Chemistry Handout 12 REF #: 012

Reduction and Oxidation

Redox reactions are chemical reactions in which oxidation and reduction takes places simultaneously. Oxidation and
reductions are reactions that involve the loss and gain of electrons.

Reduction- A substance is reduced when it: Oxidation- A substance is oxidized when it:

I. Combines with hydrogen. I. Combines with oxygen.

II. Loses oxygen with which it was combined. II. Loses hydrogen with which it was combined.
III. Gains one or more electrons. III. Loses one or more electrons.
IV. Decrease in oxidation number in its free state or within IV. Increase in oxidation number in its free state or
a compound. within a compound.

Examples of Redox Reactions in daily life

Combustion- Any time a material burns, an oxidation-reduction reaction occurs. The two equations below show what
happens when coal (which is nearly pure carbon) and gasoline (C 8 H 18 ) burn. You can see that the fuel is oxidized in each
case: C + O 2 → CO 2 2 C 8 H 18 + 25 O 2 → 16 CO 2 + 18 H 2 O

In reactions such as these, oxidation occurs very rapidly and energy is released. That energy is put to use to heat homes
and buildings; to drive automobiles, trucks, ships, airplanes, and trains; to operate industrial processes; and for
numerous other purposes.

Rusting- Most metals react with oxygen to form compounds known as oxides. Rust is the name given to the oxide of iron
and, sometimes, the oxides of other metals. The process by which rusting occurs is also known as corrosion. Corrosion is
very much like combustion, except that it occurs much more slowly. The equation below shows perhaps the most
common form of corrosion, the rusting of iron. e.g 4 Fe + 3 O 2 → 2 Fe 2 O 3

Decay- The compounds that make up living organisms, such as plants and animals, are very complex. They consist
primarily of carbon, oxygen, and hydrogen. A simple way to represent such compounds is to use the letters x, y, and z to
show that many atoms of carbon, hydrogen, and oxygen are present in the compounds. When a plant or animal dies, the
organic compounds of which it is composed begin to react with oxygen. The reaction is similar to the combustion of
gasoline shown above, but it occurs much more slowly. The process is known as decay, and it is another example of a
common oxidation-reduction reaction. The equation below represents the decay (oxidation) of a compound that might
be found in a dead plant: C x H y O z + O 2 → CO 2 + H 2 O

Biological processes- Many of the changes that take place within living organisms are also redox reactions. For example,
the digestion of food is an oxidation process. Food molecules react with oxygen in the body to form carbon dioxide and
water. Energy is also released in the process. The carbon dioxide and water are eliminated from the body as waste
products, but the energy is used to make possible all the chemical reactions that keep an organism alive and help it to
grow.

Oxidation numbers or Oxidation states

The oxidation number or oxidation state of an atom is equal to the electric charge the atom would have if the compound
the atom was in was totally ionic. Oxidation number is assigned to an ion in an ionic compound or an atom in a
molecule.

Total Transfer of Electrons- For ions, oxidation state is the number of electrons lost or gained when the ion is formed
from its element. Metals lose electrons (it is oxidized), non-metal gain electrons (it is reduced)
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 Chemistry Handout 12 REF #: 012
Partial Transfer of Electrons- For atoms in covalent compounds, the oxidation state is the number of electrons than an
atom partially gains or loses in a molecule. When two non-metal atoms share electrons, the molecule formed may be
polar (carries a charge). The electrons are pulled closer to the more electronegative atom (reduction), whiles the less
electronegative atom is oxidized. Eg- Nitrogen is more electronegative than hydrogen. Nitrogen is reduced and hydrogen
is oxidized.

Note: The oxidation state does not represent the actual charge on the atom. In describing oxidation states, the sign
always precedes the number. Example: The oxidation state of oxygen is written as -2 whiles the formal charge is written
as 2-, the oxidation state of Fe in Fe3+ is +3 whiles the formal charge is 3+.

Element Oxidation Number Notes

Group 1 metals Always +1
Group 2 metals Always +2
Hydrogen +1 Except in hydrides of metals e.g -1 in NaH and CaH 3
Cholrine Except when present in radicals e.g +1 in ClO -, +3 in ClO2-, +5 in ClO3-,
Bromine BrO3-, IO3-
Iodine -1
Oxygen -2 Except in peroxides e.g -1 in H2O2 and Na2O2
Sulphur -2 Except when present in radicals or covalent compounds. E.g +4 in SO 32-
and SO2, +6 in HSO4-, SO42- and SO3
Nitrogen -3 Except when present in radicals or covalent compounds. E.g -3 in NH 4+
and NH3, +2 in NO, +3 in NO2-, +4 in NO2 and +5 in NO3-
Carbon Varies e.g -4 in CH4; +2 in CO; +4 in HCO3- CO32- and CO2
Transition metals when Vary Oxidation number of metal appears in the name of the ion e.g Cr 2O72- is
present in radicals the dichromate (VI) ion- oxidation number of chromium is +6

Rules for Assigning Oxidation Number

1. The oxidation number of all atoms of an element in its free state is zero.
2. The oxidation number of the element as a simple monatomic ion is equal to the charge on the ion.
3. The sum of the oxidation numbers of all the atoms in a polyatomic ion equals the charge on the ion.
4. The oxidation number of elements in radicals or in covalent compounds may vary; often appear in the name of
the ion or compound.
5. The sum of all oxidation numbers of elements in a compound is equal to zero.

