metals and metal extraction
9.1 Physical Properties of Metals and Non-metals
(a) Thermal conductivity:
Metals are generally good conductors of heat. This means they transfer thermal energy efficiently.
Non-metals are generally poor conductors of heat (insulators).
(b) Electrical conductivity:
Metals are good conductors of electricity. This is due to the presence of delocalized electrons in their structure, which can carry an electrical charge.
Non-metals are generally poor conductors of electricity (insulators). Graphite is an exception, as it can conduct electricity.
(c) Malleability and ductility:
Malleability: Metals can be hammered or pressed into thin sheets without breaking (e.g., gold leaf).
Ductility: Metals can be drawn out into thin wires (e.g., copper wires).
Non-metals are generally brittle and will shatter when hammered or stretched.
(d) Melting points and boiling points:
Metals generally have high melting points and boiling points, indicating strong forces of attraction between their atoms.
Non-metals have a wider range of melting and boiling points. Some have very low melting and boiling points (e.g., gases), while others have high ones (e.g., diamond).
9.1.2 Reactions of Metals
Reactions of metals with dilute acids:
Metals react with dilute acids (e.g., hydrochloric acid, sulfuric acid) to produce a salt and hydrogen gas.
General equation: Metal + Acid → Salt + Hydrogen
Example: Magnesium + Hydrochloric acid → Magnesium chloride + Hydrogen
Reactions of metals with cold water and steam:
Some reactive metals (e.g., potassium, sodium, calcium) react vigorously with cold water to produce a metal hydroxide and hydrogen gas.
Less reactive metals (e.g., magnesium) react slowly with cold water but react more quickly with steam to produce a metal oxide and hydrogen gas.
Unreactive metals (e.g., copper, gold) do not react with water or steam.
Reactions of metals with oxygen:
Metals react with oxygen to form metal oxides.
The reactivity of metals with oxygen varies. Some metals (e.g., sodium) react rapidly at room temperature, while others (e.g., iron) require heating.
General equation: Metal + Oxygen → Metal oxide
Example: 2Mg + O2 → 2MgO
9.2 Uses of Metals
(a) Aluminium in aircraft manufacture:
Low density: Aluminium is very light, which reduces the weight of the aircraft and improves fuel efficiency.
(b) Aluminium in overhead electrical cables:
Low density: Makes the cables lighter, reducing the need for strong supports.
Good electrical conductivity: Allows the efficient transmission of electricity.
(c) Aluminium in food containers:
Resistance to corrosion: Aluminium does not react easily with air or water, protecting the food from contamination.
(d) Copper in electrical wiring:
Good electrical conductivity: Copper allows electricity to flow easily.
Ductility: Copper can be drawn into thin wires.
9.4 Reactivity Series
9.4.1 Order of the reactivity series:
The reactivity series lists metals in order of their decreasing reactivity:
Potassium, sodium, calcium, magnesium, aluminium, carbon, zinc, iron, hydrogen, copper, silver, gold
9.4.2 Reactions of metals with water and acids:
(a) Potassium, sodium, and calcium with cold water:
These metals react very vigorously with cold water, producing a metal hydroxide and hydrogen gas. The reactions are exothermic (release heat) and can be dangerous.
(b) Magnesium with steam:
Magnesium reacts with steam to produce magnesium oxide and hydrogen gas. This reaction occurs more readily with steam than with cold water.
(c) Magnesium, zinc, iron, copper, silver, and gold with dilute hydrochloric acid:
Magnesium, zinc, and iron react with dilute hydrochloric acid to produce a salt and hydrogen gas. The rate of reaction decreases as you go down the reactivity series.
Copper, silver, and gold do not react with dilute hydrochloric acid.
These reactions demonstrate the differing reactivities of metals. Metals higher in the reactivity series react more readily with acids.
9.4.3 Deducing an order of reactivity:
The order of reactivity can be deduced from experiments involving reactions with water, acids, or displacement reactions.
For example, if metal A reacts with the salt solution of metal B, then metal A is more reactive than metal B.
9.4.4 Relative reactivities and positive ion formation:
The reactivity of a metal is related to its tendency to form positive ions (cations).
More reactive metals lose electrons more easily to form positive ions.
Displacement reactions: A more reactive metal will displace a less reactive metal from its salt solution.
For example, zinc will displace copper from copper sulfate solution: Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
9.4.5 Unreactivity of aluminium:
Aluminium is a reactive metal, but it appears unreactive due to the formation of a thin, strong layer of aluminium oxide (Al2O3) on its surface.
This oxide layer is impermeable and protects the aluminium from further reaction with air, water, or dilute acids.
2.7 Metallic Bonding
2.7.1 Description of metallic bonding:
Metallic bonding is the electrostatic attraction between positively charged metal ions and a 'sea' of delocalized electrons.
Metal atoms are arranged in a giant metallic lattice structure.
The outermost electrons of the metal atoms are delocalized, meaning they are not bound to any specific atom and can move freely throughout the lattice.
2.7.2 Properties of metals explained by structure and bonding:
(a) Good electrical conductivity:
The delocalized electrons are free to move and carry an electrical charge throughout the metal lattice.
(b) Malleability and ductility:
The layers of positive ions in the metallic lattice can slide over each other without breaking the metallic bond.
The delocalized electrons maintain the bonding between the ions, even when they are moved.
9.3 Alloys
9.3.1 Description of an alloy:
An alloy is a mixture of a metal with other elements. These elements can be other metals or non-metals.
Examples:
Brass: A mixture of copper and zinc.
Stainless steel: A mixture of iron with other elements, such as chromium, nickel, and carbon.
