CHEM1011 Lecture Notes - Lecture 11: Inert Gas, Lead, Hydrogen Sulfide
2 Jul 2018
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TOPIC 11.
OXIDATION AND REDUCTION REACTIONS.
Early concepts of oxidation.
The name "oxidation" was initially applied to reactions where substances combined
with the element oxygen. Thus any substance burning in air was said to be oxidised,
the product being some type of oxide. For example, burning carbon to produce carbon
dioxide is an oxidation, as shown by the equation
C + O2 CO2
Subsequently it was realised that reactions of substances with elements other than
oxygen were of essentially the same type. For example, hydrogen can react with
oxygen to form the compound water, but equally it can react with chlorine to form the
compound hydrogen chloride. In both reactions the free element hydrogen is
converted to a compound of hydrogen and another non-metal, and so both were
classed as oxidations even though no oxygen was involved in the second case.
2H2 + O2 2H2O
H2 + Cl2 2HCl
The reverse reaction, conversion of compounds such as oxides of metals to the
elemental metal were called "reduction" reactions, for example, the reduction of
copper(II) oxide to copper by heating with charcoal (carbon).
2CuO + C 2Cu + CO2
The gain or loss of oxygen is still a useful way of recognising some oxidation or
reduction reactions, but with a knowledge of the structure of atoms, a rather different
definition is now more widely used.
Oxidation reactions as a loss of electrons.
Consider the following reaction in which the metal, magnesium, is treated with
hydrochloric acid, as discussed in Topic 6. The magnesium dissolves to form Mg2+
ions in solution and hydrogen gas is evolved. The equation given previously was
Mg(s) + 2H+(aq) Mg2+(aq) + H2(g)
Notice that the free element magnesium (Mg) has been converted to the compound
magnesium chloride, MgCl2, which is present in solution as its component ions
Mg2+(aq) and ClS(aq). Evaporating the water would isolate this compound as shown in
the following equation.
Mg2+(aq) + 2ClS(aq) MgCl2(s)
This is therefore just like the examples of oxidation given above where elements were
converted to compounds. However, when the reaction equation is written in this way
the electronic nature of the change is more apparent. The electrically neutral Mg
atoms are converted to the charged Mg2+ cations. For this to occur, each Mg atom has
XI - 1
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XI - 2
OXIDATION IS THE LOSS OF ELECTRONS
REDUCTION IS THE GAIN OF ELECTRONS
lost 2 electrons according to the following HALF EQUATION:
Mg(s) Mg2+(aq) + 2eS...... (1) Oxidation
At the same time, the H+ ions have gained 1 electron per ion to form hydrogen atoms
which have become H2 molecules according to the following half equation.
2H+(aq) + 2eS H2(g) ...... (2) Reduction
In this second half reaction the element hydrogen (H2) has been formed.
It can be seen from this example that an oxidation reaction is one in which a species
loses electrons and a reduction reaction is one in which a species gains electrons. As
electrons lost by one species must be accepted by another, then both reactions must
occur simultaneously. Thus equation (1) is an OXIDATION HALF REACTION
(electrons lost by Mg) and equation (2) is a REDUCTION HALF REACTION
(electrons gained by H+). To emphasise that an oxidation is always accompanied by a
reduction, the term REDOX REACTION is used.
The overall reaction equation is obtained by combining the separate half equations in
such a way that the number of electrons lost exactly equals the number of electrons
gained, and the final equation is called a REDOX EQUATION. In this example, the
balanced overall equation results by simply adding the two half equations.
Mg(s) + 2H+(aq) Mg2+(aq) + H2(g)
Note that the 2 electrons which appeared previously on each side of the half equations
have now cancelled out.
All reactions of the type previously regarded as oxidations on the basis of gain of
oxygen atoms have the same characteristic as in the magnesium example above - the
transfer of electrons from the species oxidised to the species reduced. Thus the most
general definitions of oxidation and reduction reactions are:
The species which causes the oxidation to occur (H+ in the above example) is called
the OXIDIZING AGENT or the OXIDANT, while the species which is oxidized is
called the REDUCING AGENT or REDUCTANT. Equations written in the manner
of (1) and (2) above are called ION-ELECTRON HALF EQUATIONS.
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XI - 3
The following are all examples of redox reactions.
Zn(s) + 2H+(aq) Zn2+(aq) + H2(g)
2Na(s) + Br2(l) 2NaBr(s)
2Fe(s) + 3Cl2(g) 2FeCl3(s)
2HgO(s) + heat 2Hg(l) + O2(g)
S(s) + O2(g) SO2(g)
In each example, free elements were converted to compounds and/or elements
combined in compounds were converted back to the free state.
Electrolysis.
At the beginning of the course, the concepts of compounds and of chemical change
were illustrated by the electrolysis of water whereby the passing of electricity through
the compound, water, caused the bonds between its component atoms, H and O, to
break and for new bonds to form leading to the free elements H2 and O2. Although not
discussed at the time, the reactions taking place at the two electrodes are oxidation and
reduction.
To most simply illustrate this, consider at high temperature molten sodium chloride
being subjected to electrolysis. The melt contains only sodium and chloride ions freed
from the constraints of the ionic lattice without the intervention of water molecules
and moving independently throughout the liquid. Note that this is not a solution of
sodium chloride such as dealt with in Topic 6 where the ions interact with water to
become aquated Na+(aq) and Cl−(aq) species which have been able to separate from
each other in the crystal as a result of the energy released in the aquation process.
Molten sodium chloride derives the required energy from the external heat supplied,
the melting point of NaCl(s) being 801oC. The separated ions can be represented by
the symbols Na+(l) and Cl−(l). At the negatively charged electrode there is an
abundance of electrons. In the melt, the ions Na+(l) and Cl−(l) are free to move so the
Na+(l) cations move towards the negative electrode and accept electrons to form solid
sodium metal. Likewise, the mobile Cl−(l) anions move to the positively charged
electrode and give up their excess electrons to become the element Cl2(g). These
processes can be represented by the following half equations
Na+(l) + e− Na(s) electrons gained ∴ reduction
2Cl−(l) Cl2(g) + 2e− electrons lost ∴ oxidation
and the overall reaction by the balanced equation
2Na+(l) + 2Cl−(l) 2Na(s) + Cl2(g) redox
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Document Summary
The name oxidation was initially applied to reactions where substances combined with the element oxygen. Thus any substance burning in air was said to be oxidised, the product being some type of oxide. For example, burning carbon to produce carbon dioxide is an oxidation, as shown by the equation. Subsequently it was realised that reactions of substances with elements other than oxygen were of essentially the same type. For example, hydrogen can react with oxygen to form the compound water, but equally it can react with chlorine to form the compound hydrogen chloride. In both reactions the free element hydrogen is converted to a compound of hydrogen and another non-metal, and so both were classed as oxidations even though no oxygen was involved in the second case. The reverse reaction, conversion of compounds such as oxides of metals to the elemental metal were called reduction reactions, for example, the reduction of copper(ii) oxide to copper by heating with charcoal (carbon).
