ATAR Chemistry Unit 3: Topic 1 Chemical Equilibrium Systems 3.1.1 Chemical Equilibrium 3.1.1.1 recognise that chemical systems may be open (allowing matter and energy to be exchanged with the surroundings) or closed (allow energy, but not matter, to be exchanged with the surroundings) 3.1.1.2 understand that physical changes are usually reversible, whereas only some chemical reactions are reversible 3.1.1.3 appreciate that observable changes in chemical reactions and physical changes can be described and explained at an atomic and molecular level 3.1.1.4 symbolise equilibrium equations by using β in balanced chemical equations 3.1.1.5 understand that, over time, physical changes and reversible chemical reactions reach a state of dynamic equilibrium in a closed system, with the relative concentrations of products and reactants defining the position of equilibrium 3.1.1.6 explain the reversibility of chemical reactions by considering the activation energies of the forward and reverse reactions 3.1.1.7 analyse experimental data, including constructing and using appropriate graphical representations of relative changes in the concentration of reactants and product against time, to identify the position of equilibrium. 3.1.2 Factors that affect equilibrium 3.1.2.1 explain and predict the effect of temperature change on chemical systems at equilibrium by considering the enthalpy change for the forward and reverse reactions 3.1.2.2 explain the effect of changes of concentration and pressure on chemical systems at equilibrium by applying collision theory to the forward and reverse reactions 3.1.2.3 apply Le Châtelier’s principle to predict the effect changes of temperature, concentration of chemicals, pressure and the addition of a catalyst have on the position of equilibrium and on the value of the equilibrium constant. 3.1.3 Equilibrium constants 3.1.3.1 understand that equilibrium law expressions can be written for homogeneous and heterogeneous systems and that the equilibrium constant (Kc), at any given temperature, indicates the relationship between product and reactant concentrations at equilibrium 3.1.3.2 deduce the equilibrium law expression from the equation for a homogeneous reaction and use equilibrium constants (Kc), to predict qualitatively, the relative amounts of reactants and products (equilibrium position) when Kc is very small the follow assumption can be made: [reactants]initial ≈ [reactants]equilibrium 3.1.3.3 deduce the extent of a reaction from the magnitude of the equilibrium constant 3.1.3.4 use appropriate mathematical representation to solve problems, including calculating equilibrium constants and the concentration of reactants and products. 3.1.4 Properties of Acids and Bases 3.1.4.1 understand that acids are substances that can act as proton (hydrogen ion) donors and can be classified as monoprotic or polyprotic depending on the number of protons donated by each molecule of the acid 1 3.1.4.2 distinguish between strong and weak acids and bases in terms of the extent of dissociation, reaction with water and electrical conductivity and distinguish between the terms strong and concentrated for acids and bases. 3.1.5 pH scale 3.1.5.1 understand that water is a weak electrolyte and the self-ionisation of water is represented by Kw = [H+][OH–]; Kw can be used to calculate the concentration of hydrogen ions from the concentration of hydroxide ions in a solution Kw is taken to be 1×10–14 at 25°C and is given in the Chemistry formula and data booklet. 3.1.5.2 understand that the pH scale is a logarithmic scale and the pH of a solution can be calculated from the concentration of hydrogen ions using the relationship pH = –log10 [H+] 3.1.5.3 use appropriate mathematical representation to solve problems for hydrogen ion concentration [H+(aq)], pH, hydroxide ion concentrations [OH–(aq)] and pOH. 3.1.6 Bronsted-Lowry Model 3.1.6.1 recognise that the relationship between acids and bases in equilibrium systems can be explained using the Brønsted-Lowry model and represented using chemical equations that illustrate the transfer of hydrogen ions (protons) between conjugate acid-base pairs 3.1.6.2 recognise that amphiprotic species can act as Brønsted-Lowry acids and bases 3.1.6.3 identify and deduce the formula of the conjugate acid (or base) of any Brønsted-Lowry base (or acid) 3.1.6.4 appreciate that buffers are solutions that are conjugate in nature and resist a change in pH when a small amount of an acid or base is added; Le Châtelier’s principle can be applied to predict how buffer solutions respond to the addition of hydrogen ions and hydroxide ions. Buffer calculations are not required. 