J At the same time as ammonia is being formed, some ammonia molecules react to re-form nitrogen and hydrogen. The rate of the reverse reaction increases as the concentration of ammonia increases.
J Eventually the forward and back reactions proceed at the same rate.
When this situation is reached ammonia is formed at exactly the same rate as it is breaking down. The concentration of ammonia, nitrogen and hydrogen will then remain constant.
Equilibrium is a dynamic state, since the forward and back reactions have not ceased. Instead they occur simultaneously at the same rate. During dynamic equilibrium:
J The amount and concentrations of chemical substances remain constant
J The total gas pressure is constant (if gases are involved)
J The temperature is constant
J The reaction is ‘incomplete’
How far so equilibrium reactions go?
Both hydrogen chloride (HCl) and ethanoic (CH3COOH) react with water in a hydrolysis reaction to form ions according to the equations:
HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq)
CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq)
Where water and hydrochloric acid nearly react completely before reaching equilibrium, whereas only about 1% of ethanoic acid reacts with water.
Solutions of hydrogen chloride and ethanoic acid will conduct electricity because they contain ions. The solution formed when Hydrogen chloride (hydrochloric acid) is a much better conductor than the ethanoic acid solution.
The difference in conductivity arises because these reactions occur to remarkably different extents. At equilibrium in a 0.10 M solution, at 25°C, almost all the Hydrogen chloride molecules are ionised before reaching equilibrium, whereas only about 1% of ethanoic acid reacts with water.
*****Different reactions proceed to different extents*****
Questions: 1
Year 12 Chemistry: Chapter 16.2 Equilibrium
16.2 The Equilibrium Law
N2(g) + 3H2(g) 2NH3(g)
Stick in table 16.1 pg 267
There is no real pattern in the above table, apart from the last column. It shows the fraction __[NH3]2_ give almost constant value for each mixture.
[N2][H2]3
This fraction is known as the concentration fraction or the reaction quotient for the equilibrium
K = __[NH3]2_
[N2][H2]3
where K is known as the equilibrium constant. The equilibrium constant for the above at 400°C is 0.052 M-2.
Consider the reaction
N2O4(g) 2NO2(g)
Stick in table 16.2 pg 267
The concentration fraction or reaction quotient that gives a constant value at equilibrium is [NO2]2
[N2O4]
The equilibrium constant K, for this reaction has a value of 4.5 M at 80°C.
Chemists have found that:
· Different chemical reactions have different values of K;
· For a particular reaction, K is constant for all equilibrium mixtures at a fixed temperature.
WHERE K IS CONSTANT = EQUILIBRIUM LAW
It is important that the equation be specified when an equilibrium constant is quoted.
Questions: 2,3
What does an equilibrium constant tell us?
The equilibrium constant, K, gives an indication of the extent of reaction.
For values of K that are:
· Between 10-4 and 104, there will be significant amounts of reactants and products present at equilibrium.
· Very large (>104) the equilibrium mixture consists mostly of products, with relative small amounts of reactants.
HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq); K = 107 at 25°C
· Very small (<10-4) the equilibrium mixture consists mostly of reactants with relative small amounts of products.
CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq); K = 1.8 x 10-5 at 25°C
Effect of temperature on equilibrium
It has been shown that the value of the equilibrium constant, K, for a particular reaction depends only upon temperature. It is not affected by actions as addition of reactants or products, changes in pressure, or the use of catalysts.
The effect of a change in temperature on an equilibrium is found to depend upon whether the reaction is exothermic or endothermic. As temperature increases:
If the temperature of the equilibrium mixture decreases the opposite effect will occur.
Since the value of K depends on temperature, it is important to specify the temperature at which an equilibrium constant has been measured.
QUESTIONS: 4, 5, 6
Calculations using equilibrium constants
Example: Calculate the value of the equilibrium constant for the equation:
H2(g) + I2(g) 2HI(g)
at 400°C, if a 2.00 L vessel contains an equilibrium mixture of 0.0860 mol of H2, 0.124 mol of I2 and 0.715 mol of HI.
[H2] =
[I2] =
[HI] =
K =
Example: The equilibrium constant for the reaction by the equation:
N2O4(g) 2NO2(g)
is 4.5M at 80°C. In an equilibrium mixture at this temperature, what is the concentration of NO2 if the concentration of N2O4 is 0.0012 M?
K=
Therefore:
The square root
The concentration of NO2 in the equilibrium mixture is
Example: At a particular temperature 0.0500 mol of SO2, 0.0100 mol of O2 and 0.1500 mol of SO3 were mixed in a 2.00 L vessel and allowed to reach equilibrium according to the equation:
2SO2(g) + O2(g) 2SO3(g)
Analysis showed that 0.1400 mol of SO3 was present in the gas mixture at equilibrium. Calculate the value of the equilibrium constant at this temperature.
