Chemistry
The course's student books as first-class content: every chapter opens the book PDF at its own page range — the same pages, citations and spec joins the tutor grounds in.
Jim Clark
LEARNING OBJECTIVES • Understand the three states of matter in terms of the arrangement, movement and energy of the particles • Understand the interconversions between the three states of matter in terms of: • the names of the interconversions • how they are achieved • the changes in arrangement, movement and energy of the particles • Know what is meant by the terms: • solvent • solute • solution • saturated solution CHEMISTRY ONLY • Know what is meant by the term solubility in the units g per 100 g of solvent. • Understand how the results of experiments involving the dilution of coloured solutions and diffusion of gases can be explained CHEMISTRY ONLY The solubility of a solid in a solvent at a particular temperature is usually defined as 'the mass of solute which must dissolve in 100g of solvent at that temperature to form a saturated solution'. Figure 1.16 A saturated solution CHEMISTRY ONLY • 6 The solubility of sodium chlorate in water was measured at a number of different temperatures. • a Use these figures to plot a solubility curve, with the temperature on the horizontal axis and the solubility on the vertical one. • b Use your graph to find the solubility of sodium chlorate at 50°C. • c Determine the maximum mass of sodium chlorate that will dissolve in 40g of water at 30 °C. • d 20g of sodium chlorate was added to 100g of water and the mixture heated to about 70 °C. It was then left to cool with the thermometer in the solution. Use your graph to answer the following questions. • iAt what temperature would crystals first appear in the solution? • ii If the solution was cooled to 17 °C, calculate the total mass of crystals formed.
LEARNING OBJECTIVES • Understand how to classify a substance as an element, compound or mixture Figure 2.1 Gold is an element, but a gold ring made from 18-carat gold only contains 75% gold. The metal is a mixture of gold and, usually, copper. Figure 2.2 Pure water is a compound, but the water we drink is a mixture of water and other dissolved substances. • Understand how a chromatogram provides information about the composition of a mixture • Understand that a pure substance has a fixed melting and boiling point, but that a mixture may melt or boil over a range of temperatures • Describe these experimental techniques for the separation of mixtures: • simple distillation • fractional distillation • filtration • crystallisation • paper chromatography
LEARNING OBJECTIVES • Know what is meant by the terms atom and molecule • Know the structure of an atom in terms of the positions, relative masses and relative charges of sub-atomic particles • Know what is meant by the terms atomic number, mass number, isotopes and relative atomic mass (Aᵣ) • Be able to calculate the relative atomic mass of an element (Aᵣ) from isotopic abundances Copper is an element. If you tried to cut it up into smaller and smaller pieces, the final result would be the smallest possible piece of copper. At that stage, you would have an individual copper atom.You can, of course, split that atom into smaller pieces (protons, neutrons and electrons), but you would no longer have copper. Therefore, an atom is the smallest piece of an element that can still be recognised as that element. Figure 3.1New atoms are produced in stars . ..
LEARNING OBJECTIVES • Understand how elements are arranged in the Periodic Table: The Periodic Table shows all the elements in the universe and is one of the most important tools that a chemist has. The arrangement of the elements allows us to understand trends in properties and make predictions. The modern Periodic Table was first presented in 1869 by a famous Russian chemist, Dmitri Mendeleev (left). This chapter explores some of the features of the Periodic Table. • Identify an element as a metal or a non-metal according to its position in the Periodic Table • in order of atomic number • in groups and periods • Understand how to deduce the electronic configurations of the first 20 elements from their positions in the Periodic Table • Understand how to use electrical conductivity and the acid-base character of oxides to classify elements as metals or non-metals • Understand how the electronic configuration of a main group element is related to its position in the Periodic Table • Understand why elements in the same group of the Periodic Table have similar chemical properties • Understand why the noble gases (Group 0) do not readily react
