E N E. Higher Chemistry R G Y. Higher Exercise Book E R S

Transcription

1 E N E Higher Chemistry R G Y M A TA T E R S Higher Exercise Book

2 Exercise 1 Standard grade revision - bonding 1. Use your data book to write down the number of electrons in the outer shell of an atom of the following elements. hydrogen oxygen phosphorus (d) fluorine 2. The arrangement for the outer electrons of an oxygen atom is shown left. Draw similar diagrams to show the arrangement of the electrons in the outer shell of an atom of the following elements. O carbon sulphur nitrogen (d) bromine 3. An electron diagram to show the overlap of electron clouds in a water molecule is shown left. Draw a diagram to show electrons in the outer shells in the following molecules. nitrogen hydride (NH 3 ) hydrogen bromide (HBr) Silane (SiH 4 ) O H H 4. Copy and complete the table. Particle Na Ca Al Na + Ca 2+ Al 3+ No. of protons No. of electrons Charge 5. Copy and complete the table Particle Cl S N Cl S 2 N 3 No. of protons No. of electrons Charge 6. Copy and complete the table Particle No. of protons No. of electrons Charge Cu 2+ Pb

3 7. Predict the type of bonding (metallic, covalent, covalent network or ionic) for each chemical. (i) water (ii) magnesium oxide (iii) ammonia (iv) hydrogen chloride (v) copper (vi) silicon carbide (vii) iron (III) sulphate (viii) ammonium chloride (ix) caesium fluoride (x) propane (xi) cyclohexane (xii) manganese (xiii) lead (IV) oxide (xiv) silicon 8. The three dimensional shape of a methane molecule is shown left. Draw similar diagrams to show the three-dimensional shapes of the following molecules. H H H C H a) silane b) ammonia c) hydrogen chloride d) water 9. Identify the following types of bonding from the following descriptions. a) White solid that melts at 160 C; dissolves in water; does not conduct electricity. b) White solid that melts at 801 C; dissolves in water; conducts electricity in solution (but not in the solid state). c) Grey solid that melts at 1535 C; does not dissolve in water; conducts electricity in solid state. d) White solid that melts at1610 C; does not dissolve in water; does not conduct electricity e) A gas with a boiling point of 165 C; does not dissolve in water; does not conduct electricity. 11. Atoms can form ions with either a positive or negative charge. A B C LI + K + Mg2+ D E F Cl Br O 2 Identify two ions with the same electron arrangement as an argon atom. Identify two ions, which combine to form an insoluble substance. 2

4 Exercise 2 Standard grade revision - formulae and equations 1. Copy and complete the table. Compound Formula Compound Formula hydrochloric acid nitric acid sulphuric acid carbonic acid ethanoic acid sodium hydroxide calcium hydroxide barium hydroxide magnesium oxide copper (II) oxide aluminium oxide calcium carbonate potassium carbonate lithium carbonate 2. Copy and complete the following to give balanced equations. a) NaOH + HCl b) H 2 SO 4 + KOH c) CuO + HCl d) MgO + HNO 3 e) CaCO 3 + HCl f) Na 2 CO 3 + H 2 SO 4 g) LiOH + HCl h) H 2 SO 4 + HNO 3 i) PbO 2 + HCl j) Al 2 O 3 + HNO 3 Exercise 3 Standard grade revision - calculations A. Titration calculations can be worked out in the following way: Example: 40 cm 3 of potassium hydroxide were neutralised by 10 cm 3 0f 2 mol/l sulphuric acid. What is the concentration of the potassium hydroxide? equation: 2KOH + H 2 SO 4 K 2 SO 4 + 2H 2 O 2 moles of KOH 1 mole of H 2 SO 4 This means that there are twice as many moles of KOH as there are H 2 SO 4 No. of moles H 2 SO 4 = concentration x volume = 2 x 0.01 = 0.02 number of moles of KOH = 0.02 x 2 = 0.04 concentration of KOH = number of moles volume = = 1 mol/l 10 cm 3 = 0.01 l 40 cm 3 = 0.04 l 1. Calculate the concentration of 20 cm 3 sulphuric acid that would exactly neutralise 30 cm 3 of 2 mol / l potassium hydroxide. 10 cm 3 of 0.2 mol/l barium hydroxide. 3