Amount of SO3 decreased therefore a net backward reaction
SO3 reacted = 0.1500 – 0.1400 =
SO2 / O2 / SO3So at equilibrium:
n(SO2) =
n(O2) =
n(SO3) =
[SO2] =
[O2] =
[SO3] =
K =
The equilibrium constant has a value of
Example: the equilibrium constant for the reaction N2O4(g) 2NO2(g) is 4.5 M at 80°C. A gas mixture in a 2.0 L vessel at 80°C contained, 0.20 mol of N2O4 and 0.30 mol of NO2. Decide if the reaction is at equilibrium, and it not, predict the direction the reaction will shift to reach equilibrium:
[N2O4] =
[NO2] =
The reaction quotient = ______
The reaction quotient doesn’t equal K, it is not at equilibrium. You need a net forward reaction. NO2 must increase and N2O4 must decrease.
Questions: 6, 7, 8.
16.3 Changing the Position of Equilibrium of a Reaction
Chemists can maximise the yield of a particular product by careful control of the conditions.
The equilibrium position of a reaction may be changed by:
J Adding or removing a reactant or product
J Changing the pressure by changing the volume (equilibria involving gases)
J Dilution (equilibria in solution)
J Changing the temperature
Adding extra reactant or product
N2(g) + 3H2(g) ® 2NH3(g)
If extra nitrogen gas were added to this reaction without changing the volume or temperature within the vessel, the mixture would momentarily not be at equilibrium. As it adjust back to equilibrium:
J The increased concentration of nitrogen gas causes the rate of the forward reaction to increase and more ammonia is formed;
J As the concentration of ammonia increases, the rate of the back reaction to re-form N2 and H2 increases;
J Ultimately the rates of the forward and back reaction become equal again and a new equilibrium is established.
Note that when equilibrium is re-established the concentrations of all substances have changed.
It is sometimes possible to arrange to remove a reactant or product from an equilibrium mixture. As you would anticipate, this has the opposite effect on an equilibrium to addition of the substance
· Addition of a reactant leads to the formation of more products (a net forward reaction);
· Addition of a product leads to the formation of more reactants (a net back reaction).
A net reaction that counteracts the effect of the change occurs:
Le Chatelier’s Principle
If an equilibrium system is subjected to a change, the system will adjust itself to partially oppose the effect of the change.
Changing the pressure
By changing the volume
The pressure of gases in an equilibrium mixture can be changed by increasing or decreasing the volume of the container whilst the temperature is kept constant.
During sulfuric acid manufacture, one step involves sulfur dioxide reacting with oxygen to form sulfur trioxide gas:
2SO2(g) + O2(g) ® 2SO3(g)
3 particles 2 particles
In this equilibrium the forward reaction involves a reduction in the number of particles of gas, causing a reduction in pressure. The back reaction involves an increase in gas particles, causing as increase in pressure.
REMEMBER Le Chatelier’s Principle, an equilibrium will respond to an increase in pressure by adjusting to reduce pressure. A net forward reaction will occur, increasing the amount of sulfur trioxide present at equilibrium.
In general the effect of a change of pressure, by changing the container volume depends on the relative number of gas particles on both sides of the equation.
Addition of an inert gas
Despite the increase in pressure due to the addition of a non-reacting gas, the concentrations of the individual chemicals involved in the equilibrium are not affected by the presence of the extra gas. The system therefore stays in equilibrium and there is no net forward or back reaction.
Questions: 9, 10, 11, 16, 17, 19, 21, 24, 25, 29, 31, 34, 35, 36, 37.
Year 12 Chemistry: Chapter 16.3 Equilibrium and Dilution
16.3 Dilution
For equilibria occurring in solution, the effect of diluting the solution by adding water is similar to changing the volume in gaseous equilibria.
A net reaction occurs in the direction the produces the greater number of particles.
Fe3+(aq) + SCN-(aq) Fe(SCN)2+(aq)
2 particles in solution 1 particle in solution
In terms of Le Chatelier’s principle, a net back reaction increases the total concentration of particles in solution, offsetting the effect of dilution. The concentration of Fe3+ and SCN- at the new equilibrium will be lower than their concentration prior to dilution as the equilibrium shift only partially offsets the change.
Changing the Temperature
Le Chatelier’s principle can also be used to determine the effect of heating on equilibrium mixtures.
2NO2(g) N2O4(g) + energy (exothermic)
brown colourless
Heating increases the energy of the substances in the mixture. Le Chatelier’s principle: the reaction opposes an increase in energy by removing energy – a net backwards reaction occurs.
For an endothermic reaction, net forward reactions occur as the temperature rises.
An increase in temperature in an equilibrium mixture results in:
J A net backward reaction (less products) for exothermic reactions.
J A net forward reaction (less reactants) for endothermic reactions.
The influence of catalysts
Catalysts increase the rate of reaction. The presence of a catalyst does not change the position of equilibrium. A catalyst therefore has no effect on the equilibrium yield of a reaction.
A catalyst greatly increases the rate at which an equilibrium is attained and it is for this reason that catalysts are employed in many industrial and biological systems.
16.4 Do all reactions reach equilibrium
Many reactions can be regarded as continuing until they are complete. These include:
J reactions that produce products such as gases that escape from the reaction mixture as they are formed. Continual loss of products drives these reactions forward.
J Reactions that form equilibria in which only minute quantities of reactants are present. The reaction of hydrogen chloride gas with water (equation in pervious handout).
READ: Extension ~ Equilibrium in the bloodstream (pg 276 – 278)
SUMMARY NOTES: pg 279 table
QUESTIONS: 9, 10, 11, 24, 25, 29, 31, 35, 36, 37