9.3.5 Structure and properties of alloys:
Alloys are harder and stronger than pure metals because the different-sized atoms in the alloy disrupt the regular arrangement of the metal lattice.
This disruption prevents the layers of atoms from sliding over each other easily, making the alloy less malleable and less ductile but stronger.
9.3.2 Uses of alloys:
Alloys are often more useful than pure metals due to their enhanced properties, such as increased hardness, strength, and resistance to corrosion.
9.3 Alloys (Continued)
9.3.3 Uses of alloys:
Alloys are used in various applications due to their specific properties:
Stainless steel: Used in cutlery because of its hardness and resistance to rusting (corrosion).
9.3.4 Representations of alloys from diagrams of structure:
Alloys can be represented in diagrams as a mixture of different sized atoms disrupting the regular lattice structure of a pure metal.
6.4 Oxidation and Reduction (Redox)
6.4.1 Oxidation number (Roman numerals):
Roman numerals are used to indicate the oxidation number of an element in a compound.
Example: Iron(III) oxide (Fe2O3) indicates that iron has an oxidation number of +3.
6.4.2 & 6.4.4 Redox reactions:
Redox reactions involve both oxidation and reduction occurring simultaneously.
Oxidation and reduction can be defined in terms of the loss and gain of oxygen.
6.4.3 Oxidation and reduction (electrons):
Oxidation is the loss of electrons.
Reduction is the gain of electrons.
In a redox reaction, one substance loses electrons (is oxidised), and another substance gains electrons (is reduced).
6.4.5 Identifying oxidation and reduction:
You should be able to identify which substances are oxidised and reduced in a redox reaction by looking at the changes in electron transfer or oxygen gain/loss.
6.6.6 Defining oxidation:
Oxidation can be defined as:
(a) Loss of electrons
(b) An increase in oxidation number
6.6.7 Defining reduction:
Reduction can be defined as:
(a) Gain of electrons
(b) A decrease in oxidation number
6.4.8 Redox reactions (electron transfer):
Redox reactions are identified as reactions involving the transfer (gain and loss) of electrons.
6.4.9 Identifying redox reactions by oxidation numbers:
Rules for assigning oxidation numbers:
(a) The oxidation number of an element in its uncombined state is zero (e.g., O2, Mg).
(b) The oxidation number of a monatomic ion is the same as the charge on the ion (e.g., Cl- is -1).
(c) The sum of the oxidation numbers in a compound is zero.
(d) The sum of the oxidation numbers in an ion is equal to the charge on the ion.
You can use these rules to track changes in oxidation numbers and identify redox reactions.
6.4.10 Identifying redox reactions by colour changes:
Redox reactions can sometimes be identified by characteristic colour changes:
Acidified aqueous potassium manganate(VII): Changes from purple to colourless when reduced.
Aqueous potassium iodide: Changes from colourless to brown when oxidised.
6.4.11 Oxidising agent:
An oxidising agent is a substance that oxidises another substance and is itself reduced.
6.4.12 Reducing agent:
A reducing agent is a substance that reduces another substance and is itself oxidised.
6.4.13 Identifying oxidising and reducing agents:
In a redox reaction, you should be able to identify the substance that is oxidised (reducing agent) and the substance that is reduced (oxidising agent).
9.5 Rusting of Iron and Steel
9.5.1 Conditions for rusting:
Rusting is the corrosion of iron and steel to form hydrated iron(III) oxide (rust).
The conditions required for rusting are:
Oxygen
Water
9.5.2 Barrier methods to prevent rusting:
Common barrier methods include:
Painting
Greasing
Coating with plastic
9.5.3 How barrier methods work:
Barrier methods prevent rusting by excluding either oxygen or water, or both, from coming into contact with the iron or steel surface.
9.5.4 Galvanising:
Galvanising is a method of rust prevention that involves coating iron or steel with a layer of zinc.
Zinc acts as both a barrier and provides sacrificial protection.
9.5.5 Sacrificial protection:
Sacrificial protection: A more reactive metal (like zinc) is used to protect a less reactive metal (like iron) from corrosion.
The more reactive metal corrodes instead of the iron.
This occurs because the more reactive metal has a greater tendency to lose electrons (get oxidised) than iron.
9.6 Metal Extraction
9.6.1 Ease of obtaining metals from their ores:
The ease with which metals can be extracted from their ores is related to their position in the reactivity series.
Unreactive metals (e.g., gold, silver) are found in the Earth's crust in their uncombined form and are easy to extract.
Reactive metals (e.g., sodium, aluminium) are found as compounds and are difficult to extract, requiring more energy.
9.6.2 Extraction of iron from hematite in the blast furnace:
Hematite (Fe2O3) is a common ore of iron.
The extraction of iron in a blast furnace involves the following steps:
(a) Burning of carbon (coke): Coke burns in the presence of air to produce heat and carbon dioxide: C + O2 → CO2
(b) Reduction of carbon dioxide: Carbon dioxide reacts with more coke to form carbon monoxide: C + CO2 → 2CO
(c) Reduction of iron(III) oxide: Carbon monoxide reduces iron(III) oxide to iron: Fe2O3 + 3CO → 2Fe + 3CO2
(d) Thermal decomposition of calcium carbonate (limestone): Limestone decomposes to produce calcium oxide and carbon dioxide: CaCO3 → CaO + CO2
(e) Formation of slag: Calcium oxide reacts with silica (sand) to form slag (calcium silicate): CaO + SiO2 → CaSiO3
Note: You are not required to know the symbol equations for the extraction of iron, but they are helpful to understand the process.
9.6.3 Extraction of aluminium:
The main ore of aluminium is bauxite.
Aluminium is extracted from bauxite by electrolysis.