LEARNING OBJECTIVES • Write word equations and balanced chemical equations (including state symbols): • for reactions studied in this course • for unfamiliar reactions where suitable information is provided • Calculate relative formula masses (including relative molecular masses) (M) from relative atomic masses (Aᵣ) • Calculate percentage yield • Understand how the formulae of simple compounds can be obtained experimentally, including metal oxides, water and salts containing water of crystallisation • Know what is meant by the terms empirical formula and molecular formula • Know that the mole (mol) is the unit for the amount of a substance • Calculate empirical and molecular formulae from experimental data • Understand how to carry out calculations involving amount of substance, relative atomic mass (A) and relative formula mass (M) • Practical: Know how to determine the formula of a metal oxide by combustion (e.g. magnesium oxide) or by reduction (e.g. copper(Il) oxide) • Calculate reacting masses using experimental data and chemical equations
LEARNING OBJECTIVES • Understand how to carry out calculations involving amount of substance, volume and concentration (in mol/dm³) of solution • Understand how to carry out calculations involving gas volumes and the molar volume of a gas (24 dm³ and 24 000 cm³ at room temperature and pressure (rtp))
LEARNING OBJECTIVES • Understand how ions are formed by electron loss or gain • Know the charges of these ions: • metals in Groups 1, 2 and 3 • non-metals in Groups 5, 6 and 7 • Ag⁺, Cu²⁺, Fe²⁺, Fe³⁺, Pb²⁺, Zn²⁺ • hydrogen (H⁺), hydroxide (OH⁻), ammonium (NH₄⁺), carbonate (CO₃²⁻), nitrate (NO₃⁻), sulfate (SO₄²⁻) • Write formulae for compounds formed between the ions listed above
LEARNING OBJECTIVES • Know that a covalent bond is formed between atoms by the sharing of a pair of electrons • Understand covalent bonds in terms of electrostatic attractions • Understand how to use dot-and-cross diagrams to represent covalent bonds in: • diatomic molecules, including hydrogen, oxygen, nitrogen, halogens and hydrogen halides • inorganic molecules, including water, ammonia and carbon dioxide • organic molecules containing up to two carbon atoms, including methane, ethane, ethene and those containing halogen atoms • Explain why substances with a simple molecular structure are gases or liquids, or solids with low melting and boiling points • Understand that the term intermolecular forces of attraction can be used to represent all forces between molecules • Explain why the melting and boiling points of substances with simple molecular structures increase, in general, with increasing relative molecular mass • Explain why substances with giant covalent structures are solids with high melting and boiling points • Explain how the structures of diamond, graphite and C6o fullerene influence their physical properties, including electrical conductivity and hardness • Know that covalent compounds do not usually conduct electricity
LEARNING OBJECTIVES • Know how to represent a metallic lattice by a 2D diagram • Understand metallic bonding in terms of electrostatic attractions • Explain typical physical properties of metals, including electrical conductivity and malleability
LEARNING OBJECTIVES • Understand why covalent compounds do not conduct electricity • sodium chloride, dilute sulfuric acid and copper(II) sulfate) and to predict the products • Understand why ionic compounds conduct electricity only when molten or in aqueous solution • Know that anion and cation are terms used to refer to negative and positive ions, respectively • Describe experiments to investigate electrolysis, using inert electrodes, of molten compounds (including lead sns s a ( (ading • Write ionic half-equations representing the reactions at the electrodes during electrolysis and understand why these reactions are classified as oxidation or reduction • Practical: Investigate the electrolysis of aqueous solutions
LEARNING OBJECTIVES • Understand how the similarities in the reactions of these elements with water provide evidence for their recognition as a family of elements • Understand how the differences between the reactions of these elements with air and water provide evidence for the trend in reactivity in Group 1 • Figure 11.3 The alkali metals M • Use knowledge of trends in Group 1 to predict the properties of other alkali metals CHEMISTRY ONLY • Explain the trend in reactivity in Group 1 in terms of electronic configurations. The elements in Group 1 of the Periodic Table are called the alkali metals. The group contains the elements shown in Figure 11.3. Francium (pronounced france-ee-um), at the bottom of the group, is radioactive. One of its isotopes is produced during the radioactive decay of uranium-235, but is extremely short-lived. At any one time scientists estimate that there is only about 20-30g of francium present in the whole of the Earth's crust and no one has ever seen a piece of francium. When you know about the rest of Group 1 you can predict what francium would be like. We will make those predictions later.
