IGCSE Chemistry CIE 9. Metals PDF

Head to savemyexams.co.uk for more awesome resources YOUR NOTES IGCSE Chemistry CIE  9. Metals CONTENT...

Head to savemyexams.co.uk for more awesome resources YOUR NOTES IGCSE Chemistry CIE  9. Metals CONTENTS 9.1 Properties, Uses & Alloys of Metals 9.1.1 Properties of Metals 9.1.2 Uses of Metals 9.1.3 Alloys 9.2 Reactivity Series & Corrosion of Metals 9.2.1 Reactivity Series 9.2.2 Explaining Reactivity 9.2.3 Rusting of Iron 9.2.4 Galvanising & Sacrificial Protection 9.3 Extraction of Metals 9.3.1 Extraction of Metals 9.3.2 Extraction of Iron from Hematite 9.3.3 Extraction of Aluminium from Bauxite Page 1 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 9.1 Properties, Uses & Alloys of Metals YOUR NOTES  9.1.1 Properties of Metals Physical Properties of Metals & Non-Metals Metals and non-metals The Periodic Table contains over 100 different elements They can be divided into two broad types: metals and non-metals Most of the elements are metals and a small number of elements display properties of both types These elements are called metalloids or semimetals The metallic character diminishes moving left to right across the Periodic Table Properties of metals Conduct heat and electricity Are malleable (can be hammered and made into different shapes) and ductile (can be drawn into wires) Page 2 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Tend to be lustrous (shiny) YOUR NOTES Have high density and usually have high melting points  Form positive ions through electron loss Form basic oxides Properties of non-metal elements Do not conduct heat and electricity Are brittle when solid and easily break up Tend to be dull and nonreflective Have low density and low melting points (many are gases at room temperature) Form negative ions through electron gain (except for hydrogen) Form acidic oxides Page 3 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Chemical Properties of Metals YOUR NOTES General chemical properties of metals  The chemistry of metals is studied by analysing their reactions with water, dilute acid and oxygen Based on these reactions, a reactivity series of metals can be produced Reactivity with water Some metals react with water, either warm or cold, or with steam Metals that react with cold water form a metal hydroxide and hydrogen gas metal + water → metal hydroxide + hydrogen For example calcium: Ca (s) + 2H2O (l) → Ca(OH)2 (aq) + H2 (g) Metals that react with steam form metal oxide and hydrogen gas, for example zinc: Zn (s) + H2O (g) → ZnO (s) + H2 (g) Reactivity with acids Most metals react with dilute acids such as HCl When acids and metals react, the hydrogen atom in the acid is replaced by the metal atom to produce a salt and hydrogen gas, for example iron: metal + acid → salt + hydrogen Fe (s) + 2HCl (aq) → FeCl2 (aq) + H2 (g) Reactivity with oxygen Unreactive metals such as gold and platinum do not react with oxygen Some reactive metals such as the alkali metals react easily with oxygen Copper and iron can also react with oxygen although much more slowly When metals react with oxygen a metal oxide is formed, for example copper: metal + oxygen → metal oxide 2Cu (s) + O2 (g) → 2CuO (s) Page 4 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 9.1.2 Uses of Metals YOUR NOTES  Uses of Metals Uses of Aluminium Uses of Copper Page 5 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 9.1.3 Alloys YOUR NOTES  Properties & Uses of Alloys An alloy is a mixture of two or more metals or metal with a non-metal such as carbon Alloys often have properties that can be very different from the metals they contain, for example, they can have more strength, hardness or resistance to corrosion or extreme temperatures These enhanced properties can make alloys more useful than pure metals The regular arrangement of a metal lattice structure is distorted in alloys Common alloys and their uses Brass is an alloy of copper and zinc and is much stronger than either metal It is used in musical instruments, ornaments and door knobs Stainless steel is a mixture of iron and other elements, for example, chromium, nickel and carbon It is used in cutlery because of its hardness and resistance to corrosion Alloys of iron with tungsten are extremely hard and resistant to high temperatures Alloys of iron mixed with chromium or nickel are resistant to corrosion Aluminium is mixed with copper, manganese and silicon for aircraft body production as the alloy is stronger but still has a low density  Exam Tip Alloys are mixtures of substances, they are not chemically combined and an alloy is not a compound. Page 6 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Explaining the Properties of Alloys YOUR NOTES EXTENDED  Alloys contain atoms of different sizes, which distorts the normally regular arrangements of atoms in metals This makes it more difficult for the layers to slide over each other, so alloys are usually much harder than the pure metal Page 7 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 9.2 Reactivity Series & Corrosion of Metals YOUR NOTES  9.2.1 Reactivity Series Reactivity Series The chemistry of the metals is studied by analysing their reactions with water