3.1.7 Dissociation Constants 3.1.7.1 recognise that the strength of acids is explained by the degree of ionisation at equilibrium in aqueous solution, which can be represented with chemical equations and equilibrium constants (Ka) 3.1.7.2 determine the expression for the dissociation constant for weak acids (Ka) and weak bases (Kb) from balanced chemical equations Students should consider hydrochloric, nitric and sulfuric acids as examples of strong acids, and carboxylic and carbonic acids (aqueous carbon dioxide) as weak acids. 3.1.7.3 analyse experimental data to determine and compare the relative strengths of acids and bases 3.1.7.4 use appropriate mathematical representation to solve problems, including calculating dissociation constants (Ka and Kb) and the concentration of reactants and products. 3.1.8 Acid-base indicators 3.1.8.1 understand that an acid-base indicator is a weak acid or a weak base where the components of the conjugate acid-base pair have different colours; the acidic form is of a different colour to the basic form 3.1.8.2 explain the relationship between the pH range of an acid-base indicator and its pKa value 3.1.8.3 recognise that indicators change colour when the pH = pKa and identify an appropriate indicator for a titration, given equivalence point of the titration and pH range of the indicator. The colour change can be considered to take place over a range of pKa ± 1. 3.1.9 Volumetric Analysis 3.1.9.1 distinguish between the terms end point and equivalence point 3.1.9.2 recognise that acid-base titrations rely on the identification of an equivalence point by measuring the associated change in pH, using chemical indicators or pH meters, to reveal an observable end point 3.1.9.3 sketch the general shapes of graphs of pH against volume (titration curves) involving strong and weak acids and bases. Identify and explain their important features, including the intercept with pH axis, equivalence point, buffer region and points where pKa = pH or pKb = pOH 2 3.1.9.4 use appropriate mathematical representations and analyse experimental data and titration curves to solve problems and make predictions, including using the mole concept to calculate moles, mass, volume and concentration from volumetric analysis data. Mandatory Practical Acid-base titration to calculate the concentration of a solution with reference to a standard solution. 3 3.1.1 Chemical equilibrium 3.1.1.1 recognise that chemical systems may be open (allowing matter and energy to be exchanged with the surroundings) or closed (allow energy, but not matter, to be exchanged with the surroundings) 3.1.1.2 understand that physical changes are usually reversible, whereas only some chemical reactions are reversible 3.1.1.3 appreciate that observable changes in chemical reactions and physical changes can be described and explained at an atomic and molecular level 3.1.1.4 symbolise equilibrium equations by using β in balanced chemical equations 3.1.1.5 understand that, over time, physical changes and reversible chemical reactions reach a state of dynamic equilibrium in a closed system, with the relative concentrations of products and reactants defining the position of equilibrium 3.1.1.6 explain the reversibility of chemical reactions by considering the activation energies of the forward and reverse reactions 3.1.1.7 analyse experimental data, including constructing and using appropriate graphical representations of relative changes in the concentration of reactants and product against time, to identify the position of equilibrium. 3.1.1.1 recognise that chemical systems may be open (allowing matter and energy to be exchanged with the surroundings) or closed (allow energy, but not matter, to be exchanged with the surroundings) Define the following and give two examples of each: Open system __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ Closed system __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ 4 Explain why dynamic equilibrium can only be established in a closed system. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ Exercises Textbook: Questions 1, 4, 5, 6 Page 13,14, Question 1 page 44 Question 8, 14 page 45 5 3.1.1.2 understand that physical changes are usually reversible, whereas only some chemical reactions are reversible Define: Reversible reaction __________________________________________________________________________________ __________________________________________________________________________________ Give two examples of physical changes that are reversible. Write chemical equations to explain the process occurring. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ Give two examples of chemical reactions that are reversible. Write chemical equations for each example. You can find some examples in your textbook. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ P 13 questions 2,3, 6 P 44 q2, 6 3.1.1.3 appreciate that observable changes in chemical reactions and physical changes can be described and explained at an atomic and molecular level Describe at a molecular level, any changes that would still occur once the system above has reached equilibrium. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ Explain, at a molecular level, any observable changes that would occur as the above reaction proceeds towards equilibrium. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ 7 When red mercury(II) oxide is heated it decomposes as shown in the equation below. Explain, at an atomic level, any observable changes that would occur in the mixture as it approaches equilibrium. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ Explain at a molecular level, what is occurring to the reactants and products of the following reaction at equilibrium. π2(π) + π2(π) β 2ππ(π) Δπ» = +180.5 ππ½πππ −1 __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ P 14 numbers 4, 8,9 P 45 no 9, P 48 no 37 38, 39 8 3.1.1.4 symbolise equilibrium equations by using β in balanced chemical equations Write an equation to represent the equilibrium that exists between liquid water and water vapour in a closed water bottle. __________________________________________________________________________________ __________________________________________________________________________________ P 46 no 20 9 3.1.1.5 understand that, over time, physical changes and reversible chemical reactions reach a state of dynamic equilibrium in a closed system, with the relative concentrations of products and reactants defining the position of equilibrium Below is the equation representing the Haber process. Explain the diagram above, with reference to the Haber process. Why has only 0.6 moles of ammonia formed rather than 2 moles, even though 1.0 moles of nitrogen gas and 3.0 moles of hydrogen gas were initially present? __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ P 14 nos 9 P 45 no 16 P 46 no 21, 32, 33 34 10 3.1.1.6 explain the reversibility of chemical reactions by considering the activation energies of the forward and reverse reactions __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ Exercises P 14 no 7 11 3.1.1.7 analyse experimental data, including constructing and using appropriate graphical representations of relative changes in the concentration of reactants and product against time, to identify the position of equilibrium. Identify on the graph below, the point at which equilibrium is established. Identify the reactants and the products from the graph below. __________________________________________________________________________________ Describe the reaction that has occurred including the changes in amounts of each substance. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ 12 Identify when the system below first reaches equilibrium. Identify how many times the system reaches equilibrium. 13 Exercises P 44 no 4 P 45 no 15 3.1.2 Factors that affect equilibrium 3.1.2.1 explain and predict the effect of temperature change on chemical systems at equilibrium by considering the enthalpy change for the forward and reverse reactions 3.1.2.2 explain the effect of changes of concentration and pressure on chemical systems at equilibrium by applying collision theory to the forward and reverse reactions 3.1.2.3 apply Le Châtelier’s principle to predict the effect changes of temperature, concentration of chemicals, pressure and the addition of a catalyst have on the position of equilibrium and on the value of the equilibrium constant. 3.1.2.1 explain and predict the effect of temperature change on chemical systems at equilibrium by considering the enthalpy change for the forward and reverse reactions __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ 14 __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ 15 Predict the effect that an increase in temperature will have on the reaction below. 2X(g) + Y2(g) X Z(g) ΔH = –100 kJ Y2 Z Keq A decrease 1 by 2 decrease increase increase B increase 1 by 2 increase decrease decrease C decrease by 2 decrease increase increase D increase by 2 increase decrease decrease The reaction between bromine and chlorine is shown below. Bromine and chlorine were placed in a sealed vessel and allowed to reach equilibrium. Predict the effect that increasing the temperature of the reaction between bromine and chlorine will have on the concentration of the contents of the vessel. Show your reasoning. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ Exercises P 31 3, P 48 no 36 16 3.1.2.2 explain the effect of changes of concentration and pressure on chemical systems at equilibrium by applying collision theory to the forward and reverse reactions __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ 17 For the equilibrium, 0.8 mol of nitrogen and 0.4 mol of oxygen were placed in a closed vessel and allowed to reach equilibrium, at which time 0.6 mol of nitrogen remained. Using the grid below, construct the graph of the relative changes in concentration in the vessel. 