LEARNING OBJECTIVES • Know that chemical reactions in which heat energy is given out are described as exothermic, and those in which heat energy is taken in are described as endothermic • Describe simple calorimetry experiments for reactions such as combustion, displacement, dissolving and neutralisation • Calculate the heat energy change from a measured temperature change using the expression Q = mc∆T • Calculate the molar enthalpy change (∆H) from the heat energy change, Q • Practical: Investigate the temperature changes accompanying some of the following types of change: Practical: Investigate the temperature changes accompanying some of the following types of change: • salts dissolving in water salts dissolving in water • neutralisation reactions neutralisation reactions • displacement reactions displacement reactions • combustion reactions combustion reactions CHEMISTRY ONLY • Draw and explain energy level diagrams to represent exothermic and endothermic reactions. Draw and explain energy level diagrams to represent exothermic and endothermic reactions. • Know that bond breaking is an endothermic process and that bond making is an exothermic process. Know that bond breaking is an endothermic process and that bond making is an exothermic process. • Use bond energies to calculate the enthalpy change Use bond energies to calculate the enthalpy change during a chemical reaction. CHEMISTRY ONLY We can also work out the unknown HCl concentration. We know that there was 0.0500 mol KOH originally present. 28.00 cm³ of HCl was required for neutralisation therefore this contained 0.0500 mol of HCl: = (0 . 02800)/(0 . 02800) CHEMISTRY ONLY • b Heptane, C7H16, is a hydrocarbon found in petrol. The equation for the combustion of heptane is: C₇H₁ ₆(I) + 11O₂(g) → 7CO₂(g) + 8H₂O(I) Δ H = - 4817 k J / mo l Draw an energy level diagram to show the combustion of heptane. Label the axes and show clearly the reactants, the products and the enthalpy change of reaction on your diagram. • c Explain in terms of breaking and making bonds why the reaction in b is exothermic. END OF CHEMISTRY ONLY • 2 a Explain what is meant by an endothermic reaction. CHEMISTRY ONLY • b Nitrogen and oxygen gases can combine to form nitrogen monoxide as follows: N₂(g) + O₂(g) → 2NO(g) Δ H = + 180 kJ/mol Draw an energy level diagram for this process. Label the axes and show clearly the reactants, the products and the enthalpy change of reaction on your diagram. END OF CHEMISTRY ONLY • 3Identify each of the following changes as exothermic or endothermic. In some cases you will have to rely on your previous knowledge of chemistry. Several reactions are likely to be completely new to you. • a The reaction between sodium and water. • b Burning methane (major constituent of natural gas). • c The reaction between sodium carbonate and ethanoic acid. A thermometer placed in the reaction mixture shows a temperature drop. • f If you dissolve solid sodium hydroxide in water, the solution becomes very hot. • 4Self-heating cans are used to provide warm food in situations where it is inconvenient to use a more conventional form of heat. By doing an internet (or other) search, find out how self-heating cans work. Write a short explanation of your findings (not exceeding 200 words). You should include equation(s) for any reaction(s) involved, and a diagram or picture if it is useful. • e CuSO4(s)→CuO(s) +SO3(g) • △H=+220kJ/mol SKILLS EXECUTIVE FUNCTION CHEMISTRY ONLY • 5 Use the bond energies in the table to calculate the amount of heat released or absorbed when the following reactions take place. In each case, say whether the change is exothermic or endothermic. • a CH4(g) + Br2(g) → CH3Br(g) + HBr(g) (The structure of CH3Br is the same as that of CH3Cl) • b H2(g) + Cl2(g) → 2HCl(g) • (H2O → H2 + O2) • d N2(g) + 3H2(g) → 2NH3(g) END OF CHEMISTRY ONLY • 6A student investigated the amount of heat given out when hexane, CH14, burns using the apparatus in Figure 19.9. Hexane is a highly flammable liquid which is one of the components of petrol (gasoline). In each case, the student calculated the amount of heat evolved per mole of hexane. Her first two experiments produced answers of -3200 and -3900 kJ/mol of heat evolved. She then decided to do a third experiment. Her results were as follows: • a Suggest a reason why the student decided to do a third experiment. • b Apart from wearing eye protection, suggest two other safety precautions the student should take during the experiment.
LEARNING OBJECTIVES • Know that a hydrocarbon is a compound made up of hydrogen and carbon only • Understand how to represent organic molecules using empirical formulae, molecular formulae, general formulae, structural formulae and displayed formulae • Know what is meant by the terms homologous series, functional group and isomerism • Understand how to name compounds (containing up to six carbon atoms) using the rules of International Union of Pure and Applied Chemistry (IUPAC) nomenclature • Understand how to write the possible structural and displayed formulae of an organic molecule given its molecular formula • Understand how to classify reactions of organic compounds as substitution, addition and combustion In the next few chapters we will be looking at organic compounds. The term 'organic' was originally used because it was believed that organic compounds could only come from living things. Now it is used for any carbon compound except for the very simplest (carbon dioxide, carbon monoxide, the carbonates and the hydrogencarbonates). Organic compounds can exist as chains, branched chains or rings of carbon atoms with hydrogens attached. When you start doing organic chemistry, you are suddenly faced with a whole lot of new compounds with strange names and unfamiliar ways of drawing them. It can be quite scary! ethanol Figure 22.3 Organic chemistry involves a lot of new compounds.
