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IB Chemistry Review Topics 7 & 8
Bronsted-Lowry acid/ base
Chemical equilibrium
Hydrogen ion
Molarity
Conjugate acid/ base
Equilibrium constant
Electron pair
Activation energy
Lewis acid/ base
Le Chatelier’s Principle
Donor
Strong acids/ base
Catalyst
Acceptor
Weak acids/ base
Haber Process
Concentration
pH
Contact process
Endothermic
System
Proton
Exothermic
Know how to determine whether something is Bronsted-Lowry acid or base.
Know how to determine the conjugate base of an acid and the conjugate acid of a base.
Be able to determine if a species will behave as a Lewis acid or Lewis base.
Explain what makes something a strong acid or strong base.
Be able to recognize the assigned strong and weak acids and bases.
Explain what equilibrium has to do with weak acids and weak bases.
Explain why a weak acid or base will have a different pH then a strong acid or base.
Explain the relationship between forward and reverse reactions and a double-ended arrow.
Be able to write an equilibrium constant from a balanced equation.
Be able to explain endothermic and exothermic reactions in terms of energy as a reactant or product.
Be able to explain how Le Chatelier’s Principle is related to pressure/volume, temperature and amounts of products and
reactants in a reaction.
Be able to explain how a catalyst affects a reaction in equilibrium.
Explain the purpose of the Haber process.
Be able to explain what factors are involved in the Haber process and their connection to Le Chatelier’s Principle.
Explain the purpose of the Contact process.
Be able to explain what factors are involved in the Contact process and their connection to Le Chatelier’s Principle.
1. The equilibrium between nitrogen dioxide, NO2, and dinitrogen tetroxide, N2O4, is shown below.
2NO2(g)
N2O4(g)
Kc = 0.01
What happens when the volume of a mixture at equilibrium is decreased at a constant temperature?
2. Which does dynamic imply about chemical equilibria?
3.
At equilibrium, __________.
4. Which one of the following is true concerning the Haber process?
5. Which one of the following will change the value of an equilibrium constant?
6. For the following reaction Kc = 1.0 × 10–5 at 30 °C.
2NOCl(g)
2NO(g) + Cl2(g)
Which relationship is correct at equilibrium at this temperature?
7. Consider the equilibrium between methanol, CH3OH(l), and methanol vapor, CH3OH(g).
CH3OH(l)
CH3OH(g)
What happens to the position of equilibrium and the value of Kc as the temperature decreases?
8. What is the equilibrium constant expression, Kc, for the following reaction? Remember, you don’t
actually need a number, just show how to get it.
N2O4(g)
2NO2(g)
9. Consider the endothermic reaction below.
5CO(g) + I2O5(g)
5CO2(g) + I2(g)
According to Le Chatelier’s principle, which change would result in an increase in the amount of CO2?
10. What is the effect of an increase of temperature on the yield and the equilibrium constant for the
following reaction?
2H2(g) + CO(g)
CH3OH(l)
∆HO = –128 kJ
11. Consider the following equilibrium reaction.
2SO2(g) + O2(g)
2SO3(g)
∆Ho = –197 kJ
Which change in conditions will increase the amount of SO3 present when equilibrium is reestablished?
12. Consider the following reversible reaction.
Cr2O72–(aq) + H2O(l)
2CrO42–(aq) + 2H+(aq)
What will happen to the position of equilibrium and the value of Kc when more H+ ions are added at
constant temperature?
13. What effect will an increase in temperature have on the Kc value and the position of equilibrium in the
following reaction?
N2(g) + 3H2(g)
2NH3(g)
ΔH = –92 kJ
14. An increase in temperature increases the amount of chlorine present in the following equilibrium.
PCl5(s)
PCl3(l) + Cl2(g)
What is the best explanation for this (keep in mind kinetics as well)?
15. What will happen when at a constant temperature if more iodide ions, I–, are added to the equilibrium
below?
I2(s) + I–(aq)
I3–(aq)
16. What changes occur when the temperature is increased in the following reaction at equilibrium? Include
both the direction of the shift and how Kc is affected.
Br2(g) + Cl2(g)
2BrCl(g) ∆Hο = +14 kJ mol–1
17. Methanol may be produced by the exothermic reaction of carbon monoxide gas and hydrogen gas.
CO(g) + 2H2(g)
CH3OH(g) ∆HO = –103 kJ
 State the equilibrium constant expression, Kc, for the production of methanol.
State and explain the effect of changing the following conditions on the amount of methanol present at
equilibrium:
 increasing the temperature of the reaction at constant pressure.
 increasing the pressure of the reaction at constant temperature.
 The conditions used in industry during the production of methanol are a temperature of 450 °C
and pressure of up to 220 atm. Explain why these conditions are used rather than those that could
give an even greater amount of methanol.
18. An example of a homogeneous reversible reaction is the reaction between hydrogen and iodine.
H2(g) + I2(g)
2HI(g)
 Outline the characteristics of a homogeneous chemical system that is in a state of equilibrium.
 Deduce the expression for the equilibrium constant, Kc.
 Predict what would happen to the position of equilibrium and the value of Kc if the pressure is
increased from 1 atm to 2 atm.