5 2. Calculate the concentration of 10 cm 3 hydrochloric acid that would exactly neutralise: 20 cm 3 of 0.4 mol / l sodium hydroxide solution? 5.0 cm 3 of 0.2 mol / l calcium hydroxide solution? B. Calculations of mass from equations can be worked out in the following way: Example: What is the mass of CO 2 produced by burning 4.4g of C 3 H 8? Equation: C 3 H 8 + 5O 2 3CO 2 + 4H 2 O 1 mole of C 3 H 8 3 moles of CO 2 Formula = C 3 H 8 R.F.M. = = 44 Number of moles of C 3 H 8 = Mass of C 3 H 8 R.F.M. of C 3 H 8 = = 0.1 Formula = CO 2 R.F.M. = = 44 Number of moles of CO 2 = 3 x 0.1 = 0.3 Mass of CO 2 = 0.3 x 44g = 13.2g 3. What mass of carbon dioxide is produced by completely burning: - 36 g of carbon 1.5 g of ethane 4. What mass of mass of water is made by burning completely: g of methane 10 tonnes of ethane 30 g of hydrogen Exercise 4 Standard grade revision - acids and alkalis 1. Ann added 20 cm 3 0f 1mol/l sodium hydroxide solution to 20 cm 3 of 1mol/l sulphuric acid. Identify the statement(s), which can be applied to this experiment. A B C D E F The number of H + (aq) ions in the beaker decreased. The ph of the solution decreased. 2 The number of SO 4 (aq) ions in the beaker decreased. Water molecules formed during the reaction. A precipitate formed during the reaction. The final solution contained equal numbers of H + (aq) and OH (aq) ions. 2. Chromium (VI) oxide dissolves in water to form chromic acid. CrO 3 + H 2 O H 2 CrO 4 Suggest what is unusual about the reaction between chromium (VI) oxide and water. Chromic acid forms a salt with potassium hydroxide. (i) (ii) How many moles of potassium hydroxide would be needed to neutralise completely 1 mole of chromic acid? Name the salt formed in the reaction. (You may wish to use page 7 of the data booklet to help you.) 4

6 3. Many chemical compounds contain ions. A B C BaO CaBr 2 CuO D E F KCl NaBr SrCl 2 Identify the two compounds, which are bases 4. Many chemical reactions involve water. Identify the neutralisation reaction. A C 2 H 5 OH C 2 H 4 + H 2 O B CH 4 + 2O 2 CO 2 + 2H 2 O C C 2 H 2 + H 2 O C 2 H 4 O D 2H 2 + O 2 2H 2 O E C 6 H 12 O 6 6C + 6H 2 O F HNO 3 + KOH KNO 3 + H 2 O 5. The grid shows pairs of reactants. A B C Mg (s) + HCl (aq) CuSO 4 (aq) + Na 2 CO 3(aq) Fe (s) + CuSO 4(aq) D E F NaOH (aq) + HCl (aq) NH 4 NO 3(s) + NaOH (aq) MgO (s)) + H 2 SO 4(aq) Identify the pair(s) which is/are neutralisation reactions. Identify the reaction in which a precipitate is formed. Identify the reaction(s) in which a gas is formed. 6. The ph values of 1 mol/l solutions of some salts are shown in the table. Salt iron (II) sulphate aluminium chloride zinc (II) sulphate copper (II) nitrate sodium chloride potassium sulphate sodium carbonate potassium carbonate ph Identify the statement(s), which can be made about salts in the table. A B C D E The salts are neutral. The salts of transition metals are acidic. The sulphates are acidic. The salts of group 1 metals are neutral. The salts of hydrochloric acid are neutral. 5

7 7. 2 mol/l hydrochloric acid is often used in the chemistry laboratory. Identify the correct statement(s) about this solution. A B C D E It produces hydrogen when electrolysed. It contains more H + (aq) ions than Cl (aq) ions. It does not react with magnesium. 20 cm 3 is neutralised by 40 cm 3 of 1 mol/l sodium hydroxide. 200 cm 3 contains one mole of hydrogen chloride. 8. Chlorine reacts with water. Cl 2(g) + H 2 O (l) 2H + (aq) + OCl (aq) + Cl (aq) What happens to the ph of the water in a swimming pool when it is chlorinated? 9. Zinc reacts with sulphuric acid to produce hydrogen. Adding copper sulphate solution is thought to speed up the reaction. You are given the following reactants: 1 mol/l sulphuric acid 1 mol/l copper sulphate zinc lumps Describe how you would investigate the effect of adding copper sulphate solution on the speed of the reaction between zinc and sulphuric acid. 10. A solution of sodium hydroxide used in photography has a concentration of 80g/l. Calculate the concentration of sodium hydroxide in mol/l. 11. Magnesium sulphate can be converted to magnesium nitrate in 2 stages as follows. Stage 1 MgSO 4(aq) + Na 2 CO 3(aq) Na 2 SO 4(aq) + MgCO 3(s) Stage 2 MgCO 3(s) + HNO 3(aq) Mg(NO 3 ) 2(aq) + H 2 O (l) + CO 2(g) A B C reduction oxidation neutralisation D E F precipitation addition displacement Identify the type of reaction involved in Stage 1. Identify the type of reaction involved in Stage 2. 6

8 12. Solid sodium hydroxide often contains sodium carbonate as an impurity. This is caused by some of the sodium hydroxide reacting with carbon dioxide from the air. Instruction Card Preparation of approximately 1 mol/l sodium hydroxide solution CARE: Sodium hydroxide is highly corrosive. 1. Weigh out 41 g of sodium hydroxide pellets. 2. Dissolve in about 400 cm 3 of water. 3. Add a few drops of barium chloride solution to remove the carbonate ions. 4. Filter 5. Make the volume of solution up to exactly 1 litre. The equation foe Step 3 is: BaCl 2(aq) + Na 2 CO 3(aq) BaCO 3(s) + 2NaCl (aq) (i) (ii) Name an ion, which will be an impurity in the final sodium hydroxide solution. Rewrite the equation as an ionic equation omitting spectator ions. 7