LEARNING OBJECTIVES • Know the colours, physical states (at room temperature) and trends in physical properties of the halogens • Use knowledge of trends in Group 7 to predict the properties of other halogens • Figure 12.2 The halogens • Figure 12.3 lodine is a grey solid but has a purple vapour. CHEMISTRY ONLY • Explain the trend in reactivity in Group 7 in terms of electronic configurations.
LEARNING OBJECTIVES • Know the approximate percentages by volume of the four most abundant gases in dry air • Understand how to determine the percentage by volume of oxygen in air using experiments involving the reactions of metals (e.g. iron) and non-metals (e.g. phosphorus) with air • Describe the combustion of elements in oxygen, including magnesium, hydrogen and sulfur • Describe the formation of carbon dioxide from the thermal decomposition of metal carbonates, including copper(II) carbonate • Know that carbon dioxide is a greenhouse gas and that increasing amounts in the atmosphere may contribute to climate change • Practical: Determine the approximate percentage by volume of oxygen in air using a metal or a non-metal
LEARNING OBJECTIVES • Understand how metals can be arranged in a reactivity series based on their reactions with: • water • dilute hydrochloric or sulfuric acid Figure 14.2 Gold is so unreactive that it will remain chemically unchanged in contact with air or water basicallyforever. • Understand how the rusting of iron may be prevented by • barrier methods • galvanising • sacrificial protection • Understand how metals can be arranged in a reactivity series based on their displacement reactions between: • metals and metal oxides • Understand the terms: • oxidation • reduction • metals and aqueous solutions of metal salts • redox • Know the order of reactivity of these metals: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver, gold • Know the conditions under which iron rusts • oxidising agent • reducing agent in terms of gain or loss of oxygen and loss or gain of electrons • Practical: Investigate reactions between dilute hydrochloric and sulfuric acids and metals (e.g. magnesium, zinc and iron) CHEMISTRY ONLY Carbon is included in the reactivity series because it is important in extracting several metals (including iron) from metal oxides. If the metal is less reactive than carbon (below carbon in the reactivity series), then heating with carbon can be a cheap way of removing oxygen from the oxide to leave the metal. Copper isn't, in fact, extracted like this. This reaction is simply a lab illustration that carbon is above copper in the reactivity series. END OF CHEMISTRY ONLY
LEARNING OBJECTIVES • Know that most metals are extracted from ores found in the Earth's crust and that unreactive metals are often found as the uncombined element • Explain how the method of extraction of a metal is related to its position in the reactivity series, illustrated by carbon extraction for iron and electrolysis for aluminium • Be able to comment on a metal extraction process, given appropriate information (detailed knowledge of the processes used in the extraction of a specific metal is not required)
LEARNING OBJECTIVES • Describe the use of litmus, phenolphthalein and methyl orange to distinguish between acidic and alkaline solutions • Know that acids in aqueous solution are a source of hy n sls sn an a sl n s ra source of hydroxide ions • Understand how the pH scale, from 0–14, can be used to classify solutions as strongly acidic (0-3), weakly acidic (4–6), neutral (7), weakly alkaline (8–10) and strongly alkaline (11–14) • Describe the use of universal indicator to measure the approximate pH value of an aqueous solution • Know that alkalis can neutralise acids CHEMISTRY ONLY • Describe how to carry out an acid-alkali titration. END OF CHEMISTRY ONLY • State the colours of methyl orange, phenolphthalein and litmus in each of the following solutions: • a sodium hydroxide solution • b hydrochloric acid solution. • 2 Table 16.5 gives the pH of some solutions. Classify each as strongly acidic, strongly alkaline, weakly acidic, weakly alkaline or neutral by copying the table and putting ticks in the appropriate boxes. Table 16.5 • 3Write equations for the reactions between: • a sodium oxide and nitric acid b potassium hydroxide and sulfuric acid. • Explain what type of reaction is occurring in 3. CHEMISTRY ONLY • 5 A student carried out an acid-alkali titration. They used the following set of instructions. Copy out the paragraph and fill in the blanks. Measure out 25.0 cm³ of potassium hydroxide solution using a Transfer the potassium hydroxide solution to a . Add a few drops of an Put sulfuric acid into the Add the acid to the alkali until the END OF CHEMISTRY ONLY
