C9 Metals - summary notes.pdf

C9 METALS Most of the elements in the Periodic Table are metals. (all the elements to the left of the zigzag - - - line in the Periodic Table shown below). General Physical Properties of Metals Property Metals Non-...

C9 METALS Most of the elements in the Periodic Table are metals. (all the elements to the left of the zigzag - - - line in the Periodic Table shown below). General Physical Properties of Metals Property Metals Non-metals thermal Good thermal conductors, as metals Poor thermal conductors, as there are conductivity have delocalised electrons that can carry no free electrons. heat throughout the structure. (Graphite is an exception) electrical Good electrical conductors, due to the Poor electrical conductors, as there conductivity presence of delocalised electrons to are no free electrons. carry the charge. (Graphite is an exception) malleability Malleable (can be hammered and made Brittle when solid and easily break up. and ductility into different shapes) and ductile (can (Graphite is an exception) be drawn into wires), because the layers of metal ions are able to slide over each other without breaking the structure. melting and Have high melting and boiling points Low melting and boiling points boiling points because the strong electrostatic because the weak intermolecular forces attraction between the positive metal do not require a lot of energy to be ions and delocalised electrons (metallic overcome. bond) requires lots of energy to break. (Diamond is an exception). Alloys An alloy is a mixture of a metal with other elements (mostly metals, some contain non-metals as well) Alloys often have properties that are different to the metals they contain. They can be stronger, harder and resistant to corrosion / extreme temperatures. These enhanced properties can make alloys more useful than pure metals. Structure of Alloys Pure metals are very soft because the layers can slide over each other easily. Alloys have different sized atoms, which make the lattice structure irregular. Therefore, the layers cannot slide over each other easily. This makes the alloy stronger (and more brittle) than pure metals. Common Alloys - Composition, Properties and Uses General Chemical Properties of Metals Reaction with Water Reactive metals react vigorously with cold water, producing a metal hydroxide and hydrogen gas. Na + H2O → NaOH + H2 Less reactive metals react with steam to produce a metal oxide and hydrogen gas. Mg + H2O → MgO + H2 Reaction with Dilute Acids Many metals react with dilute acids to produce a salt and hydrogen gas. Zn + HCl → ZnCl2 + H2 Reactivity Series The order of reactivity of metals can be deduced by making experimental observations of reactions between metals and water, acids and oxygen For a more reactive metal: ○ the reaction is more vigorous ○ the temperature change is greater Metal Reaction with cold water Reaction with acid Reaction with oxygen Potassium Reacts violently Reacts quickly Sodium Reacts fast with cold water Reacts quickly Calcium Reacts vigorously Magnesium Slow reaction (reacts with Reacts vigorously Reacts readily Aluminium steam) Reacts readily Carbon Zinc Very slow reaction (reacts Reacts less strongly Reacts Iron slowly with steam) Reacts less strongly Hydrogen Copper No reaction with steam or Reacts slowly Silver No reaction water Gold No reaction *The non-metals hydrogen and carbon are also included in the reactivity series as they are used to extract metals from their oxides. Displacement Reactions A more reactive metal will displace a less reactive metal 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. E.g. When magnesium metal is added to a solution of copper sulfate, magnesium displaces the copper ions from the salt, and forms magnesium sulfate and copper metal. magnesium + copper sulfate → magnesium sulfate + copper Mg (s) + CuSO4 (aq) → MgSO4 (aq) + Cu (s) Displacement reactions can be used to determine the order of reactivity of different metals. Rusting of Iron Rusting 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 Rust Prevention Methods Rust can be prevented by coating iron with barriers that prevent the iron from coming into contact with water and oxygen. Barrier Methods Common barriers include: ○ Paint ○ Oil/ Grease ○ Plastic If the coatings are removed, the iron is once again exposed to water and oxygen and will rust. Galvanising The iron to be protected is coated with a layer of zinc 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. Sacrificial Protection A more reactive metal is attached to a less reactive metal The more reactive metal will oxidise and corrode first, protecting the less reactive metal from corrosion. E.g. Iron is less reactive therefore will not lose its electrons as easily so it is