18 Exercises P31 no 3, 4,5 P 45 no 10, 11, P 45 no 1`7 P 46 no 18, 19, 3.1.2.3 apply Le Châtelier’s principle to predict the effect changes of temperature, concentration of chemicals, pressure and the addition of a catalyst have on the position of equilibrium and on the value of the equilibrium constant. State Le Chatelier’s Principle __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ The contact process is an important industrial process for making sulfuric acid (H2SO4). This process occurs in three stages. a) Explain which of the stages would be affected by a change in pressure. (2 marks) _______________________________________________________________________ _______________________________________________________________________ _______________________________________________________________________ _______________________________________________________________________ 19 20 What effect do each of the following changes have on the equilibrium position for this reversible reaction? ππΆπ5(π) + βπππ‘ β ππΆπ3(π) + πΆπ2(π) (a) (b) (c) (d) addition of Cl2 increase of pressure removal of heat removal of PCl3 __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ How is the equilibrium position of this reaction affected by the following changes? πΆ(π ) + π»2 π(π) + βπππ‘ β πΆπ(π) + π»2(π) (a) (b) (c) (d) lowering the temperature increasing the pressure removing H2 adding H2 __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ 21 __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ What effect does each change have on the equilibrium position of this reaction? π2(π) + 3π»2(π) β 2ππ»3(π) + 92 ππ½ (a) (b) (c) (d) addition of heat increase in pressure addition of catalyst removal of heat __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ Exercises P 22 numbers 1,2,3 4 P 30 numbers 1,2 P 31 no 7,8,9 P 44 no 3,7, 22 3.1.3 Equilibrium constants 3.1.3.1 understand that equilibrium law expressions can be written for homogeneous and heterogeneous systems and that the equilibrium constant (Kc), at any given temperature, indicates the relationship between product and reactant concentrations at equilibrium 3.1.3.2 deduce the equilibrium law expression from the equation for a homogeneous reaction and use equilibrium constants (Kc), to predict qualitatively, the relative amounts of reactants and products (equilibrium position) when Kc is very small the follow assumption can be made: [reactants]initial ≈ [reactants]equilibrium 3.1.3.3 deduce the extent of a reaction from the magnitude of the equilibrium constant 3.1.3.4 use appropriate mathematical representation to solve problems, including calculating equilibrium constants and the concentration of reactants and products. 3.1.3.1 understand that equilibrium law expressions can be written for homogeneous and heterogeneous systems and that the equilibrium constant (Kc), at any given temperature, indicates the relationship between product and reactant concentrations at equilibrium Explain what Kc tells you about the relationship between products and reactants in a system at equilibrium. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ 23 Compare homogeneous and heterogenous systems and give examples of each by writing an equilibrium equation highlighting both types of systems. Why is it important to understand the difference between them in terms of writing the equilibrium expressions? __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ P 41 no 1,2,3 P 44 no 5, 6 , 12 P 46 no 22, 23, 3.1.3.2 deduce the equilibrium law expression from the equation for a homogeneous reaction and use equilibrium constants (Kc), to predict qualitatively, the relative amounts of reactants and products (equilibrium position) when Kc is very small the follow assumption can be made: [reactants]initial ≈ [reactants]equilibrium Consider the following reaction: 2H2(g) + S2(g) β 2H2O(g) + 2Cl2(g) Kc = 9.4 x 105 at 750°C A mixture of the above was allowed to come to equilibrium in a closed 2.0L container at 750°C. The equilibrium concentrations of H2 and H2S gases were analysed and found to be 0.234 M and 0.442 M respectively. a. State what the value of the equilibrium constant for this reaction tells you about the extent of the reaction. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ b. Show an expression for the equilibrium constant for this reaction. __________________________________________________________________________________ 24 c. Calculate the equilibrium concentration of S2(g) in the mixture. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ P 41 no 4, 5, 6,7,8,9, 10, P 45 no 13 P 46 no 24, 26, 27, 28, 29,30 3.1.3.3 deduce the extent of a reaction from the magnitude of the equilibrium constant Complete the following: If Kc < 10-4, there is a negligible __________ reaction and mainly ____________ present at equilibrium. The position of equilibrium is to the ___________. If 10-4 < Kc < 104, the equilibrium mixture consists of significant amounts of both __________ and ____________. If Kc > 104, there is extensive ____________ reaction and mainly ______________ present at equilibrium. The position of equilibrium is to the __________. P 42 no 11,12 25 3.1.3.4 use appropriate mathematical representation to solve problems, including calculating equilibrium constants and the concentration of reactants and products. An equilibrium mixture of CH4, H2O, H2 and CO was prepared in a 2.00L flask by adding 0.600 mol of CH4 and 0.400 mol of H2O to the reaction vessel. At equilibrium there was 0.110 mol of CO present. a. Determine the amounts of CH4, H2O and H2 at equilibrium. b. Calculate the equilibrium concentration of CH4, H2O, H2 and CO. c. Calculate the equilibrium constant, Kc, for the reaction. Exercises P 42 no 13, 14, 15, 16, 17,28 P 46 no 25, 31, P 48 no 35 P 49 no 40 26 3.1.4 Properties of Acids and Bases 3.1.4.1 understand that acids are substances that can act as proton (hydrogen ion) donors and can be classified as monoprotic or polyprotic depending on the number of protons donated by each molecule of the acid 3.1.4.2 distinguish between strong and weak acids and bases in terms of the extent of dissociation, reaction with water and electrical conductivity and distinguish between the terms strong and concentrated for acids and bases. 3.1.4.1 understand that acids are substances that can act as proton (hydrogen ion) donors and can be classified as monoprotic or polyprotic depending on the number of protons donated by each molecule of the acid Complete the table below Type of Acid Monoprotic Number of protons to donate Examples Equation for dissociation HNO3, CH3COOH HNO3(aq) + H2O(l) → NO3- (aq) + H3O+(aq) Diprotic Triprotic P57 no 5,6,8 P62 no 3 P 83 no 12 27 3.1.4.2 distinguish between strong and weak acids and bases in terms of the extent of dissociation, reaction with water and electrical conductivity and distinguish between the terms strong and concentrated for acids and bases. Strong Acids Readily donates protons and completely dissociates in water. E.g. HCl(aq) + H2O(l) → Cl- (aq) + H3O+(aq) Weak Does not readily __________ protons and _________ dissociates in water. Bases Readily _________ protons and __________ dissociates in water Does not readily ________ protons and _____________ dissociates in water. Draw a diagram to illustrate the differences between the terms strong and concentrated. Explain how measurements of electrical conductivity can be used to distinguish between strong and weak acids. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ 28 __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ Exercises P 62 no 1,2,4.,5 P83 no 1,13 3.1.5 pH Scale 3.1.5.1 understand that water is a weak electrolyte and the self-ionisation of water is represented by Kw = [H+][OH–]; Kw can be used to calculate the concentration of hydrogen ions from the concentration of hydroxide ions in a solution Kw is taken to be 1×10–14 at 25°C and is given in the Chemistry formula and data booklet. 3.1.5.2 understand that the pH scale is a logarithmic scale and the pH of a solution can be calculated from the concentration of hydrogen ions using the relationship pH = –log10 [H+] 3.1.5.3 use appropriate mathematical representation to solve problems for hydrogen ion concentration [H+(aq)], pH, hydroxide ion concentrations [OH–(aq)] and pOH. 3.1.5.1 understand that water is a weak electrolyte and the self-ionisation of water is represented by Kw = [H+][OH–]; Kw can be used to calculate the concentration of hydrogen ions from the concentration of hydroxide ions in a solution Kw is taken to be 1×10–14 at 25°C and is given in the Chemistry formula and data booklet. Complete the following reaction to represent the self-ionisation of water. H2O(l) + H2O(l) β Calculate the pH of a 0.001M NaOH solution. P 70 no 2,4,5 P 84 no 16, 29 3.1.5.2 understand that the pH scale is a logarithmic scale and the pH of a solution can be calculated from the concentration of hydrogen ions using the relationship pH = –log10 [H+] Calculate the pH for each of the following solutions. Solution 0.010 M HCl Calculation 0.050 M H2SO4 0.010 M CH3COOH 0.010 M NaOH 0.00050 M Ba(OH)2 P 70 no 3,6 P 84 no 17, 18 19,20, 30 3.1.5.3 use appropriate mathematical representation to solve problems for hydrogen ion concentration [H+(aq)], pH, hydroxide ion concentrations [OH–(aq)] and pOH. Complete the following mathematical relationships. a. pH = b. [H3O+] = c. pOH = d. [H3O+] x [OH-] = e. AT 25°C, the value of Kw = f. At 25°C, pH + pOH = Calculate the pH of a solution formed by mixing 20.00 mL of 0.30 M HCl with 30.00 mL or 0.15 M HNO3 Calculate the pH of the final solution obtained by mixing 20 mL of 0.20 M KOH solution with 80 mL of 0.20 M HCl solution. 