Questions: Find the oxidation state of

a) S in H2SO4 e) S in SO32-

b) N in Mg(NO3)2 f) V in VO2+

c) Cr in CrO42- g) N in N2H4

d) Mn in Mn2O7 h) P in H3PO3

Oxidizing and Reducing Agents

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 Chemistry Handout 12 REF #: 012
An oxidizing agent is a substance which brings about oxidation of another substance. In the process, the oxidizing agent
is reduced (it accepts electrons).

A reducing agent brings about the reduction of another substance. In the process the reducing agent is oxidized (it
donates (loses) electrons).

Half Equations/ Electron half equations/ Ionic half equations

Half equations show the transfer of electrons during a redox reaction. There is an oxidation half equation and a reduction
half equation. Half equations are not true chemical equations. A half reaction does not occur by itself, at least two such
reactions must be coupled so that the electron released by one reactant is accepted by another in order to complete the
reaction. 

Questions:

Write balanced ionic-half equations for the following changes. Then determine whether oxidation or reduction occurs.

a) Chlorine to chloride ions e) Tin (IV) to tin (II) ions h) Iodide ions to iodine
b) Oxygen to oxide ions f)Chromium (II) ions to chromium
c) Aluminum ions to aluminum (III) ions
d) Magnesium to magnesium ions g) Calcium to calcium ions

Redox Equations and Half Equations

Rules for Recognizing Redox Reactions:

 Write a balanced Equation for the reaction

 Assign an oxidation number to each different element in its free state or combined within a compound.
 Radicals which remain unchanged during a reaction may be ignored
 Determine the element whose oxidation number has increased- this has been oxidized by the other reactant.
The other reactant is therefore, the oxidizing agent.
 Determine the element whose oxidation number has decreased- this has been reduced by the other reactant.
The other reactant is therefore, the reducing agent.
 If the oxidation number of all elements remains unchanged, the reaction is not a redox reaction.

Questions: Determine the specie that has been reduced, oxidized, the oxidizing and reducing agent from the equations below.

1. Iron (II) Oxide (s) + Oxygen 3. Copper (II) Oxide (s) + Hydrogen

2. Zinc (s) + Copper (II) Sulphate (aq) 4. H2S (g) + Cl2 (g)  2HCl (g) + S (s)

Which of the following equations represent redox reactions? Identify the oxidizing agent, the reducing agent, the
substance being oxidized and the substance being reduced. Write ionic-half equations where appropriate.

a. Ca (s) + Cl2 (g)  CaCl2 (s) c. Ba(NO3)2 (aq) + Na2SO4 (aq)  BaSO4 (s) + 2NaNO3 (aq)

b. Cl2 (g) + H2S (g)  2HCl (g) + S (s) d. CH4 (g) + 2O2 (g)  CO2 (g) + 2H2O (l)

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e. Ca(OH)2 (aq) + CO2 (g)  CaCO3 (s) + 2H2O (l) g. Cl2O7 (g) + H2O (l)  2HClO4 (aq)

f. Ni (s) + CuSO4 (aq)  NiSO4 (aq) + Cu (s) h. 2CuCl (aq)  CuCl2 (aq) + Cu (s)

Common Tests for Oxidizing and Reducing Agents

To test for an oxidizing agent, mix the substance being tested with a known reducing agent that gives a visible change
when it acts as a reducing agent i.e. it is oxidized. To test for a reducing agent, mix the substance being tested with a
known oxidizing agent that gives a visible change when it acts as an oxidizing agent i.e. it is reduced.

An oxidizing agent will oxidize iodide ions to iodine. A solution of potassium iodide in water with a little starch added
turns black with an oxidizing agent below iodine in the reactivity series, because the iodine produced reacts with the
starch. Another test is to add a solution of ammonium iron (II) sulphate, which contains Fe2+ ions. With an oxidizing
agent (for example, aqueous bromine), iron (III) ions produced, which turn potassium thiocyanate dark red. In dilute
solution the brown colour of the iron (III) ions is not noticeable.

A reducing agent will turn potassium dichromate (VI) from orange to green. Chromium (VI) is reduced to chromium (III),
which is green. A filter paper soaked in an acidic solution of potassium dichromate (VI) turns green when placed in
sulphur dioxide. Another test is to add an acidic solution of potassium manganate (VII). The solution becomes almost
colourless. Manganese (VII) is reduced to manganese (II), which is pale pink. In dilute solution the pink colour is not
noticeable.

Substances that behave as both an oxidizing and reducing agents

Sulphur Dioxide, SO2

In the presence of water, sulphur dioxide is usually a reducing agent; if reacted with a stronger reducing agent, it acts as
an oxidizing agent:

 With Chlorine, it acts as a reducing agent

SO2 (g) + 2H2O (l) + Cl2 (g)  H2SO4 (aq) + 2HCl (aq)

 With hydrogen sulphide, a stronger reducing agent, sulphur dioxide acts as an oxidizing agent

2H2S (g) + SO2 (g)  3S (s) + 2H2O (l)

Acidified Hydrogen Peroxide, (H2O2/H+)

This is usually an oxidizing agent, if reacted with a stronger oxidizing agent, it acts as a reducing agent.

- With Potassium Iodide, solution, it acts as an oxidizing agent, oxidizing the iodide ions to iodine.
- With acidified potassium manganate (VII), a stronger oxidizing agent, it acts as a reducing agent, reducing
purple MnO4- ion to colourless Mn2+ ion.

Hydrogen gas, H2

Hydrogen gas acts as an oxidizing agent when combined with metals and as a reducing agent when it reacts with
nonmetals.

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Chemistry Handout 12 REF #: 012

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