O level Chemistry Notes (Page 5)
Oxygen and Oxides
Oxygen makes up to about 21% of the Earths Atmosphere, Nitrogen about 78%, Argon about 0.9%, Carbon dioxide about 0.04% and also variable amounts of water vapour.
Finding the % oxygen in air
This can be shown by an experiment using two gas syringes:
The one on the left starts with (say) 80 cm3 of air. The copper wire is heated, and the air is shunted back and forth through the combustion tube where the oxygen reacts with the iron.
When the apparatus has cooled back to room temperature the final volume of the gas is (say) 63.2 cm3, which means the proportion of oxygen is = 21%.
Industrial Preparation of Oxygen
Oxygen is prepared industrially by the fractional distillation of liquid air. Air is first cooled to -78 degrees Celsius where water and CO2 are removed as solids to avoid blocking the tubes. Air is further cooled to –200 degrees Celsius by a process of rapid contraction and expansion and then fractionally distilled. Nitrogen boils off first at –196 degrees Celsius and Oxygen which has a boiling point of -183 degrees Celsius is collected as a liquid from the bottom of the fractionating column.
Laboratory Preparation of Oxygen
In the laboratory oxygen is preparation from hydrogen peroxide,
2H2O2 → O2 + 2H2O
MnO2 is added as a catalyst and the oxygen is collected by downward displacement of water.
Uses of Oxygen
Oxygen is used in:
- steel manufacture (removal of carbon from molten iron in the oxygen-lance process)
- welding and cutting metals (e.g. oxy-acetylene torch)
- breathing apparatus (e.g. subaqua and oxygen masks in hospitals).
Oxidation and Reduction in terms of oxygen
The term oxidation is refers to the gain of oxygen by a substance.
Reduction is defined as the loss of oxygen by a substance.
Oxidation and reduction always occur together in a chemical reaction. If one substance is oxidised, another must be reduced. Reactions involving oxidation and reduction are called redox reactions.
Reactions of Oxygen
Reactions of with metals.
- Metals react with oxygen to form alkaline metal oxides.
2Ca + O2 → 2CaO(s)
Reactions of with non-metals.
- Non-metals react with oxygen to form acidic non-metal oxides.
C + O2 → CO2 (g)
Reactions of with hydrocarbons.
- Hydrocarbons undergo complete and incomplete combustion when reacted with oxygen.
C3H8 + 5O2 → 3CO2 + 4H2O
Carbon dioxide is a colourless gas which is more dense than air. It is responsible for acid rain, the greenhouse effect and photosynthesis in plants.
Laboratory Preparation of Carbon Dioxide
Dilute hydrochloric acid is reacted with calcium carbonate.
2HCl (aq) + CaCO3(s) → CaCl2(aq) + H2O(l) + CO2(g)
If necessary the carbon dioxide gas can be dried by passing it through concentrated sulphuric acid.
Reactions of Carbon dioxide
Carbon dioxide reacts with water to form carbonic acid (the main component of acid rain).
CO2(g) + H2O(l) → H2CO3(aq)
Being a non-metal oxide carbon dioxide is acidic and so it reacts with alkalis to make salts (called carbonates).
CO2(g) + Ca(OH)2(aq) → CaCO3(s) + H2O(l)
Ca(OH)2 is commonly known as limewater. A cloudy precipitate of calcium carbonate is produced when it reacts with carbon dioxide. This reaction is the test for carbon dioxide
Uses of carbon dioxide
Carbon dioxide is used:
- in carbonated drinks.
- Fire extinguishers (as it is non flammable and does not support combustion).
- It is also used in its solid form dry ice as a refrigerant.
Sulphur dioxide is a pollutant formed from the combustion of sulphur impurities in Coal and petrol reaction with water.
SO2 + H2O → H2SO3
[Sulphurous acid – also found in acid rain]
Nitrogen dioxide, NO2 is formed in the internal combustion engine from the combustion of nitrogen in air. It causes acid rain and photochemical smog.
Hydrogen has atoms with only one electron each.