and acids Based on these reactions a reactivity series of metals can be produced The series can be used to place a group of metals in order of reactivity based on the observations of their reactions with water and acids The non-metals hydrogen and carbon are also included in the reactivity series as they are used to extract metals from their oxides Table of Metal Reactions Page 8 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources YOUR NOTES  The reactivity series mnemonic Observations from the table above allow the following reactivity series to be deduced The order of this reactivity series can be memorised using the following mnemonic “Please send cats, monkeys and cute zebras into hot countries signed Gordon" Page 9 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources YOUR NOTES  You can learn the reactivity series with the help of a silly phrase Reactions of Metals Reaction with cold water The more reactive metals will react with cold water to form a metal hydroxide and hydrogen gas Potassium, sodium and calcium all undergo reactions with cold water as they are the most reactive metals: metal + water → metal hydroxide + hydrogen For example, calcium and potassium: Ca (s) + 2H2O (l) → Ca(OH)2 (aq) + H2 (g) K (s) + H2O (l) → KOH (aq) + H2 (g) Reaction with steam Page 10 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Metals just below calcium in the reactivity series do not react with cold water but will react YOUR NOTES with steam to form a metal oxide and hydrogen gas, for example, magnesium:  Mg (s) + H2O (g) → MgO (s) + H2 (g) Reaction with dilute acids Only metals above hydrogen in the reactivity series will react with dilute acids Unreactive metals below hydrogen, such as gold, silver and copper, do not react with acids The more reactive the metal then the more vigorous the reaction will be Metals that are placed high on the reactivity series such as potassium and sodium are very dangerous and react explosively with acids When acids react with metals they form a salt and hydrogen gas: The general equation is: metal + acid ⟶ salt + hydrogen Some examples of metal-acid reactions and their equations are given below: Acid-Metal Reactions Table Reaction with oxygen Some reactive metals, such as the alkali metals, react easily with oxygen Silver, copper and iron can also react with oxygen although much more slowly When metals react with oxygen a metal oxide is formed, for example, copper: metal + oxygen → metal oxide 2Cu (s) + O2 (g) → 2CuO (s) Gold does not react with oxygen Deducing the order of reactivity The order of reactivity of metals can be deduced by making experimental observations of reactions between metals and water, acids and oxygen The more vigorous the reaction of the metal, the higher up the reactivity series the metal is A combination of reactions may be needed, for example, the order of reactivity of the more reactive metals can be determined by their reactions with water The less reactive metals react slowly or not at all with water, so the order of reactivity would need to be determined by observing their reactions with dilute acid Page 11 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Temperature change in a reaction can also be used to determine the order of reactivity YOUR NOTES The greater the temperature change in a reaction involving a metal, the more reactive the  metal is Page 12 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 9.2.2 Explaining Reactivity YOUR NOTES  Explaining Reactivity EXTENDED Metal atoms form positive ions by loss of electrons when they react with other substances The tendency of a metal to lose electrons is a measure of how reactive the metal is A metal that is high up on the series loses electrons easily and is thus more reactive than one which is lower down on the series Displacement reactions between metals and aqueous solutions of metal salts Any metal will displace another metal that is below it in the reactivity series from a solution of one of its salts This is because more reactive metals lose electrons and form ions more readily than less reactive metals, making them better reducing agents The less reactive metal is a better electron acceptor than the more reactive metal, thus the less reactive metal is reduced. (OIL-RIG: reduction is gain of electrons)l Example: Magnesium + copper sulfate Magnesium is a reactive metal and can displace copper from a copper sulfate solution Magnesium loses its electrons more easily and the ion of the less reactive metal, copper, will gain these electrons to form elemental copper This is easily seen as the more reactive metal slowly disappears from the solution, displacing the less reactive metal magnesium + copper sulfate → magnesium sulfate + copper Mg (s) + CuSO4 (aq) → MgSO4 (aq) + Cu (s) The blue colour of the CuSO4 solution fades as colourless magnesium sulfate solution is formed Copper coats the surface of the magnesium and also forms solid metal which falls to the bottom of the beaker Page 13 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources YOUR NOTES  Diagram showing the colour change when magnesium displaces copper from copper sulfate By combining different metals and metal