O level Chemistry Notes (Page 6)
O level Chemistry Notes covering types of oxides, oxidising agents, reducing agents and extraction of metals.
An oxide is a chemical compound of a element or radical combined with oxygen. Some examples of oxides are ZnO (zinc oxide), CO2 (carbon dioxide), H2O (water).
Types of Oxides
Oxides can be classified based on their chemical properties.
- SO3(s) + H2O(l) → H2SO4(aq)
Acidic oxides are also known as acid anhydrides.
Other examples are:
- SO2 – Sulphur dioxide (which dissolves in water to form sulphurous acid)
- CO2 – Carbon dioxide (which dissolves in water to form carbonic acid)
- P2O5 – Phosphorus pentoxide (which dissolves in water to form phosphoric acid).
Acidic oxides react with alkalis to form salts and water only. For example, sulphur trioxide reacts with sodium hydroxide to form sodium sulphate and water as follows:
- SO3(s) + 2NaOH(aq) → Na2SO4(aq) + H2O(l)
Basic oxides are oxides of metals which neutralise acids. For example, copper oxide neutralises sulphuric acid to form copper(ii) sulphate and water as follows:
CuO(s) + H2SO4(aq) → CuSO4(aq) + H2O(l)
Other examples of basic oxides are:
- CaO – calcium oxide
- MgO – magnesium oxide
Some basic oxides such as iron(III) oxide (Fe2O3) and copper(II) oxide CuO do not react with water and appear neutral to litmus. However oxides of metals high in the reactivity series such as lithium, pottasium, sodium, calcium, magnesium dissolve in water to form alkalis. For example: Sodium oxide dissolves in water to form sodium hydroxide (an alkali) as follows:
- Na2O(s) + H2O(l) → 2NaOH(aq)
Neutral oxides are oxides which are neither acidic nor basic. Examples are:
- H2O – water
- CO – carbon monoxide
- N2O – dinitrogen oxide
- NO – nitrogen oxide
Amphoteric oxides are oxides which have both acidic and basic propertities. Examples are aluminium oxide, zinc oxide and lead oxide. They form salts with acids and complex salts with alkalis. For example: aluminium oxide reacts with hydrochloric acid to form aluminium chloride and water, as follows:
- Al2O3(s) + 6HCl(aq) → 3H2O(l) + 2AlCl3(aq).
Aluminium oxide also reacts with sodium hydroxide to form sodium aluminate as follows:
- Al2O3(s) + 2NaOH(aq) + 3H2O(l) → 2NaAl(OH)4(aq)
Oxidising and reducing agents
A reducing agent loses electrons and is oxidised while reducing another substance.
An oxidising agent gains electrons and is reduced while oxidising another substance.
Zn + CuSO4 → ZnSO4 + Cu
Can be split into ionic equations:
Zn → Zn2+ + 2e–
and Cu2+ + 2e– → Cu
Zn has been oxidised to Zn2+ because it lost 2 electrons and those 2 electrons where gained by Cu2+ which was reduced to Cu. Zn is however the reducing agent and Cu2+ is the oxidising agent.
Mg + 2HCl → MgCl2 + H2
Can be split into ionic equations:
Mg → Mg2+ + 2e–
and 2H+ + 2e– → H2
Mg is oxidised to Mg2+ and Mg is the reducing agent. 2H+ reduced to H2 and 2H+ is the reducing agent.
Extraction of metals
Only the least reactive metals (e.g. gold) are found uncombined in their native form. The majority are found as compounds such as oxides, carbonates, chlorides and sulphides naturally. However oxide ores are the most common.
An ore is a naturally available impure compound from which a metal can be extracted economically. In some cases sulphide ores and carbonate ores need to be roasted in air to be converted into the metal oxides from which the metals are eventually extracted.
For metals in the middle of the reactivity series, the most common and economic method of extraction is by heating the oxide with carbon or carbon monoxide. This is partly because carbon is plentiful and cheap, in the form of coke or charcoal, and it forms a gaseous product (carbon dioxide) which is easy to remove.
Metals above zinc are too reactive for their oxides to be reduced by carbon, and and are usually extracted by electrolysis, which is a more expensive method.
Extraction using metal displacement reactions
The Thermite Process for the extraction of Iron
Iron can be extracted from iron(II) oxide using a potentially explosive mixture called Thermite. The reaction is only performed on a small scale and is still used today to weld railway tracks together.
