Reduction and Oxidation
Reduction and oxidation reactions are often known as Redox reactions. These reactions can involve the transfer of oxygen between reactants in a reaction as they form the products. They can also involve the transfer of electrons between the reactants to form the products.
Higher level questions and answer will focus on the transfer of electrons during redox reactions rather than oxygen. However, you will be expected to explain them in terms of both oxygen and electron transfers.
Oxidation: The gaining of oxygen by a chemical during a reaction.
Or: The loss of electrons.
Reduction: The loss of oxygen by a chemical during a reaction.
Or: The gain of electrons.
This can be remembered using OIL RIG Oxidation is loss Reduction is gain
2Fe2O3 + 3C → 4Fe + 3CO2 (part of the process of extracting iron from iron ore).
Fe2O3 is converted to Fe - the oxygen has been lost from this molecule and so it has been reduced to Fe (iron).
C is converted to CO2 - the carbon has gained oxygen in the reaction and so has been oxidized.
The Fe in the Fe2O3 is an ion of Fe with 3 less electrons Fe3+, during the reaction these electrons are replaced (gaining electrons), so it converts back to a neutral Fe atom.
Fe3+ + 2e- → Fe the iron has gained electrons so it is reduced, (oil__RIG__).
The C begins as a neutral carbon atom, in the reaction it merges its outer electrons with the oxygen atoms and in effect loses them.
C - 4e- → C4- the carbon has lost electrons so it is oxidized, (OIL rig).
(Note:it is not forming an ion as it forms a covalent bond with the two oxygen atoms).
Not all redox reactions involve the transfer of oxygen in the reaction. Displacement reactions where a more reactive metal displaces a less reactive metal are also examples of redox reactions as they involve the transfer (gain and loss) of electrons between the metals.
Magnesium + Copper sulfate → Magnesium sulfate + Copper.
Mg + CuSO4 → MgSO4 + Cu
Mg - 2e- → Mg2+
The magnesium atom has been oxidised to a magnesium ion by the loss of electrons, (__OIL__rig).
Cu2+ + 2e- → Cu
The copper ion has been reduced to a copper atom by gaining electrons, (oil__RIG__).
Oxidising and Reducing Agents
An oxidising agent is a chemical that adds oxygen or removes electrons from another chemical.
In the above example the copper is the oxidising agent as it has removed the electrons from the magnesium.
A reducing agent is a chemical that removes oxygen or adds electrons to another chemical.
In the above example the magnesium is the reducing agent as it adds electrons to the cooper to reduce it.
Acids and Alkalis
Acids and alkalis are often thought of as being “opposites”, normally due to the fact that they appear at opposite ends of the pH scale and when combined produce a neutral solution. However, they are two separate groups of chemicals that produce different ions when dissolved in water.
Acids are chemicals that dissolve into water and to release hydrogen ions (H+), the more hydrogen ions that are released the stronger the acid.
Hydrogen chloride dissolves into water to produce H+ ions and Cl- ions forming hydrochloric acid.
Sulfuric Acid (H2SO4) produces 2H+ ions and an SO42- group ion, (sulfate group).
Nitric Acid (HNO3), produces an H+ ion and an NO3- ion, (nitrate group).
Alkanes are produced when a bases (solids) dissolves to release hydroxide ions (OH-).
The greater the concentration of OH- ions the stronger the alkaline solution.
Sodium hydroxide (NaOH) dissolves to produce an Na+ ion and an OH- ion.
Magnesium hydroxide (Mg(OH)2) will dissociate in water to give Mg2+ ion and two OH- ions.
The pH scale is used to measure if a substance is acid, neutral or alkaline on a scale of 1 to 14.
1 - 6: Acids
pH is a measure of the hydrogen ion concentration in a solution, actually it the inverse of the hydrogen ion concentration
The pH scale is an example of a logarithmic scale, this means that a__ x10 increase in the hydrogen ion concentration results in a decrease of 1 on the pH scale__. Thus an acid of pH 2 is ten times stronger than an acid of pH 3 and 1000 times stronger than an acid of pH 4.
- Hydrochloric acid produced in the stomach - pH 2
- Lemon juice (citric acid) - pH 3
- ¼ M HCl - pH 4 /5
- Distilled water - pH 7
- 1 M NaOH - pH 9/10
- Bleach - pH 12
There are several ways to test a substance to find its pH.
- Litmus is an indicator solution that turns from Blue to Red to indicate an acid or from Red to Blue to indicate an alkaline. This can be a liquid to add to a solution or as a dye dried onto paper strips.
- Universal indicator is also an indicator that changes colour. The colours range from dark reds through green to blue and purples to give an indication of the pH from 1 to 14.
- Electronic sensors or probes can also measure the pH and produce a digital read out with the pH normally to 1 or 2 decimal places.
A neutralisation reaction is one in which a low pH chemical (acids) reacts with a high pH chemical (alkalis or bases) to form a neutral solution.
Hydrochloric acid + Sodium hydroxide → Sodium chloride + water
2HCl + 2NaOH → 2NaCl + 2H2O
The hydrogen ions from the HCl combine with the hydroxide ions of the NaOH to form water and are thus neutralised.
2H+ + 2OH- → 2H2O
2Na++ 2Cl- → 2NaCl
The general formula for this type for neutralisation reaction is:
Acid + Base → Salt + Water
There are other types of neutralisation reactions that have general formula.
Alkalis containing the carbonate group (CO3) react as follows with an acid;
Acid + Carbonate → Carbonate salt + Carbon dioxide + water
Hydrochloric acid + Calcium carbonate → Calcium carbonate + Carbon dioxide + water
2HCl + CaCO3 → CaCl2 + CO2 + H2O
Metals can also react with acids to form a neutral solution.
Acid + Metal → Salt + Hydrogen
Magnesium + Sulfuric acid → Magnesium sulfate + Hydrogen H2SO4 + Mg → MgSO4 + ↑H2(g)
Note: A salt is a ionic compound that contains a metal ion bonded to a salt group:
- Cl group - chlorides (NaCl sodium chloride) from hydrochloric acid
- SO4 group - sulfates (CuSO4 copper sulfate) from sulfuric acid
- NO3 group - nitrates (KNO3 potassium nitrate) from nitric acid
- CO3 group - carbonates ( CaCO3 calcium carbonate) from, carbonic acid
There are many other alt groups, the four above are the most commonly used examples in a GCSE examination.