19. A state of equilibrium can exist when a piece of copper metal is placed in a solution of copper(II)
sulfate. Outline the characteristics of a chemical system in dynamic equilibrium.
20. For an exothermic reaction state how an increase in temperature would affect both Kc and the position
of equilibrium.
21. In carbonated drinks containing dissolved carbon dioxide under high pressure, the following dynamic
equilibrium exists. CO2(aq)
CO2(g)
Describe the effect of opening a carbonated drink container and outline how this equilibrium is affected.
22. Consider the following equilibrium.
2SO2(g) + O2(g)
2SO3(g)
ΔHo = –198 kJ mol–1
 Deduce the equilibrium constant expression, Kc, for the reaction.
 State and explain the effect of increasing the temperature on the yield of sulfur trioxide.
 State the effect of a catalyst on the value of Kc.
 State and explain the effect of a catalyst on the position of equilibrium.
23. Consider the following reaction taking place at 375 °C in a 1.00 dm3 closed container.
Cl2(g) + SO2(g)
SO2Cl2(g)
∆HO = –84.5 kJ
 Deduce the equilibrium constant expression, Kc, for the reaction.
 If the temperature of the reaction is changed to 300 °C, predict, stating a reason in each case,
whether the equilibrium concentration of SO2Cl2 and the value of Kc will increase or decrease.
24. The equation for the main reaction in the Haber process is: N2(g) + 3H2(g)
2NH3(g)
- ∆H



State and explain the effect on the equilibrium yield of ammonia with increasing the pressure and the
temperature.
In practice, typical conditions used in the Haber process involve a temperature around 450 °C and a pressure
of 250 atm. Explain why these conditions are used rather than conditions that would give a higher yield.
Suggest why this reaction is important for humanity.

A chemist claims to have developed a new catalyst for the Haber process, which increases the yield of
ammonia. State the catalyst normally used for the Haber process, and comment on the claim made by this
chemist.
1. Consider the equilibrium below.
CH3CH2COOH(aq) + H2O(l)
CH3CH2COO–(aq) + H3O+(aq)
Which species represent a conjugate acid-base pair?
2. The pH of a solution changes from pH = 2 to pH = 5. What happens to the concentration of the
hydrogen ions during this pH change?
3. Which species behave as Brønsted-Lowry acids in the following reversible reaction?
H2PO4–(aq) + CN–(aq)
HCN(aq) + HPO42–(aq)
4. Which of the following are weak acids in aqueous solution?
I.
CH3COOH
II.
H2CO3
III. HCl
5. What is the conjugate base of H2CO3 according to the Brønsted-Lowry theory?
6. For equal volumes of 1.0 mol dm–3 solutions of hydrochloric acid, HCl(aq), and methanoic acid,
HCOOH(aq), which statements are correct?
I. HCl dissociates more than HCOOH
II. HCl is a better electrical conductor than HCOOH
III. HCl will neutralize more NaOH than HCOOH
7. Which are definitions of an acid according to the Brønsted-Lowry and Lewis theories?
8. A Br∅nsted-Lowry base is defined as a substance that __________
9. A substance that is capable of acting as both an acid and as a base is __________.
10. The molar concentration of hydronium ion (or H+) in pure water at 25°C is __________.
11. Water has a very lowKc (also referred to as Kw). This indicates that __________.
12. What would the pH be of a solution with a high concentration of hydroxide ions?
13. Nitric acid is a strong acid. This means that __________.
14. Of the following acids, __________ is not a strong acid.
15. What is the conjugate acid of NH3?
16. The conjugate base of HSO4- is __________.
17. The conjugate acid of HSO4- is __________.
18. What is the conjugate base of OH- ?
19. Which list contains only strong acids?
20. What is the formula of the conjugate base of the hydrogenphosphate ion, HPO42–?
21. Define the terms acid and base according to the Brønsted-Lowry theory and state one example of a
weak acid and one example of a strong base.
22. Describe two different methods, one chemical and one physical, other than measuring the pH, that
could be used to distinguish between ethanoic acid and hydrochloric acid solutions of the same
concentration.
23. Black coffee has a pH of 5 and toothpaste has a pH of 8. Identify which is more acidic and deduce
how many times the [H+] is greater in the more acidic product.
24. Define the terms acid and base according to the Brønsted-Lowry theory. Distinguish between a
weak base and a strong base. State one example of a weak base.
25. Define an acid in terms of the Lewis theory.
Deduce, giving a reason, whether NF3 is able to function as a Lewis acid or as a Lewis base.
26. Describe two different properties that could be used to distinguish between a 1.00 mol dm–3 solution
of a strong monoprotic acid and a 1.00 mol dm–3 solution of a weak monoprotic acid.
27. Explain, using the Brønsted-Lowry theory, how water can act either as an acid or a base. In each
case identify the conjugate acid or base formed.
28. Define a Brønsted-Lowry acid.
29. Deduce the two acids and their conjugate bases in the following reaction:
H2O(l) + NH3(aq)
OH–(aq) + NH4+(aq)
30. Explain why the following reaction can also be described as an acid-base reaction.
F–(g) + BF3(g)
BF4–(s)