LEARNING OBJECTIVES • Know the general rules for predicting the solubility of ionic compounds in water: • common sodium, potassium and ammonium compounds are soluble • all nitrates are soluble • common chlorides are soluble, except those of silver and lead(II) • common hydroxides are insoluble except for those of sodium, potassium and calcium (calcium hydroxide is slightly soluble) • Understand acids and bases in terms of proton transfer • Understand that an acid is a proton donor and a base is a proton acceptor • Describe the reactions of hydrochloric acid, sulfuric acid and nitric acid with metals, bases and metal carbonates (excluding the reactions between nitric acid and metals) to form salts CHEMISTRY ONLY • Describe an experiment to prepare a pure, dry sample of a soluble salt, starting from an acid and alkali. • Describe an experiment to prepare a pure, dry sample of an insoluble salt, starting from two soluble reactants. • Practical: Prepare a sample of pure, dry lead(II) sulfate CHEMISTRY ONLY • 6 a Read the following description of a method for making sodium sulfate crystals, Na2SO4.10H2O, and then explain the reasons for each of the underlined phrases or sentences. 25 cm³ of sodium carbonate solution was transferred to a conical flask using a pipette, and a few drops of methyl orange were added. Dilute sulfuric acid was run in from a burette until the solution became orange. The volume of acid added was noted. That same volume of dilute sulfuric acid was added to a fresh 25 cm³ sample of sodium carbonate solution in a clean flask, but without the methyl orange. The mixture was evaporated until a sample taken on the end of a glass rod crystallised on cooling in the air. The solution was left to cool. The crystals formed were separated from the remaining solution and dried. END OF CHEMISTRY ONLY • 10 All of the following equations represent acid-base reactions. For each of the equations, state which substance is the acid and which the base. • a CO3²⁻(s) + H₂O(l) → HCO₃⁻(aq) + H₂O(l) • b MgO(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂O(l) • c HNO₃(aq) + NH₃(aq) → NH₄⁺(aq) + NO₃⁻(aq) • d H₂SO₄(aq) + H₂O(l) → HSO₄⁻(aq) + H₃O⁺(aq) • e NH₃(g) + HCl(g) → NH₄Cl(s) • f CH₃COOH(aq) + NH₃(aq) → CH₃COO⁻(aq) + NH₄⁺(aq)
LEARNING OBJECTIVES • Describe tests for these gases: • hydrogen • oxygen • carbon dioxide • ammonia • chlorine • Describe how to carry out a flame test • Know the colours formed in flame tests for these cations: • Li+ (red) • Na+ (yellow) • K+ (lilac) • Ca²+ (orange-red) • Cu²+ (blue-green) • Describe tests for these cations: • NH4+ using sodium hydroxide solution and identifying the gas evolved • Cu²+, Fe²+ and Fe³+ using sodium hydroxide solution • Describe tests for these anions: • Cl-, Br- and I- using acidified silver nitrate solution • SO42- using acidified barium chloride solution • CO32- using hydrochloric acid and identifying the gas evolved • Describe a test for the presence of water using anhydrous copper(Il) sulfate • Describe a physical test to show whether or not a sample of water is pure
LEARNING OBJECTIVES • Describe experiments to investigate the effects of changes in the surface area of a solid, the concentration of a solution, the temperature and the use of a catalyst on the rate of a reaction • Describe the effects of changes in the surface area of a solid, the concentration of a solution, the pressure of a gas, the temperature and the use of a catalyst on the rate of a reaction • Explain the effects of changes in the surface area of a solid, the concentration of a solution, the pressure of a gas and the temperature on the rate of a reaction in terms of particle collision theory • Know that a catalyst is a substance that increases the rate of a reaction, but is chemically unchanged at the end of the reaction • Figure 20.4 The weathering of limestone and the formation of stalagmites and stalactites takes a very long time. • Know that a catalyst works by providing an alternative pathway with lower activation energy CHEMISTRY ONLY • Draw and explain reaction profile diagrams showing △H and activation energy. • Practical: Investigate the effect of changing the surface area of marble chips and of changing the concentration of hydrochloric acid on the rate of reaction between marble chips and dilute hydrochloric acid • Practical: Investigate the effect of different solids on the catalytic decomposition of hydrogen peroxide solution