31 Exercises P 84 no 21,22,23,24,25.26,27, P 85 no 31 3.1.6 Brønsted-Lowry acids and bases 3.1.6.1 recognise that the relationship between acids and bases in equilibrium systems can be explained using the BrønstedLowry model and represented using chemical equations that illustrate the transfer of hydrogen ions (protons) between conjugate acid-base pairs 3.1.6.2 recognise that amphiprotic species can act as Brønsted-Lowry acids and bases 3.1.6.3 identify and deduce the formula of the conjugate acid (or base) of any Brønsted-Lowry base (or acid) 3.1.6.4 appreciate that buffers are solutions that are conjugate in nature and resist a change in pH when a small amount of an acid or base is added; Le Châtelier’s principle can be applied to predict how buffer solutions respond to the addition of hydrogen ions and hydroxide ions. Buffer calculations are not required. 3.1.6.1 recognise that the relationship between acids and bases in equilibrium systems can be explained using the Brønsted-Lowry model and represented using chemical equations that illustrate the transfer of hydrogen ions (protons) between conjugate acidbase pairs Complete the table below. Ionic equation NH3(aq) + H2O(l) β Conjugate pairs NH3 and NH4+, H2O and OH- HCl(aq) + OH-(aq) → H2SO4(aq) + H2O(l) → CH3COOH(aq) + OH- → P 57 no 3,7, P 83 no 7,9,10,11, 32 3.1.6.2 recognise that amphiprotic species can act as Brønsted-Lowry acids and bases What is an amphiprotic substance? __________________________________________________________________________________ __________________________________________________________________________________ Give two examples of substances which are amphiprotic and write chemical equations to explain. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ P 57 no 6,8,9 3.1.6.3 identify and deduce the formula of the conjugate acid (or base) of any BrønstedLowry base (or acid) Label the acid, base, conjugate acid and conjugate base for the equation below. HCl(aq) + OH-(aq) → Cl-(aq) + H2O(l) P 57 no 1,2,4 5 33 3.1.6.4 appreciate that buffers are solutions that are conjugate in nature and resist a change in pH when a small amount of an acid or base is added; Le Châtelier’s principle can be applied to predict how buffer solutions respond to the additionn of hydrogen ions and hydroxide ions. Buffer calculations are not required. Define ‘buffer’ __________________________________________________________________________________ __________________________________________________________________________________ Explain why a mixture of HCl and NaOH does not form a buffer solution. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ 34 P 72 no 2,3 4, 35 3.1.7 Dissociation Constants 3.1.7.1 recognise that the strength of acids is explained by the degree of ionisation at equilibrium in aqueous solution, which can be represented with chemical equations and equilibrium constants (Ka) 3.1.7.2 determine the expression for the dissociation constant for weak acids (Ka) and weak bases (Kb) from balanced chemical equations Students should consider hydrochloric, nitric and sulfuric acids as examples of strong acids, and carboxylic and carbonic acids (aqueous carbon dioxide) as weak acids. 3.1.7.3 analyse experimental data to determine and compare the relative strengths of acids and bases 3.1.7.4 use appropriate mathematical representation to solve problems, including calculating dissociation constants (Ka and Kb) and the concentration of reactants and products. 3.1.7.1 recognise that the strength of acids is explained by the degree of ionisation at equilibrium in aqueous solution, which can be represented with chemical equations and equilibrium constants (Ka) Compare strong and weak acids. Use equations and provide examples to illustrate your answer. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ P 76 no 1, P 83 no 4,5 36 3.1.7.2 determine the expression for the dissociation constant for weak acids (Ka) and weak bases (Kb) from balanced chemical equations Students should consider hydrochloric, nitric and sulfuric acids as examples of strong acids, and carboxylic and carbonic acids (aqueous carbon dioxide) as weak acids. Write the expression for the dissociation constant for the following acids and bases. CH3COOH H2CO3 NH3 P 76 no 1,2,3,4 3.1.7.3 analyse experimental data to determine and compare the relative strengths of acids and bases Revise all questions from experiments 3.1.4 and 3.1.5 where you measured the pH of some acids and bases, and you measured the conductivity of acids and bases. When using pH or electrical conductivity to compare strong and weak acids an bases, it is important to keep the concentration of each substance the same. Explain why. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ 37 3.1.7.4 use appropriate mathematical representation to solve problems, including calculating dissociation constants (Ka and Kb) and the concentration of reactants and products. The value of Ka or Kb is a measure of the __________ of an acid or base, or the degree of dissociation of an acid or base. The ____________ the value, the stronger the acid or base. Four acids, W, X, Y and Z, with concentrations of 0.10 M, have pH values of 2.4, 1.0, 6.6 and 3.1 respectively. Determine and explain: Which, if any, is completely ionised. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ Which if any is almost completely un-ionised. __________________________________________________________________________________ __________________________________________________________________________________ Their order in increasing acid strength (weakest to strongest) __________________________________________________________________________________ __________________________________________________________________________________ 38 Lactic acid, HC3H5O3, which is found in milk, has a Ka of 1.38 x 10-4. Calculate the percentage dissociation of lactic acid in a 0.100 M solution. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ The concentration of CH3COOH at 25°C is 1.00 M. The value of Ka is 1.75 x 10-5. Calculate the concentration of the ethanoate ion at this temperature. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ 39 Calculate the pH of a 0.500 M solution of ammonia if the Kb at 25°C is 1.80 x 10-5. __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ 3.1.8 Acid-base indicators 3.1.8.1 understand that an acid-base indicator is a weak acid or a weak base where the components of the conjugate acidbase pair have different colours; the acidic form is of a different colour to the basic form 3.1.8.2 explain the relationship between the pH range of an acid-base indicator and its pKa value 3.1.8.3 recognise that indicators change colour when the pH = pKa and identify an appropriate indicator for a titration, given equivalence point of the titration and pH range of the indicator. The colour change can be considered to take place over a range of pKa ± 1. 3.1.8.1 understand that an acid-base indicator is a weak acid or a weak base where the components of the conjugate acid-base pair have different colours; the acidic form is of a different colour to the basic form. By representing the acid form of an indicator as HIn, write a reaction to show the reversible reaction which occurs when the acid dissociates in water to form its conjugate base, In -. 40 3.1.8.2 explain the relationship between the pH range of an acid-base indicator and its pKa value Consider the following indicator solution equilibrium: HIn(aq) + H2O(l) β H3O+(aq) + In-(aq) Ka at 25 = 1.0 x 10-7 Write an expression for the acidity constant for the above reaction. Calculate the pH of the solution when the acid and base forms are present in equal concentrations, i.e. [HIn] = [In-]. What would you see happening at this point? __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ P82 no 2,3,4,6,7 P 83 no 6,14 41 3.1.8.3 recognise that indicators change colour when the pH = pKa and identify an appropriate indicator for a titration, given equivalence point of the titration and pH range of the indicator. The colour change can be considered to take place over a range of pKa ± 1. 42 P82 no 5,7b P 83 no 2, 3, 15 P 84 no 29,30 P 100 no 10, 11 P 113 no 3, 3.1.9 Volumetric Analysis 3.1.9.1 distinguish between the terms end point and equivalence point 3.1.9.2 recognise that acid-base titrations rely on the identification of an equivalence point by measuring the associated change in pH, using chemical indicators or pH meters, to reveal an observable end point 3.1.9.3 sketch the general shapes of graphs of pH against volume (titration curves) involving strong and weak acids and bases. Identify and explain their important features, including the intercept with pH axis, equivalence point, buffer region and points where pKa = pH or pKb = pOH 3.1.9.4 use appropriate mathematical representations and analyse experimental data and titration curves to solve problems and make predictions, including using the mole concept to calculate moles, mass, volume and concentration from volumetric analysis data. 3.1.9.1 distinguish between the terms end point and equivalence point __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ __________________________________________________________________________________ P 99 no 2, 3.1.9.2 recognise that acid-base titrations rely on the identification of an equivalence point by measuring the associated change in pH, using chemical indicators or pH meters, to reveal an observable end point P 99 no 3, P 113 no 5,6.7, 8, P 113 no 15,19 43 3.1.9.3 sketch the general shapes of graphs of pH against volume (titration curves) involving strong and weak acids and bases. Identify and explain their important features, including the intercept with pH axis, equivalence point, buffer region and points where pKa = pH or pKb = pOH P84 no 28, P 113 no 1,2, P 114 no 13. 20, 44 3.1.9.4 use appropriate mathematical representations and analyse experimental data and titration curves to solve problems and make predictions, including using the mole concept to calculate moles, mass, volume and concentration from volumetric analysis data. P 99 no 4,6, 7, 8, 9, P 109 no 4,5,6,7 P 113 no 4, P 114 no 9,10,11,12,16,17, 18, 21 45