In this reaction aluminium metal is used to displace iron from iron(III) oxide because aluminium is more reactive than iron. In general, a more reactive metal is a good reducing agent for a lesser reactive metal. The process is so exothemic that the heat generated from the reaction is enough to melt the iron.
Fe2O3 + 2Al → 2Fe + Al2O3
The molten iron can then be placed between the railway tracks and when it cools it solidifies joining the railway tracks.
Extraction of iron using Carbon Reduction
In the Blast Furnace, Iron(III) oxide (haematite) is mixed with coke (almost pure carbon) and limestone (calcium carbonate) to form a mixture called a charge. The charge enter the blast furnace from the top, and a blast of hot air is blown in at the bottom.
Near the bottom of the furnace, some of the carbon is oxidised to overally produce carbon monoxide, giving out much heat.
2C + O2 → 2CO
Iron ore is then reduced by both CO and C:
Fe2O3 + 3CO → 2Fe + 3CO2
Fe2O3 + 3C → 2Fe + 3CO
Limestone (CaCO3) is decomposed by the heat to form calcium oxide which then remove the sandy impurities (SiO2 is Iron to form slag, calcium silicate (CaSiO3)
CaCO3 → CaO + CO2
CaO + SiO2 → CaSiO3 (molten slag)
The molten iron is runoff at the bottom of the furnace, where it is tapped off from time to time and the molten slag, which floats on top of the iron is also removed periodically.
Rusting of Iron
Most metals tend to oxidise slowly if left out in the air, especially if the air is wet and if it contains salt. This type of corrosion is most common for iron and it is known as rusting. Rust is iron(III) oxide, Fe2O3.
Rusting can only occur if iron comes into contact with both air and water. We usually try to prevent rusting by keeping out both air and water using the following methods:
- Painting: it is cheap, and covers the surface. But the paint may chip, and once surface is exposed rusting begins.
- Coating with zinc: (galvanising). This involves dipping the iron into molten zinc. It forms a tougher coating than paint, and once scratched rusting is slow, because the zinc is more reactive than iron and oxidises in preference (a form of sacrificial protection). Used in corrugated iron, galvanised buckets etc.
- Sacrificial protection: fastening blocks of a more reactive metal such as magnesium to the steel, at intervals (e.g. to a steel mast on a yacht). The more reactive metal oxidises in preference to the iron.
Extracting Metals using Electrolysis
Electrolysis is the passage of electricity through a liquid called an electrode accompanied by a chemical change taking place at the electrodes.
The electrodes are the conductors (made of inert metal or graphite) by which the current enters or leaves the electrolyte.
The positive electrode is known as the anode and the negative electrode is known as the cathode.
For a liquid electrolyte to be able to conduct electricity, it must contain ions. The ions must be free to move, which is why electrolysis takes place in solutions or molten electrolytes, not solids.
Ions that are attracted to the cathode are called cations and they are positively charged and ions that are attracted to the anode are called anions and they are negatively charged.
When a positively charged ion reaches the cathode, it accepts enough electrons to make it a neutral atom:
e.g. Na+ + e– → Na(l);
or Al3+ + 3e– → Al(l)
At the anode, negative ions give up their extra electrons to form stable nuetral molecules:
e.g. Br– → Br + e–
then Br + Br → Br2
the overall reaction is:
2Br– → Br2 + 2e–
Metals and hydrogen form positive ions, and are set free at the cathode. Non-metals ( except hydrogen ) form negative ions and are set free at the anode.
Extraction of Aluminium
Most of the very reactive metals are extracted by electrolysing their molten chlorides. However, aluminium is extracted from bauxite, which is mainly Al2O3.
Aluminium oxide is insoluble in water, and its melting point is very high. It is dissolved in a molten cryolite (formula Na3AlF6) at 900 degrees Celsius.
At the cathode, aluminium ions each accept three electrons to change them to neutral aluminium atoms:
Al3+ + 3e– → Al(l)
At the anode, oxide ions each give up two electrons to form oxygen atoms. These combine in pairs to form oxygen molecules.
2O2– → O2 + 4e–
In practice, a carbon anode is used, and the oxygen reacts with the carbon to form carbon dioxide.
C + O2 → CO2
And the overall reaction is
2Al2O3 → 4Al + 3O2