CHEMISTRY ONLY You can show the effect of a catalyst on a reaction profile diagram (Figure 20.25 and Figure 20.26). Adding a catalyst gives the reaction an alternative way to happen that involves a lower activation energy. Figure 20.25 A reaction profile diagram END OF CHEMISTRY ONLY Cat a e telp substances to react quickly at lower temperatures and pressures than would otherwise be needed. This saves money. • 1 A student carried out an experiment to investigate the rate of a reaction between an excess of dolomite (magnesium carbonate) and 50 cm³ of dilute hydrochloric acid. The dolomite was in small pieces. The reaction is: Mg CO₃(s) + 2HCl (aq) → Mg Cl₂(aq) + H₂O(l) + CO₂(g) He measured the volume of carbon dioxide given off at regular intervals, with the results shown in the table below. • Draw a diagram of the apparatus you would use for this experiment and label it. • b Plot these results on a piece of graph paper, with time on the x-axis and volume of gas on the y-axis. • c At what time is the gas being given off most quickly? Explain why the reaction is fastest at that time. • d Use your graph to find out how long it took to produce 50 cm³ of gas. • e Use your graph to calculate the average rate of this reaction in the first 80 seconds. • f In each of the following questions, state and explain what would happen to the initial rate of the reaction and to the total volume of gas given off if various changes were made to the experiment. • i The mass of dolomite and the volume and concentration of acid were kept constant, but the dolomite was in one big lump instead of small pieces. • ii The mass of dolomite was unchanged and it was still in small pieces. 50cm³ of hydrochloric acid was used, which had half the original concentration. • iii The dolomite was unchanged again. This time 25 cm³ of the original acid was used instead of 50 cm³. • iv The acid was heated to 40 °C before the dolomite was added to it. • 2 The effect of concentration and temperature on the rate of a reaction can be explored using the reaction between magnesium ribbon and dilute sulfuric acid: Mg (s) + H₂SO₄(aq) → Mg SO₄(aq) + H₂(g) A student dropped a 2 cm length of magnesium ribbon into 25 cm³ of dilute sulfuric acid (a large excess of acid) in a boiling tube. She stirred the contents of the tube continuously and timed how long it took for the magnesium to disappear. • a What would you expect to happen to the time taken for the reaction if the student repeated the experiment using the same length of magnesium with a mixture of 20 cm³ of acid and 5 cm³ of water? Explain your answer in relation to the collision theory. • b What would you expect to happen to the time taken for the reaction if she repeated the experiment using the original quantities of magnesium and acid, but first heated the acid to 50 °C? Explain your answer in relation to the collision theory. • C Give a reason why is it important to keep the reaction mixture stirred continuously. CHEMISTRY ONLY • d The reaction between magnesium ribbon and dilute sulfuric acid is exothermic. Draw a reaction profile diagram for this reaction, label the axes clearly and show both the activation energy and enthalpy change on your diagram. END OF CHEMISTRY ONLY • 3Catalysts speed up reactions, but can be recovered chemically unchanged at the end of the reaction. • a Explain how a catalyst works to increase the rate of reactions. • b Describe how you would find out whether iron(Ill) oxide was a catalyst for the decomposition of hydrogen peroxide solution. Note: you need to show not only that it speeds the reaction up, but that it is chemically unchanged at the end. CHEMISTRY ONLY • c Draw a clearly labelled reaction profile diagram showing the effect of usinga catalyst in a reaction. END OF CHEMISTRY ONLY
LEARNING OBJECTIVES • Know that some reactions are reversible and that this is indicated by the symbol ⇒ in equations • Describe reversible reactions such as the dehydration of hydrated copper(II) sulfate and the effect of heat on ammonium chloride CHEMISTRY ONLY • Know that a reversible reaction can reach dynamic equilibrium in a sealed container. • Know that the characteristics of a reaction at dynamic equilibrium are: • the forward and reverse reactions occur at the same rate • the concentrations of reactants and products remain constant. END OF CHEMISTRY ONLY SKILLS CRITICAL THINKING SKILLS SKILLS REASONING REASONING CHEMISTRY ONLY • 2 For the following reactions at equilibrium, predict in which direction (Left or Right or No change) the position of equilibrium might shift if the conditions were