It is a non-metal, which forms H2 molecules, and is a gas at room temperature. It has an unusual position in the periodic table as it is sometimes placed above Lithium in group 1. This is because although it is a non-metal, it has one electron in its outer shell and so, like the alkali metals, it can form ions with a charge of +1 (valency 1).
Laboratory preparation of hydrogen
Hydrogen is made by reacting zinc with dilute sulphuric acid and collected by upward delivery. If dry hydrogen is needed, the gas is passed through a drying agent, silica gel or concentrated sulphuric acid.
Combustion of Hydrogen
Combustion of hydrogen with air or oxygen produces water as the only product.
2H2(g) + O2(g) → 2H2O(l)
The reaction also gives out useful energy and hydrogen is often refered to as ‘the fuel of the future’ as water is its only combustion product.
Test for water
There are a number of tests for the presence of water.
- Cobalt chloride paper Turns from Blue to Pink.
- Anhydrous copper(II) sulphate Turns from White to Blue.
The test for pure water is its boiling point, 100 degrees celsius and its melting point 0 degrees celsius.
The Reactivity Series
Below is part of the reactivity series starting with the most reactive metal.
* Al forms an inert coating of aluminium oxide, which makes it seem less reactive than it really is.
Evidence about the reactivity of metals can be gained from how vigorously they react with water or acids.
The most reactive metals react violently with cold water.
- potassium melts, fizzes around the surface, catches fire and burns with a lilac flame
- sodium melts to a silver ball, fizzes around the surface, giving hydrogen
- calcium reacts by fizzing gently.
2K + 2H2O → 2KOH + H2
2Na + 2H2O → 2NaOH + H2
Ca + 2H2O → Ca(OH)2 + H2
The other metals below Ca in the series react with steam, to form the metal oxide and hydrogen:
Zn + H2O → ZnO + H2
Some metals will react to form metal salts and hydrogen.
Mg + 2HCl → MgCl2 + H2
It is often more reliable to compare the reactivity of metals directly. Displacement reactions are often used to compare the reactivity of metals.
For example, magnesium, when heated with copper(II) oxide, gives out much heat and forms copper metal:
Mg(s) + CuO(s) → MgO(s) + Cu(s)
This shows that magnesium is more reactive than copper. If we tried heating copper and magnesium oxide, the reaction would not work.
Even when a metal is dipped into a solution of the salt of a second metal. If the first metal is able to displace the second, it is more reactive.
e.g. If an iron wire is dipped into copper sulphate solution it becomes coated in reddish copper, while the blue colour fades, which shows that iron is more reactive than copper.
Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)
The process of adding oxygen to copper to form copper oxide is called oxidation.
2Cu + O2 → 2CuO
When copper reacts with oxygen, each copper atom loses 2 electrons to each oxygen atom to form an ionic compound of copper(ii) oxide as follows.
2Cu + O2 → 2Cu2+ + 2O2-
Therefore the process involves removing two electrons from each copper atom to form copper ions as follows:
Cu → Cu2+ + 2e–
This means that oxidation is the gain of oxygen or removal of electrons.
It can also involve removal of hydrogen from a molecule. In electrolysis, the process which occurs at the anode always involves removal of electrons, and so it is oxidation.
Reduction is the reverse of oxidation and is defined as the loss of oxygen, gain of electrons or gain of hydrogen. In electrolysis the cathode process involves positive ions gaining electrons, and so it is reduction.
Oxidation and reduction always occur together in a chemical reaction: if one substance is oxidised, another must be reduced. Reactions involving oxidation and reduction are called redox reactions.
In the examples below the underlined atom is being oxidised, while the one with a double underline is being reduced:
2Cu + O2 → 2CuO
copper is oxidised and oxygen is reduced
Zn + CuSO4 → ZnSO4 + Cu
This is a displacement reaction which when written in ionic form becomes:
Zn + Cu2+ → Zn2+ + Cu
Zinc is oxidised because it loses electrons and Copper is reduced because it gains electrons.
Mg + 2HCl → MgCl2 + H2
Becomes Mg + 2H+ → Mg2+ + H2
Magnesium is oxidised and hydrochloric acid is reduced.