salts solutions it is possible to come up with a relative reactivity order Metal Solutions Displacement Table From this table we can deduce the order of reactivity: Magnesium and zinc are both more reactive than iron but magnesium is more reactive than zinc Copper and silver are both less reactive than iron but silver is less reactive than copper The order of reactivity of the metals tested can be therefore be deduced as: Page 14 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Mg > Zn > Fe > Cu > Ag YOUR NOTES Reactivity of aluminium  Aluminium is high in the reactivity series, but in reality, it does not react with water and the reaction with dilute acids can be quite slow This is because it reacts readily with oxygen, forming a protective layer of aluminium oxide which is very thin This layer prevents reaction with water and dilute acids, so aluminium can behave as if it is unreactive Page 15 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 9.2.3 Rusting of Iron YOUR NOTES  Rusting of Iron Rust is a chemical reaction between iron, water and oxygen to form the compound hydrated iron(III) oxide (rust) Oxygen and water must be present for rust to occur During rusting, iron is oxidised iron + water + oxygen → hydrated iron(III) oxide Investigating rusting To investigate the conditions required for rusting, prepare three test tubes as shown in the diagram The oil in the 2nd tube keeps out air and the water has been boiled so that no air is left in it The calcium chloride in the 3rd tube is used to remove any moisture in the air After a few days, the iron nail in the 1st tube will be the only nail to show signs of rust Diagram showing the requirements of oxygen and water for rust to occur: only the nail on the left rusts Rust prevention methods Barrier methods Rust can be prevented by coating iron with barriers that prevent the iron from coming into contact with water and oxygen However, if the coatings are washed away or scratched, the iron is once again exposed to water and oxygen and will rust Page 16 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources YOUR NOTES   Exam Tip Only iron or steel (an alloy made from iron) can rust. If any other metal oxidises in air causing the metal to break down, you should say that the metal has corroded. Page 17 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 9.2.4 Galvanising & Sacrificial Protection YOUR NOTES  Galvanising & Sacrificial Protection EXTENDED Iron can be prevented from rusting using the reactivity series Sacrificial Protection A more reactive metal can be attached to a less reactive metal The more reactive metal will oxidise and therefore corrode first, protecting the less reactive metal from corrosion E.g. using zinc bars on the side of steel ships: Diagram to show the use of zinc bars on the sides of steel ships as a method of sacrificial protection Zinc is more reactive than iron therefore will lose its electrons more easily than iron and is oxidised more easily: Zn → Zn2+ + 2e- The iron is less reactive therefore will not lose its electrons as easily so it is not oxidised; the zinc is sacrificed to protect the steel For continued protection, the zinc bars have to be replaced before they completely corrode Galvanising Galvanising is a process where the iron to be protected is coated with a layer of zinc This can be done by electroplating or dipping it into molten zinc ZnCO3 is formed when zinc reacts with oxygen and carbon dioxide in the air and protects the iron by the barrier method If the coating is damaged or scratched, the iron is still protected from rusting by sacrificial protection Page 18 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Exam Tip YOUR NOTES  You maybe asked to explain why a metal is/is not suitable as a method of preventing  an iron/steel object from rusting. Remember that if it is higher in the reactivity series than iron, it will be suitable for sacrificial protection as it will be oxidised instead of iron. If it is lower in the reactivity series than iron, it would not be suitable as iron would be oxidised, causing it to rust. Page 19 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 9.3 Extraction of Metals YOUR NOTES  9.3.1 Extraction of Metals Extraction of Metals The Earth’s crust contains metals and metal compounds such as gold, copper, iron oxide and aluminium oxide Useful metals are often chemically combined with other substances forming ores A metal ore is a rock that contains enough of the metal to make it worthwhile extracting They have to be extracted from their ores through processes such as electrolysis, using a blast furnace or by reacting with more reactive material In many cases the ore is an oxide of the metal, therefore the extraction of these metals is a reduction process since oxygen is being removed Common examples of oxide ores are iron and aluminium ores which are called hematite and bauxite respectively Unreactive metals do not have to be extracted chemically as they are often found as the uncombined element This occurs as they do not easily react with other substances due to their chemical stability They are known as native metals and examples include gold and platinum which can both