9 Exercise 5 Reaction rates 1. Excess marble chips were added to 25 cm 3 of 2 mol l 1 hydrochloric acid. Which measurement, taken at regular intervals and plotted against time, would give the graph shown below? Measurement taken Time A. Temperature B. Volume of gas produced C. ph of solution D. Mass of the beaker and contents 2. 1 g of magnesium ribbon was added to excess dilute hydrochloric acid. The mass of hydrogen gas produced was measured and the results plotted in the graph below Mass of hydrogen gas released / g (d) Time / s From the graph estimate (i) the total mass of gas produced (ii) time taken for the reaction to finish. Calculate the average rate of reaction for the first 20 seconds. Calculate the average rate of reaction for the period 20 seconds to 40 seconds. Explain the change in reaction rate as the reaction proceeds. 3. Some marble chips (excess) were added to 50 cm 3 of 0.1 mol l 1 hydrochloric acid and the volume of carbon dioxide produced was measured over a period of time. The results were as follows. Time (seconds) Volume of CO 2 (cm 3 ) Draw a diagram of the apparatus that could be used to make these measurements. 8

10 Question 3 (continued) Draw a line graph of the measurements. (d) Using your graph, calculate the average reaction rate for the first 30 seconds of the experiment. The experiment was repeated with excess calcium carbonate powder and 25 cm 3 of 0.1 mol l 1 hydrochloric acid. Draw a dotted line on the above graph to show the graph you would expect to get with this second experiment. 4. Sodium thiosulphate was added to hydrochloric acid in a series of experiments. The time taken for the reaction was recorded. The results were recorded in a table: Volume of Na 2 S 2 O 3 (aq) used / cm 3 Volume of water added / cm 3 Relative concentration of Na 2 S 2 O 3 (aq) Volume of HCl (aq) added / cm 3 Time (t) for X to vanish / s Rate / s 1 A B C D E Describe how the time taken for the reaction was actually measured. Copy and complete the table by calculating the rate of reaction for experiments A to E. From the first two columns of the table, which factor is being kept constant throughout the series of experiments? (d) (i) What conclusion can be drawn concerning the concentration of the thiosulphate solution and reaction rate? (ii) Explain why changing the concentration affects the reaction rate. (e) State one other way of increasing the reaction rate. Exercise 6 Reaction rates 1. The following potential energy diagram represents the energy changes in a chemical reaction. The activation energy for the reaction in kj mol 1, is Potential energy / kj mol 1 A. 20 B B. 50 D. 70 reactant products Reaction 9

11 2. Which of the following is not a correct statement about the effect of a catalyst? The catalyst A. Provides an alternative route to the products B. Lowers the energy which molecules need for successful collisions C. Provides energy so that more molecules have successful collisions D. Forms bonds with reacting molecules. 3. When copper carbonate is reacted with excess acid, carbon dioxide is produced. The curves shown opposite were obtained under different conditions. The change from P to Q could be brought about by A. increasing the concentration of the acid B. decreasing the mass of copper carbonate C. decreasing the particle size of the copper carbonate D. adding a catalyst. 4. For any chemical reaction, the temperature is a measure of Volume of CO 2 Time P Q A. The average kinetic energy of the particles which react B. The average kinetic energy of all the particles C. The activation energy D. The minimum kinetic energy required before a reaction occurs. 5. An experiment was carried out at four temperatures. The table shows the times taken for the reaction to occur. The results show that Temperature / C Time / s A. The reaction endothermic B. The activation energy increases with increasing temperature C. The rate of reaction is directly proportional to the temperature D. A small rise in temperature results in a large increase in reaction rate. 6. The diagram on the right could represent A. the fermentation of glucose B. neutralisation of an acid by an alkali C. the combustion of sucrose D. the reaction of a metal with acid. Rate of reaction Temperature 7. The potential energy diagram opposite refers to the reversible reaction involving reactants R and products P. What is the enthalpy change in kj mol 1 for the reverse reaction P R? A. +30 B. +10 C. -10 D. -40 Potential energy /kj mol Reaction pathway 10

12 8. The enthalpy change for the forward reaction in the graph can be represented by A. x B. y C. x + y D. x y Potential energy x y 9. A B C Reaction pathway rate rate rate temperature temperature D E F temperature rate rate rate temperature temperature temperature Identify the graph, which shows how the rate varies with temperature in the reaction between sodium thiosulphate and dilute sulphuric acid the action of amylase on starch 10. The relative volumes of hydrogen produced in a given time for three reactions are plotted on the graph. volume of hydrogen II I III time Curve I is for the reaction of excess zinc powder with 100 cm 3 of 0.1 mol l 1 hydrochloric acid. A B C Excess magnesium powder 50 cm 3 of 0.1 mol l 1 hydrochloric acid 50 cm 3 of 0.2 mol l 1 sulphuric acid D E F Excess iron powder 200 cm 3 of 0.1 mol l 1 hydrochloric acid 100 cm 3 of 0.05 mol l 1 sulphuric acid Identify the two chemicals that would react to give the results plotted in curve II. Identify the two chemicals that would react to give the results plotted in curve III. 11

13 11. Many factors influence the rates of reactions. A B C Particle size of reactants temperature Surface area available for reaction D E F Activation energy concentration Average kinetic energy of reactant molecules Identify the factor which, if increased, causes an increase in the factor shown in box F. Identify the factor(s) which, if increased, would make the reaction slower. 12. Hydrogen peroxide can be used to clean contact lenses. In this process, the enzyme catalase is added to break down hydrogen peroxide. The equation for the reaction is: 2H 2 O 2 2H 2 O + O 2 The rate of oxygen production was measured in three laboratory experiments using the same volume of hydrogen peroxide at the same temperature. Experiment Concentration of H 2 O 2 / mol l -1 Catalyst used A 0.2 Yes B 0.4 Yes C 0.2 no The curve obtained from experiment A is shown Volume of oxygen produced / cm time / s A Calculate the average rate of the reaction over the first 40 s. Copy the graph (no graph paper needed) and add curves to the graph to show the results of experiments B and C. Label each curve clearly. 13. Graph 1 shows the distribution of kinetic energies of molecules in a gas at 30 C. Add a dotted line to Graph 1 to show the distribution of kinetic energies at 20 C. Number of molecules Graph 2 Graph 1 Kinetic energy 12

14 Question 13 (continued) In Graph 2, the shaded area represents the number of molecules with the required activation energy E A, for the reaction to occur. Draw a line to show how a catalyst affects the energy of activation. Number of molecules Kinetic energy E A 14. Copy and complete the following table on industrial reactions and the catalysts used: Catalyst Process Reaction Product(s) Vanadium (V) oxide Contact sulphuric acid Haber 4NH 3 + 5O 2 4NO + 6H 2 O Breaking down long chain hydrocarbon molecules nickel Unsaturated oils + H 2 saturated fats Exercise 7 Reaction rates 1. Magnesium ribbon reacts with dilute hydrochloric acid to form a salt, and hydrogen gas. Write a balanced equation for the above reaction. Indicate using labelled diagrams, two methods of following the course of the reaction. In each case, state what measurements are necessary to determine the rate of the chemical reaction. Draw graphs (no graph paper required) of the expected results from each of the experiments described in g of marble chips were added to 50 cm 3 of 1 mol l 1 hydrochloric acid. State three ways in which the rate of the chemical reaction can be increased. 3. Write a brief note explaining the difference between catalyst poisoning and catalyst regeneration. 4. Powdered silver metal behaves as a heterogeneous catalyst in the decomposition of 5 volume hydrogen peroxide solution to give water and oxygen as the only products. Note: Concentrations of peroxide solutions are quoted by volume number with higher numbers relating to solutions of greater concentrations. Write a balanced equation for the decomposition of hydrogen peroxide solution. What does a heterogeneous catalyst mean? Explain any observed differences in the rates of oxygen production, with time, in each of the following pairs of silver metal/hydrogen peroxide solution reactions. (i) (ii) (iii) 1 g silver powder in 250 cm 3 5 volume solution at 20 C, and then at 40 C. 1 g of silver powder in 250 cm 3 5 volume solution at 20 C, then with 250 cm 3 10 volume solution at the same temperature. 250 cm 3 10 volume solution at 30 C with 1 g of silver powder, then with 1 g of silver foil at the same temperature. 13

15 5. All new cars are fitted with a catalytic converter to deal with exhaust emissions. (d) Name two metals which are used in the construction of catalytic converters in cars. Name three gases emitted by car exhausts, which are considered to be harmful. Explain how the catalyst changes these gases into harmless ones. Explain why cars fitted with a catalytic converter must use unleaded fuel. 6. The stomach produces an enzyme called pepsin. This helps digest food while it is in the stomach. Define the term enzyme. What are the two optimum conditions for the enzyme pepsin? Which type of food does pepsin digest? 7. Study the potential energy diagram for an uncatalysed reaction. Using a dotted line, add a line to the diagram to show energy change for a catalysed reaction. Using the values shown on the diagram, calculate the energy values for the uncatalysed reaction. (i) (ii) E A H 150 Potential energy / 80 kj mol 1 40 reactants Reaction products (d) What type of enthalpy reaction has accompanied this reaction? Explain. Explain why the catalyst speeds the rate of the reaction. 8. Copy and complete the following table to identify the functions of the enzymes. Enzyme zymase papain cellulase sucrase lipase amylase Reactions 14

16 Exercise 8 Excess chemicals g of methane was mixed with 0.2 mol of oxygen gas and exploded. Which chemical was in excess? mol of magnesium was added to 100 cm 3 of 0.5 mol l 1 of hydrochloric acid. Which chemical was in excess? g of sodium carbonate was added to 100 cm 3 of 0.01 mol l 1 of sulphuric acid. Which chemical was in excess? 4. 1 g of propane was mixed with 10 g of oxygen and exploded. Calculate the mass of carbon dioxide produced g of zinc metal was added to 100 cm 3 of 0.02 mol l 1 sulphuric acid. Calculate the mass of hydrogen released g of magnesium was added to 50 cm 3 of 0.4 mol l 1 of hydrochloric acid. Which chemical was in excess and hence calculate the mass of the remaining reactant. 7. A 1 g sample of marble (impure calcium carbonate) was added to excess nitric acid. The reaction produced 0.4 g of carbon dioxide. Calculate the mass of the calcium carbonate that has reacted and hence, the percentage of calcium carbonate in the marble. Exercise 9 Enthalpy 1. Copy and complete the following equation: H = H (products) 2. Complete the labelling of the energy diagram by identifying X and Y. Potential energy X Y Reaction pathway Is this reaction exothermic or endothermic? Explain. 3. Redraw the energy diagram in question 2 and add a dotted line to show the reaction when a catalyst is used. 4. Write out the balanced equations to show the standard enthalpies (do not forget to show the state symbols): combustion of ethane solution of copper (II) sulphate addition of nitric acid to potassium hydroxide 15

17 5. The following is the equation for calculating the energy change in an experiment E h = c x m x t Explain what is meant by the symbols c, m and T. 6. When 2 g of sodium hydroxide were added to water 1.68 kj was liberated. Calculate the standard enthalpy of solution of sodium hydroxide kj of heat was released when 0.24 g of methanol (CH 3 OH) completely burned. Calculate the standard enthalpy of combustion of methanol. 8. When 100 cm 3 of dilute nitric acid, concentration 0.1 mol l 1, is neutralised by 100 cm 3 of 0.1 mol l 1 sodium hydroxide solution, 570 J of heat is liberated. Calculate the enthalpy of neutralisation kj of heat was produced when 1 g of sulphur burned. Calculate the enthalpy of combustion of sulphur. Exercise 10 Enthalpy 1. When 0.1 g of propanol (C 3 H 7 OH) was burned, the heat produced raised the temperature of 100 cm 3 water from 18 C to 23.2 C. Use the above data to calculate the enthalpy of combustion of propanol. A data booklet gives the value for the enthalpy combustion of propanol as H = 2010 kj mol 1 Explain the differences between this and your calculated value. 2. Discuss how the terms lattice-breaking and hydration are involved in the enthalpy of solution of an ionic solid. 3. Dissolving sodium hydroxide in 100 cm 3 of water, releases 4.2 kj of heat energy in the solution. Use the data to calculate the rise in temperature. 4. The mixing of 50 cm 3 of 0.5 mol l 1 hydrochloric acid and 50 cm 3 of 0.5 mol l 1 of sodium hydroxide produced a temperature rise of 3.2 C. Calculate the enthalpy of neutralisation. 5. The following results are taken from a notebook of a pupil. Experiment: Addition of 50 cm 3 of 2.0 mol l 1 sodium hydroxide to 50 cm 3 of 1.0 mol l 1 hydrochloric acid. NaOH (aq) + HCl (aq) NaCl (aq) + H 2 O (l) Mass of sodium hydroxide solution = 50 g Mass of hydrochloric acid solution = 50 g Initial temperature of NaOH (aq) = 21.9 C Initial temperature of HCl (aq) = 21.5 C Average temperature on mixing = X Highest temperature during experiment = 28.5 C Work out which solution is in excess. Calculate the enthalpy of neutralisation of hydrochloric acid by sodium hydroxide. 16

18 6. When some propane was completely burned 222 kj of heat was produced. If the enthalpy of combustion of propane is 2220 kj mol l 1, what mass of propane was burned? moles of a hydrocarbon was completely burned to give kj of heat. Calculate the enthalpy of combustion of the hydrocarbon and, using the data booklet, identify the hydrocarbon. 8. In an experiment 1.4 g of ethene were completely burned and the heat produced was kj. What would be the temperature rise if the heat was absorbed by 500 cm 3 of water. Exercise 11 Periodic Table 1. What is meant by the first ionisation energy? (i) Is the process involved exothermic or endothermic? (ii) Explain your answer. Write equations corresponding to the first ionisation energies of sodium and calcium and alongside each equation write in the H value with its appropriate sign. 2. Write equations corresponding to: B. the second ionisation energy of magnesium C. the third ionisation energy of aluminium and alongside each equation write in the H value with its appropriate sign. 3. Explain why the second ionisation energy of each alkali metal is much greater than that of the halogen in the same period. Explain why the third ionisation energy of magnesium is greater than that for aluminium. Why is there no value given in the data booklet for the fourth ionisation energy of lithium? 4. Calculate the energy required to bring about each of the following changes. Al (g) Al 3+ (g) + 3e K + (g) K 3+ (g) + 2e 5. Which three elements (in the first 20) have the lowest first ionisation energies? (i) In which group are these elements? (ii) Explain why they have the lowest first ionisation energies? 6. Which two elements (in the first 20) have the highest first ionisation energies? (i) In which group are these elements? (ii) Explain why they have the highest first ionisation energies? 7. Which elements show the greatest difference between their first ionisation energies and their second ionisation energies? Explain your answer. 17

19 Exercise 12 Periodic Table 1. Which element would require the most energy to convert 1 mole of gaseous atoms into gaseous ions carrying two positive charges? A. scandium B. titanium C. vanadium D. chromium 2. Which element has the greatest attraction for bonding electrons within a bond? A. caesium B. oxygen C. fluorine D. iodine 3. Which equation represents the first ionisation enthalpy for chlorine? A. Cl (g) + e Cl (g) B. Cl (g) Cl + (g) + e C. Cl + (g) + e Cl (g) D. Cl (g) Cl (g) + e 4. How does the pattern of density change going down group 7 of the Periodic Table? Explain your answer. 5. What is the trend in boiling points going down group 8 of the periodic table? Explain your answer. 6. Explain how the atomic size changes as you go across the periodic table from lithium to fluorine. 7. What is meant by a periodic property? 8. What is meant by the term electronegativity? 9. For each pair of elements identify the atom which has the greater electronegativity. oxygen and sulphur strontium and calcium caesium and neon (d) lithium and iodine Exercise 13 Bonding; structure and properties 1 Decide whether each of the following molecules is polar or non-polar. oxygen Hydrogen chloride Carbon monoxide (d) Hydrogen 2. By drawing the shape of and considering the symmetry, decide whether each of the following molecules is polar. water (H 2 O) ammonia (NH 3 ) methane (CH 4 ) (d) chloroform (CHCl 3 ) (e) methanol (CH 3 OH) (f) carbon dioxide (CO 2 ) (g) carbon tetrachloride (CCl 4 ) 18

20 3. Hydrogen bonds will be set up between polar molecules, when these molecules contain a hydrogen atom covalently bonded to an atom such as nitrogen, oxygen and fluorine. Which of the following substances will exhibit hydrogen bonding? hydrogen bromide (HBr) propan 1 ol (CH 3 -CH 2 -CH 2 -OH) aminoethane (CH 3 -CH 2 -NH 2 ) (d) methoxymethane (CH 3 -O-CH 3 ) (e) tetrafluoromethane (CF 4 ) (f) propanone (CH 3 -CO-CH 3 ) (g) ethanoic acid (CH 3 -COOH) 4. Ammonia, NH 3, exhibits a high degree of hydrogen bonding, yet there is no such bonding in hydrogen. Explain why this so. How do the strengths of the forces involved in hydrogen bonding compare with the strengths of van der Waals forces? 5. Analysis of hydrogen fluoride shows the existence of molecules with relative molecular masses of 20, 40 and 60. Explain the origin of these molecules. Exercise 14 Bonding; structure and properties 1. Taking water as an example, describe what happens when a covalent bond is formed between the atoms. Illustrate your answer. 2. Draw shapes of the following molecules: methane ammonia water (d) hydrogen fluoride 3. Explain what is meant by the following terms: Ionic bond Dipole van der Waals forces (d) Hydrogen bond 4. State what metallic bonding means. Illustrate your answer. 5. For each chemical, state whether the type of bonding is metallic, covalent, covalent network, polar or ionic. Chemicals (i) water (ii) methane (iii) sodium chloride (iv) methanol (v) silicon (vi) carbon tetrachloride (vii) potassium (viii) hydrogen (ix) copper fluoride gas (x) copper (xi) aluminium (xii) aluminium oxide (xiii) carbon dioxide chloride fluoride (xiv) ammonia (xv) ethene Which of the above chemicals contain non-polar molecules even though the bonding is polar covalent. 19

21 6. When the following solids melt, what forces of attraction are overcome by the heating? sodium metal hydrogen gas water (d) pentane (e) propene (f) potassium chloride (g) hydrogen fluoride (h) silicon dioxide 7. Explain why propanone has a higher boiling point than butane even though they have the same relative molecular mass. 8. Explain why sodium chloride has a melting point of 800 C while that of aluminium chloride has a melting point of 192 C. Exercise 15 Bonding; structure and properties 1 The diagram illustrates the trend in the boiling points of Group 7 hydrides HCl, HBr and HI Boiling Point/K HF HCl HBr HI Why is the boiling point of HI higher than that of HBr? Estimate the value for the boiling point of hydrogen fluoride (HF) if it were to follow the trend in the graph. The actual boiling point of HF is 292 K Account for this unexpectedly high value. 2. Liquid ammonia boils at 33 C but liquid phosphine, PH 3, boils at 87.5 C. Explain this difference in terms of bonding. 3. A pupil is comparing the boiling points of an alkane with ethanol, C 2 H 5 OH, to examine the influence of hydrogen bonding. To examine the influence of hydrogen bonding, the boiling point of which alkane should be compared with that of ethanol? Explain your answer. Which of the two compounds has hydrogen bonding between the molecules? 20