changed: Change in conditions N2O4(g)=2NO2(g) 2NH3(g) =N2(g) + 3H2(g) Endothermic Increase in temperature Endothermic Increase in pressure N2(g) + O2(g) = 2NO(g) Endothermic Decrease in temperature )()+()() Endothermic Decrease in pressure CO(g)+ 2H2(g) =CH3OH(g)Exothermic Adding a catalyst • 3An important stage in the Contact process for the production of sulfuric acid involves a reversible reaction: 2SO₂(g) + O₂(g)leftharpoons 2SO₃(g) Δ H = - 196 k J / mo l At 450 °C the reaction mixture consists of a dynamic equilibrium involving sulfur dioxide, oxygen and sulfur trioxde. • a Explain what is meant by the term dynamic equilibrium. Be sure that you have explained what both of the words mean. • b Predict and explain the effect of an increase in pressure on the proportion of the sulfur trioxide present in the equilibrium mixture. • c Predict and explain the effect of lowering the temperature of the mixture on the position of equilibrium of the reaction. • 4Hydrogen can be made by the reaction between methane (natural gas) and steam. The reaction can be carried out by passing a mixture of methane and steam over a nickel catalyst at pressures between 2 and 30 atmospheres and a temperature of about 1000 °C. CH₄(g) + H₂O(g)leftharpoons CO(g) + 3H₂(g) Δ H = + 210 k J / mo l • a The pressure used is relatively low. What would be the effect on the conversion of the methane into carbon monoxide and hydrogen if the pressure was higher? • b Explain why a high temperature is used to get a good conversion of methane into hydrogen. • c Explain the use of the nickel catalyst in the production of hydrogen, referring to both the rate of the reaction and the position of equilibrium. • 5 Ammonia, NH3, is manufactured by passing a mixture of nitrogen and hydrogen over an iron catalyst at a pressure of 200 atmospheres or more, and a temperature of 450 °C. • a Explain why this reaction will produce a higher percentage conversion into ammonia if the pressure is very high. • b 200 atmospheres is a high pressure, but not very high. Give a reason why most ammonia manufacturers don't use a pressure of, for example, 1000 atmospheres. (Hint: think about the risk of using high pressure.) • c Predict and explain whether you would get the best yield of ammonia in the equilibrium mixture at a low or a high temperature. • d The temperature used, 450 °C, is neither very high nor very low. Suggest a reason why a manufacturer might choose a temperature which gave less than an ideal percentage conversion. (Hint: think about rates of reaction.)
LEARNING OBJECTIVES • Know that crude oil is a mixture of hydrocarbons • Describe how the industrial process of fractional distillation separates crude oil into fractions • Know the names and uses of the main fractions obtained from crude oil: refinery gases, gasoline, kerosene, diesel, fuel oil and bitumen • Know the trend in colour, boiling point and viscosity of the main fractions • Know that a fuel is a substance that, when burned, releases heat energy • Know the possible products of complete and incomplete combustion of hydrocarbons with oxygen in the air • Know that in car engines the temperature reached is high enough to allow nitrogen and oxygen from air to react, forming oxides of nitrogen • Explain how the combustion of some impurities in hydrocarbon fuels results in the formation of sulfur dioxide • Understand how sulfur dioxide and oxides of nitrogen contribute to acid rain • Describe how long-chain alkanes are converted to alkenes and shorter-chain alkanes by catalytic cracking (using silica or alumina as the catalyst and a temperature in the range of 600–700 °C) • Explain why cracking is necessary in terms of the balance between supply and demand for different fractions
LEARNING OBJECTIVES • Know the general formula for alkanes Know the general formula for alkanes • Explain why alkanes are classified as saturated hydrocarbons Explain why alkanes are classified as saturated hydrocarbons • Figure 24.1 The small alkanes are gases and are burned as fuels.
LEARNING OBJECTIVES • Know that alkenes contain the functional group >C=C< • Know the general formula for alkenes • Explain why alkenes are classified as unsaturated hydrocarbons • Describe the reactions of alkenes with bromine to produce dibromoalkanes • Describe how bromine water can be used to distinguish between an alkane and an alkene The alkenes are a homologous series of hydrocarbons which contain a carbon-carbon double bond. The C=C bond is the functional group of the alkenes. The first three members of the alkene series are shown in Table 25.1.