not oxidised; zinc is sacrificed to protect the steel For continued protection, the zinc bars must be replaced before they completely corrode. Uses of Metals Uses of metals depend on their properties. Aluminium Aluminium is above hydrogen in the reactivity series, so it is a reactive metal. However, aluminium quickly reacts with oxygen in the air to form a protective layer of aluminium oxide which sticks to the metal surface. Hence, aluminium appears to be unreactive/ resistant to corrosion. Use of Aluminium Property aeroplane bodies high strength, low density overhead power cables good electrical conductor, low density saucepans good thermal conductor food cans non-toxic, resistant to corrosion Copper Copper is below hydrogen in the reactivity series, so it is an unreactive metal. However, it is widely used due to its different useful properties. Use of Copper Property electrical wiring very good conductor of electricity and ductile pots and pans very good conductor of heat, unreactive, malleable non-toxic, unreactive (does not react with water) and water pipes malleable surface in hospitals antibacterial properties Extraction of Metals Metals are found in the Earth’s crust in different forms. Unreactive metals are usually found as the uncombined element (native element), and can be mined directly from the Earth’s crust. ○ E.g. gold, silver, platinum Most of the useful metals are chemically combined with other substances, such as oxygen. ○ E.g. iron can be found as iron oxide, aluminium as aluminium oxide. A rock that contains enough of the metal to make it feasible to extract is called a metal ore. ○ Iron ore is called haematite, aluminium ore is called bauxite. Metal oxides can be extracted by removing the oxygen (called reduction). The method of extraction depends on the position of the metal in the reactivity series. ○ Oxides of metals below carbon can be reduced by heating the ore with carbon. ○ Oxides of metals above carbon need to be extracted using a method called electrolysis. Potassium Extracted by electrolysis of their molten Sodium salts. Calcium A large amount of electricity is required Magnesium hence the process is expensive. Aluminium Carbon Zinc Extracted by heating with a reducing agent Iron such as carbon, or carbon monoxide in a Lead blast furnace Copper The process is relatively cheap. Silver Occur native as pure metals. Gold Platinum Iron Extraction The main ore of iron is haematite, Fe2O3. The iron is obtained by reduction of Fe2O3 in a blast furnace. Raw Materials: Iron Ore (Haematite), Coke, Limestone and Air Iron ore, coke and limestone are mixed together and fed into the top of the blast furnace. Hot air is blasted into the bottom of the blast furnace. Coke is used as the starting material. It is an impure carbon and it burns in the hot air blast to form carbon dioxide. This is a strong exothermic reaction: C (s) + O2 (g) → CO2 (g) At the high temperatures in the furnace, carbon dioxide reacts with coke to form carbon monoxide: CO2 (g) + C (s) → 2 CO (g) Carbon monoxide (the reducing agent) 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: Fe2O3 (s) + 3CO (g) → 2Fe (l) + 3CO2 (g) Limestone is added to the furnace to remove acidic impurities in the ore. The calcium carbonate in the limestone decomposes to form calcium oxide: CaCO3 (s) → CaO (s) + CO2 (g) The calcium oxide reacts with the silicon dioxide, which is an impurity in the iron ore, to form calcium silicate. This melt and collects as a molten slag floating on top of the molten Iron, which is tapped off separately: CaO (s) + SiO2 (s) → CaSiO3 (l) Aluminium Extraction Aluminium is a reactive metal which sits above carbon on the reactivity series. It cannot be extracted from its ore (bauxite) by carbon reduction, so electrolysis is used. Raw materials: Aluminium Ore (Bauxite) Method of extraction: Electrolysis of molten aluminium oxide The Bauxite is first purified to produce Aluminium Oxide Al2O3. Aluminium oxide has a very high melting point, so it is first dissolved in molten cryolite Cryolite is used to decrease the melting point of aluminium oxide and to dissolve it, thereby reducing the energy use and expenses.. The electrolyte is a solution of aluminium oxide in molten cryolite at a temperature of about 1000 °C. A lot of electricity is required, soit is an expensive process. Reaction at the negative electrode: Al3+ + 3e- → Al The Aluminium melts and collects at the bottom of the cell and is then tapped off. Reaction at the positive electrode: 2O2- - 4e- → O2 Some of the oxygen produced at the positive electrode then reacts with the graphite (carbon) electrode at high temperature to produce carbon dioxide gas: C(s) + O2 (g) → CO2 (g) This causes the carbon anodes to burn away, so they must be replaced regularly.