22 4. Arrange the molecules below in order of increasing viscosity. CH 3 CH 2 OH CH 2 CH CH 2 CH 2 CH 2 OH OH OH OH OH ethanol propane 1,2,3 triol ethane 1,2, - diol 5. Taking sodium chloride as your example, explain, in terms of bonding, what happens to an ionic crystal when it dissolves in water. 6. Explain the following: Potassium chloride dissolves readily in water but not in cyclohexane. tetrachloromethane dissolves in hexane but not in water. 7. Lithium iodide is moderately soluble in non-polar solvents whereas caesium fluoride is not. On the basis of this evidence, deduce the difference in bonding in the two compounds. Exercise 16 The Mole 1. Calculate the mass of each of the following: 0.5 mol of sulphur dioxide, SO 2. 5 mol of calcium nitrate, Ca(NO 3 ) 2. 2 mol of sodium carbonate. (d) 0.1 mol of copper (II) carbonate. 2. Calculate the number of moles which are contained in each of the following: 14g of nitrogen, N g of potassium nitrate g of ammonium sulphate (d) 10.2 g of aluminium oxide. 3. How many moles of sodium hydroxide must be dissolved to make each of the following solutions? 200 cm 3 of 1 mol l cm 3 of 0.5 mol l cm 3 of 2 mol l 1 4. What is the concentration of each of the following solutions of hydrochloric acid? 2 mol of hydrogen chloride dissolved in 200 cm 3 of solution. 0.5 mol of hydrogen chloride dissolved in 250 cm 3 of solution. 0.1 mol of hydrogen chloride dissolved in 500 cm 3 of solution. 5. What is the volume of each of the following solutions sodium hydroxide? 2 mol l 1 solution which contains 0.4 mol solute? 0.5 mol l 1 solution which contains 0.1 mol solute? 0.1 mol l 1 solution which contains 0.2 mol solute? 21

23 6. How many grams of each of the following substances must be weighed out to make up each of the following solutions? 500 cm 3 of 0.1 mol l 1 magnesium sulphate, MgSO cm 3 of 2 mol l 1 sodium nitrate, NaNO 3. 2 litres of 0.1 mol l 1 sodium hydroxide, NaOH. 7. What is the concentration of each of the following solutions? (d) g of potassium carbonate, K 2 CO 3, dissolved in 2 litres of solution g of copper (II) sulphate dissolved in 250 cm 3 of solution. 10 g of ammonium nitrate dissolved in 100 cm 3 of solution g of hydrogen chloride dissolved in 250 cm 3 of solution. Exercise 17 The Mole 1. Decide which has more atoms: 16 g of sulphur or 27 g of aluminium 2.3 g of sodium or 2.3 g of potassium 10 g of argon or 10 g of neon 2. Decide which has more molecules: 28 g of carbon monoxide or 32 g of sulphur dioxide 32 g of oxygen or 1 g of hydrogen 6.4 g of sulphur dioxide or 6.4 g of carbon dioxide 3. Decide which has more atoms: 64 g of sulphur dioxide or 28 g of carbon monoxide 4.4 g of carbon dioxide or 8 g of argon 20 g of neon or 2 g of hydrogen 4. Decide which has more hydrogen atoms: 17 g of ammonia or 16 g of methane 3 g of ethane or 3 g of water 0.2 g of hydrogen or 2 g of hydrogen fluoride 5. Decide which has more ions: 16 g of copper (II) sulphate or 58 g of magnesium hydroxide 48 g of ammonium carbonate or 35 g of potassium hydroxide 45 g of sodium sulphate or 45 g of potassium nitrate 6. Calculate the number of atoms in each of the following elements. 10 g of calcium 24 g of carbon 2.43 g of magnesium (d) 4 g of argon 22

24 7. Calculate the number of molecules in each of the following compounds. 1.1 g of carbon dioxide 34 g of ammonia 360 g of water (d) 3.2 g of methane 8. Calculate the number of formula units in each of the following compounds g of potassium oxide 14.8 g of calcium hydroxide 5.85 g of sodium chloride (d) g of iron (II) sulphate Exercise 18 The Mole 1. Decide which has the greater volume, measurements being made at the same temperature and pressure. 17 g of ammonia or 17 g of methane 20 g of neon or 20 g of nitrogen 4 g of hydrogen or 44 g of carbon dioxide (d) 6.4 g of sulphur dioxide or 8 g of oxygen 2. The volume of 0.22 g of propene is 118 cm 3. Calculate the volume of 2 mol of propene at this temperature and pressure. 3. Calculate the volume of 2.4 g of ethene. (molar volume = 23.6 litres mol 1 ) 4. The volume of 1 g of hydrogen is 11.6 litres. Calculate the volume of 4 mol of hydrogen at this temperature and pressure. 5. Using the densities given in the data booklet, calculate the volume of 10 g of: hydrogen argon 6. 3 g of an alkane occupies a volume of 2.24 litres. What is the molecular formula for the alkane? (molar volume = 22.4 litres mol 1 ) 7. Using the densities given in the data booklet, calculate the number of moles in 10 litres of: helium nitrogen 23