LEARNING OBJECTIVES • Know that alcohols contain the functional group -OH • Know that ethanol can be manufactured by: • Understand how to draw structural and displayed formulae for methanol, ethanol, propanol (propan-1-ol only) and butanol (butan-1-ol only), and name each compound • Know that ethanol can be oxidised by: • burning in air or oxygen (complete combustion) • reaction with oxygen in the air to form ethanoic acid (microbial oxidation) • heating with potassium dichromate(Vl) in dilute sulfuric acid to form ethanoic acid What we commonly refer to as alcohol' is just one member of a large family (homologous series) of similar compounds. Alcohols all contain the functional group -OH covalently bonded to a carbon chain. The familiar alcohol in drinks is C2H5OH (or, better, showing the structure, CH3CH2OH) and should properly be called ethanol. • reacting ethene with steam in the presence of a phosphoric acid catalyst at a temperature of about 300 °C and a pressure of about 60–70 atm • the fermentation of glucose, in the absence of air, at an optimum temperature of about 30 °C and using the enzymes in yeast • Understand the reasons for fermentation in the absence of air and at an optimum temperature
LEARNING OBJECTIVES • Know that carboxylic acids contain the functional group -OH • Understand how to draw structural and displayed formulae for unbranched-chain carboxylic acids with up to four carbon atoms in the molecule, and name each compound Acids such as ethanoic acid are known as carboxylic acids and they all contain the functional group -CoOH The most familiar one is ethanoic acid (old name acetic acid). Vinegar is an aqueous solution containing ethanoic acid. Carboxylic acids are formed by oxidation of the corresponding alcohols, for example ethanoic acid is formed when ethanol is left in the air or oxidised using potassium dichromate(Vl) in dilute sulfuric acid. If you need a reminder, have at look at Chapter 26, pages 289–290. • Describe the reactions of aqueous solutions of carboxylic acids with metals and metal carbonates • Know that vinegar is an aqueous solution containing ethanoic acid
LEARNING OBJECTIVES • Know that esters contain the functional group • Know that ethyl ethanoate is the ester produced when ethanol and ethanoic acid react in the presence of an acid catalyst • Understand how to write the structural and displayed formulae of an ester, given the name or formula of the alcohol and carboxylic acid from which it is formed and vice versa • Know that esters are volatile compounds with distinctive smells and are used as food flavourings and in perfumes • Understand how to write the structural and displayed formulae of ethyl ethanoate • Practical: Prepare a sample of an ester such as ethyl ethanoate Esters are organic compounds formed by the reaction of an alcohol with a carboxylic acid. They have the functional group This can be written in condensed form as -Coo-.
LEARNING OBJECTIVES • Know that an addition polymer is formed by joining up many small molecules called monomers • Explain problems in the disposal of addition polymers, including: • Understand how to draw the repeat unit of an addition polymer, including poly(ethene), poly(propene), poly(chloroethene) and poly(tetrafluoroethene) • Understand how to deduce the structure of a monomer from the repeat unit of an addition polymer and vice versa CHEMISTRY ONLY • Krn d a sd o a s w a e s o do wter. • Understand how to write the structural and displayed formula of a polyester, showing the repeat unit, given the formulae of the monomers from which it is formed, including the reaction of ethanedioic acid and ethanediol: • Know that some polyesters, known as biopolyesters, are biodegradable. CHEMISTRY ONLY • 3A polymer has the structure shown below: SKILLS INTERPRETATION • a Identify the type of the polymer. • b Draw the structures of the monomers that are used to make this polymer. SKILLS INTERPRETATION PROBLEM SOLVING CRITICAL THINKING INTERPRETATION • 4 Polyesters such as Terylene (for clothes) or PET (commonly used to make drinks bottles) are made by condensation polymerisation from ethane-1,2diol and terephthalic acid (properly known as benzene-1,4-dicarboxylic acid). The structures of these are: • a Draw a chain of the polymer with two repeat units. • b Write a balanced chemical equation for the formation of PET. Ethane-1,2-diol: Terephthalic acid: HO-CH_2CH_2-OH (The structure of the CH4 group is complicated, and you can write it simply as C₆H4. You can draw this as a block diagram if you wish, but it is more satisfying (and no more difficult in this case) to draw the structure properly.) • 5(This question contains new material and is designed to look difficult. It is actually not too difficult as long as you understand about polyester.) Nylon-6,6 is made by a condensation polymerisation of the monomers 1,6-diaminohexane (H2NCH2CH2CH2CH2CH2CH2NH₂) and hexane-1,6-dioic acid. The amine group -NH2 reacts in a very similar manner to the -OH gron n od on n n n ic acids, forming amide functional groups which join the monomers together into a polymer chain. • Explain what is meant by condensation polymerisation and how it differs from addition polymerisation. • ii Using a block diagram, draw one repeat unit for nylon-6,6. SKILLS ANALYSIS, REASONING • b Nylon-6,10 is made from 1,6-diaminohexane and a longer chain acid, decanedioic acid, containing a total of 10 carbon atoms: HOOCCH₂CH₂CH₂CH₂CH₂CH₂CH2CH₂COOH. • How will a chain of nylon-6,10 differ from one of nylon-6,6? Refer to the diagram you drew in a ii. • ii In what way(s) will the two chains be the same? Again, refer to the diagram you drew in a ii. END OF CHEMISTRY ONLY