be mined directly from the Earth’s crust The position of the metal on the reactivity series influences the method of extraction Those metals placed higher up on the series (above carbon) have to be extracted using electrolysis Metals lower down on the series can be extracted by heating with carbon The Extraction Method Depends on the Position of a Metal in the Reactivity Series Page 20 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources YOUR NOTES  Page 21 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 9.3.2 Extraction of Iron from Hematite YOUR NOTES  Extraction of Iron from Hematite Iron is extracted in a large container called a blast furnace from its ore, hematite Modern blast furnaces produce approximately 10,000 tonnes of iron per day The process is demonstrated and explained below: Diagram showing the carbon extraction of iron The raw materials: iron ore (hematite), coke (an impure form of carbon), and limestone are added into the top of the blast furnace Hot air is blown into the bottom Zone 1: Coke burns in the hot air forming carbon dioxide The reaction is exothermic so it gives off heat, heating the furnace Page 22 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources carbon + oxygen → carbon dioxide YOUR NOTES Zone 2:  At the high temperatures in the furnace, more coke reacts with carbon dioxide forming carbon monoxide Carbon dioxide has been reduced to carbon monoxide carbon + carbon dioxide → carbon monoxide Zone 3: Carbon monoxide reduces the iron(III) oxide in the iron ore to form iron This will melt and collect at the bottom of the furnace, where it is tapped off: iron(III) oxide + carbon monoxide → iron + carbon dioxide Limestone (calcium carbonate) is added to the furnace to remove impurities in the ore. The calcium carbonate in the limestone thermally decomposes to form calcium oxide calcium carbonate → calcium oxide + carbon dioxide The calcium oxide formed reacts with the silicon dioxide, which is an impurity in the iron ore, to form calcium silicate This melts and collects as a molten slag floating on top of the molten iron, which is tapped off separately calcium oxide + silicon dioxide → calcium silicate  Exam Tip For Core students, the symbol equations are not needed for the different reactions involved in the extraction of iron from hematite. Page 23 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources Equations for Extraction of Iron from Hematite YOUR NOTES EXTENDED  The symbol equations for the different stages of the extraction of iron from hematite are: Zone 1: The burning of carbon (coke) to provide heat and produce carbon dioxide: C (s) + O2 (g) → CO2 (g) Zone 2: The reduction of carbon dioxide to carbon monoxide: CO2 (g) + C (s) → 2CO (g) Zone 3: The reduction of iron(III) oxide by carbon monoxide: Fe2O3 (s) + 3CO (g) → 2Fe (I) + 3CO2 (g) The thermal decomposition of calcium carbonate (limestone) to produce calcium oxide: CaCO3 (s) → CaO (s) + CO2 (g) The formation of slag: CaO (s) + SiO2 (s) → CaSiO3 (l) Page 24 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources 9.3.3 Extraction of Aluminium from Bauxite YOUR NOTES  Extraction of Aluminium from Bauxite Aluminium is a reactive metal, above carbon in the reactivity series Its main ore, is bauxite, which contains aluminium oxide Aluminium is higher in the reactivity series than carbon, so it cannot be extracted by reduction using carbon Instead, aluminium is extracted by electrolysis Diagram showing the extraction of aluminium by electrolysis  Exam Tip If you are a Core student, you do not need to explain the process of extraction of aluminium by electrolysis. Page 25 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to savemyexams.co.uk for more awesome resources The Process of Aluminium Extraction by Electrolysis YOUR NOTES EXTENDED  Bauxite is first purified to produce aluminium oxide, Al2O3 Aluminium oxide is then dissolved in molten cryolite This is because aluminium oxide has a melting point of over 2000°C which would use a lot of energy and be very expensive The resulting mixture has a lower melting point without interfering with the reaction The mixture is placed in an electrolysis cell, made from steel, lined with graphite The graphite lining acts as the negative electrode, with several large graphite blocks as the positive electrodes At the cathode (negative electrode): Aluminium ions gain electrons (reduction) Molten aluminium forms at the bottom of the cell The molten aluminium is siphoned off from time to time and fresh aluminium oxide is added to the cell Al3+ + 3e- → Al At the anode (positive electrode): Oxide ions lose electrons (oxidation) Oxygen is produced at the anode: 2O2- → O2 + 4e- The overall equation for the reaction is: 4Al + 3O2 → 2Al2O3 The carbon in the graphite anodes reacts with the oxygen produced to produce CO2 C (s) + O2 (g) → CO2 (g) As a result the anode wears away and has to be replaced regularly A lot of electricity is required for this process of extraction, this is a major expense  Exam Tip Use OIL RIG to remember whether oxidation or reduction has occurred at the electrodes: Page 26 of 26 © 2015-2023 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers