ACSI HL Equilibrium 2022 (Teacher) (1)

IBDP Chemistry HL/ Equilibrium Page 1 Anglo − Chinese School (Independent) Year 5 (2022) IBDP Chemistry HL (IBDP syllabus Topic 7 ) 7.1 Equilibrium - Essential Idea: Many reactions are reversible. These reactions will reach a state of equilibrium when the rates of the forward and reverse reaction are equal. The position of equilibrium can be controlled by changing the conditions. (IBDP syllabus Topic 1 7 ) 1 7 .1 The equilibrium law - Essential Idea: The position of equilibrium can be quantified by the equilibrium law. The equilibrium constant for a particular reaction only depends on the temperature. TOPIC 7 EQUILIBRIUM TEACHER COPY – WITH SUGGESTED SOLUTIONS

IBDP Chemistry HL/Equilibrium Page 2 7.1 Equilibrium Nature of science: Obtaining evidence for scientific theories— isotopic labelling and its use in defining equilibrium. (1.8) Common language across different disciplines —the term dynamic equilibrium is used in other contexts, but not necessarily with the chemistry definition in mind. (5.5) Understandings: • A state of equilibrium is reached in a closed system when the rates of the forward and reverse reactions are equal. • The equilibrium law describes how the equilibrium constant (Kc) can be determined for a particular chemical reaction. • The magnitude of the equilibrium constant indicates the extent of a reaction at equilibrium and is temperature dependent. • The reaction quotient (Q) measures the relative amount of products and reactants present during a reaction at a particular point in time. Q is the equilibrium expression with non- equilibrium concentrations. The position of the equilibrium changes with changes in concentration, pressure, and temperature. • A catalyst has no effect on the position of equilibrium or the equilibrium constant. Applications and skills: • The characteristics of chemical and physical systems in a state of equilibrium. • Deduction of the equilibrium constant expression ( Kc) from an equation for a homogeneous reaction. • Determination of the relationship between different equilibrium constants ( Kc) for the same reaction at the same temperature. • Application of Le Châtelier’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. Guidance: • Physical and chemical systems should be covered. • Relationship between Kc values for reactions that are multiples or inverses of one another should be covered. • Specific details of any industrial process are not required. 17.1 The equilibrium law Nature of science: Employing quantitative reasoning— experimentally determined rate expressions for forward and backward reactions can be deduced directly from the stoichiometric equations and allow Le Châtelier’s principle to be applied. (1.8, 1.9) Understandings: • Le Châtelier’s principle for changes in concentration can be explained by the equilibrium law. • The position of equilibrium corresponds to a maximum value of entropy and a minimum in the value of the Gibbs free energy. • The Gibbs free energy change of a reaction and the equilibrium constant can both be used to measure the position of an equilibrium reaction and are related by the equation, ∆G = –RT lnK

IBDP Chemistry HL/Equilibrium Page 3 Applications and skills: • Solution of homogeneous equilibrium problems using the expression for Kc. • Relationship between ∆G and the equilibrium constant. • Calculations using the equation ∆G = –RT lnK Guidance: • The expression ∆G = –RT lnK is given in the data booklet in section 1. • Students will not be expected to derive the expression ∆G = –RT lnK. • The use of quadratic equations will not be assessed. 7.1 Equilibrium 7.1.1 Reversible Chemical Reaction • Consider the following example: When heated strongly in the open, calcium carbonate decomposes to form calcium oxide and carbon dioxide: CaCO3 (s) CaO (s) + CO2 (g) As the carbon dioxide, CO 2 escapes, it is impossible for the calcium oxide, CaO to be converted back to calcium carbonate, CaCO 3. We say that the above reaction is irreversible. However, if calcium carbonate is heated in a closed evacuated container (closed system), the following reverse reaction becomes possible. CaO (s) + CO 2 (g) CaCO3 (s) In the closed system, both the forward and reverse reactions take place simultaneously, the reaction is known as a reversible reaction. Reversible reactions are chemical reactions that can take place in both directions and they are denoted by the use of a double headed arrow, ⇌. CaCO 3 (s) ⇌ CaO (s) + CO2 (g) • Reversible reactions have the following characteristics: o Reversible reactions can only occur in a closed system, where there is no gain or loss of chemicals to or from the surroundings. o The reaction that proceeds from left to right is known as the forward reaction. o The reaction that proceeds from right to left is known as the reverse reaction. o At equilibrium, the rates of the forward and reverse reactions are equal. o Equilibrium can be achieved from either direction: you can start with either reactants or products. heat

IBDP Chemistry HL/Equilibrium Page 4 7.1.2 Dynamic Equilibrium • In a reversible reaction, dynamic equilibrium is achieved when the rates of the forward and reverse reaction are equal. At this stage, the concentration of all the reactants and the products remain constant if the temperature is kept constant. • Consider the following dynamic equilibrium: H2 (g) + I2 (g) ⇌ 2HI (g) • The dynamic nature of the equilibrium can be demonstrated by adding a small amount of radioactive iodine (I2*) to the above reaction at a state of equilibrium. Measurements will show that after some time, the mixture contains radioactive hydrogen iodide (H I*), indicating that the reaction is still continuing. After a period of time, the forward reaction rate becomes equal to the reverse reaction rate. A state of dynamic equilibrium is re– established. • At a microscopic (molecular) scale, reactions continue but both the forward and reverse reaction rates are equal. At dynamic equilibrium, there is no change in macroscopic properties such as colour and density of the reaction mixture. (TOK: Refer to Pg 313 of chemistry course companion.) 7.1.3 Graph of Concentration vs Time and Rate vs Time • Consider the decomposition of the colourless gas dinitrogen tetroxide (N 2O4) to form dark brown nitrogen dioxide (NO2) gas. N2O4 (g) ⇌ 2NO2 (g) Colourless Brown • The reaction begins with an initial concentration of 0.0400 mol dm–3 of N2O4 and the formation of NO2 is indicated by the appearance of a brown colour. The progress of the reaction can be monitored by measuring the intensity of the brown colour with a spectrophotometer.

IBDP Chemistry HL/Equilibrium Page 5 • According to the chemical equation, 2.0 mol of NO2 forms for each mole of N 2O4 that decomposes, so the concentration of N2O4 remaining at any time can be computed by subtracting half the concentration of NO2 formed from the initial concentration of N 2O4. As time passes, the concentration of N 2O4 decreases and the concentration of NO 2 increases until both concentrations level off at constant equilibrium values. • The concentrations of the reactants and products reach constant values, not because the reactions stop, but because the rates of the forward and reverse reactions become equal. A dynamic equilibrium is reached. • Because reaction rates depend on concentrations, the rate of the forward reacti