25 Exercise 19 The Mole (d) 1. From the following data calculate the approximate formula of the hydride gas X. Mass of plastic bottle (empty) = g Mass of plastic bottle + gas X = g Capacity of plastic bottle = 1 litre Molar volume of gas X = 23.6 litres mol A flask, capacity 600 cm 3, was used to calculate the molar volume of sulphur dioxide. The following data was obtained Mass of evacuated flask = g Mass of flask + sulphur dioxide = g Calculate the volume of 1 mol of sulphur dioxide (under the experimental conditions). 2. Mg (s) + 2HCl (aq) MgCl 2(aq) + H 2(g) What volume of hydrogen gas is produced when 5 g of magnesium is added to 100 cm 3 of 2 mol l 1 hydrochloric acid? (Molar volume = 23.6 litres mol 1.) 3. What volume of oxygen is needed to completely burn 4 g of sulphur to form sulphur dioxide? (Use the density of oxygen given in the data booklet page 3.) 4. What mass of magnesium oxide is obtained when excess magnesium is ignited in 5 litres of oxygen? (Molar volume = 22.8 litres mol 1.) 5. What mass of water is produced when 2 litres of hydrogen is burned in excess oxygen? (Use the density of hydrogen given in the data booklet page 3.) 6. 2 litres of hydrogen is ignited with 2 litres of oxygen. Calculate the mass of water produced. (Use the density of hydrogen and oxygen given in the data booklet page 3.) 7. Calculate the volume of oxygen required for the complete combustion of 1 g of ethene. (molar volume = 23.2 litres mol 1.) 8. Using the density of hydrogen given in the data booklet, what volume of hydrogen gas is produced when 13 g of zinc is added to excess acid? 9. The first stage in the extraction of copper from sulphide ores such as chalcopyrite (CuFeS 2 ) involves heating the ore with sand and oxygen. 4 CuFeS 2(s) + 2SiO 2(s) + 5O 2(g) 2Cu 2 S.FeS (l) + 2FeSiO 3(l) + 4SO 2(g) Calculate the volume of oxygen required to react completely with 1472 kg chalcopyrite. (Molar volume = 24.8 litres mol 1.) 11. The flow diagram shows how benzene (C 6 H 6 ) can be obtained from naphtha. hydrogen naphtha hexane benzene C 6 H 6 C 6 H 14 catalytic reforming Calculate the volume of hydrogen produced when 1 kg of hexane is reformed. (Use the density of hydrogen given in the data booklet page 3.) 24

26 Exercise 20 The Mole (e) 1. In each of the following reactions decide the ratio of the volume of products to the volume of reactants. H 2(g) + Cl 2(g) 2HCl (g) N 2(g) + 3H 2(g) 2NH 3(g) 2C (s) + O 2(g) 2CO (g) (d) C 2 H 4(g) + 3O 2(g) 2CO 2(g) + 2H 2 O (l) (e) CuO (s) + CO (g) Cu (s) + CO 2(g) 2. N 2(g) + 2O 2(g) NO 2(g) What volume of nitrogen dioxide is produced when 100 cm 3 of nitrogen is sparked in excess oxygen? 3. When each of the following volumes of gases is burned completely, calculate the volume of oxygen required the volume of carbon dioxide produced. (i) (ii) (iii) (iv) 100 cm 3 of methane 2 litres of carbon dioxide 250 cm 3 of ethene 150 cm 3 of propane cm 3 of butane gas is mixed with 75 cm 3 of oxygen and the mixture exploded. C 4 H 10(g) + 13/2 O 2(g) 4CO 2(g) + 5H 2 O (l) Calculate the volume and composition of the resulting gas mixture. 5. A mixture of 60 cm 3 of hydrogen and 40 cm 3 of carbon monoxide is passed over hot copper (II) oxide until no further reaction occurred. Calculate the volume of the resulting gas. 6. The reaction of sulphur dioxide with oxygen to produce sulphur trioxide does not go to completion. Assuming only 50% of the available sulphur dioxide is converted to sulphur trioxide, calculate the volume and composition of the resulting gas mixture when 100 cm 3 of sulphur dioxide is mixed with 100 cm 3 of oxygen. 7. If 100 cm 3 of propene is burned completely with 900 cm 3 of oxygen, what will be the composition of the resulting mixture? cm 3 of ethyne, C 2 H 2, is completely burned in 220 cm 3 of oxygen. What will be the volume and composition of the resulting gas mixture? What will be the volume if the reaction is repeated in a container heated to 200 C? 9. A mixture of 80 cm 3 CO and 150 cm 3 of O 2 is exploded. After cooling, the residual gas is shaken with sodium hydroxide solution. Which gas is absorbed by the sodium hydroxide solution? What is the reduction in volume of the residual gas on shaking with the solution? What volume of gas remains? 25