CHEMICAL REACTIONS ANDD ENERGY CHANGES

A chemical reaction is a chemical change of two or more substances (reactants) in which the products will have
completely different properties from the reactants.

A+B  AB

Reactant product

Reversible reactions Symbol: ⇌

A reversible reaction is a reaction that proceeds in the forward and backward directions. It follows that if the reaction is
exothermic, then the reverse reaction will be endothermic and vice versa.

Example: A + B ⇌ C + D

The double headed arrow shows that the reaction is a reverse reaction

TYPES OF CHEMICAL REACTIONS

1) SYNTHESIS
This is where one or two substances combine to form just one compound e.g. iron reacting with Sulphur to form
iron sulphide
Fe (s) + S(s)  FeS (s)
2) DECOMPOSITION REACTION
This is the breaking down of reactant constituents into simpler substances. For example
2AgCl (s)  2Ag (s) + Cl2 (s)
3) DOUBLE DECOMPOSITION REACTION
This is where two substances break down and they exchange radicals or metals in their compound. E.g.
NaSO4 (aq) + BaCl2 (aq)  NaCl (aq) + BaSO4 (aq)
4) COMBUSTION REACTIONS
This is the burning of substance. This reaction produces light, heat and sound. E.g.
2Mg(s) + O2 (g)  MgO (s)
5) DISPLACEMENT REACTION
This is the type of reaction were an element swaps (displaces) another element of less reactivity e.g.
Na(s) + AgNO3 (aq)  NaNO3 (aq) + Ag (s)
6) PRECIPITATION REACTION
This is a reaction giving an insoluble product
K2CO3 (aq) + 2AgNO3 (aq)  2KNO3 (aq) + AgCO3
7) NEUTRALIZATION REACTION
This is a reaction between an acid and a base to produce salt and water only.
NaOH (aq) + HCl (aq)  NaCl (aq) + H2O (L)
8) REDOX REACTIONS

A redox reaction: is the one where both oxidation and reduction occur simultaneously.

OXIDATION AND REDUCTION

Oxidation and reduction can be defined in terms of oxygen, hydrogen, electrons and oxidation number.

1) Oxidation in terms of oxygen: is the gain in oxygen. Any substance which gains oxygen is oxidized

Oxidizing agent: is a substance which adds oxygen to another substance. An oxidizing agent is always reduced.
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 (I) 2Mg(s) + O2(g) → 2MgO(s)

Magnesium has been oxidized to magnesium oxide since oxygen has been added

(II) 4FeO(s) + 3O2(g) → 2Fe2O3(s)

Iron (II) oxide (FeO) has been oxidized to iron (III) oxide (Fe2O3) since oxygen has been added.

2) Oxidation in terms of hydrogen: is the loss in hydrogen. Any substance which loses hydrogen is oxidized

Examples

(I) 2NH3(g) → N2(g) + 3H2(g)

Ammonia (NH3) has been oxidized to nitrogen (N2) since hydrogen has been removed.

(II) H2S(g) + Cl2(g) → 2HCl(g) + S(s)

Hydrogen sulphide (H2S) has been oxidized to Sulphur (S) since hydrogen has been removed.

Oxidizing agent: is a substance which gains hydrogen. An oxidizing agent is always reduced.

3) oxidation in terms of electrons: is the loss of electrons, Any substance which loses electrons is oxidized

(I) Cu → Cu2+ + 2e

Copper atoms (Cu) have been oxidized to copper (II) ions (Cu2+) since two electrons have been lost.

(II) Al → Al3+ + 3e

Aluminum atoms (Al) have been oxidized to aluminum ions (Al3+) since three electrons have been lost.

(III)2Cl- → Cl2 + 2e

Chloride ions (Cl-) have been oxidized to chlorine molecules (Cl2) since electrons have been lost.

Oxidizing agent: is a substance which gains electrons. An oxidizing agent is always reduced.

4) Oxidation in terms of oxidation number: is the increase in oxidation number. Any substance whose oxidation
state increases is said to be oxidized.

Oxidation number: Oxidation number (oxidation state) is the total number of electrons that an atom either gains or
losses in order to form a chemical bond with another atom.

Rules to consider when assigning the oxidation number

1. The oxidation number of neutral particles like atoms and molecules is equal to Zero.

Examples

Cu0 Oxidation number = 0

Al0 Oxidation number = 0
Na0 Oxidation number = 0
Oxidation number = 0
Oxidation number = 0

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2. The oxidation number of hydrogen in all compounds except metallic halides is +1

3. The oxidation number of oxygen in all compounds except in peroxides and in OF2 is -2.

4. For some metals the oxidation number is equal to the group number on the periodic table or the valence.

5. The oxidation number of a simple ion is equal to the charge it carries, example

Ion Ion formula Oxidation number

Aluminium ion Al3+ +3
Calcium ion Ca2+ +2
Nitride ion N3- -3
Oxide ion O2- -2
Sodium ion Na+ +1
Chloride ion Cl- +1

1. In neutral compounds, the sum of individual elements present is equal to zero.

Examples
a) Find the oxidation number of
(i) S in SO2 (ii) S in H2SO4
Solution
(i) x + (-2) 2 = 0
x + (-4) = 0 x – 4 = 0
x = +4
(ii) (1x2) + x + (-2) 4 =0
2+x–8=0 x–6=0
x = +6

7. The sum of all oxidation numbers of all elements in a complex ion is equal to the charge on the ion.

Examples

a) Find the oxidation number of

(i) C in HCO3-

(ii) S in SO4 2-

Solution

(i) 1 + x + (-2)3 = -1
x+1–6=-1
x – 5 = -1
x = +5 - 1
x = +4
(ii) x + (-2)4 = -2
x – 8 = -2
x = +8 - 2
x = +6

Oxidation in terms of oxidation number; Oxidation is the increase in the oxidation number of a substance.

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Examples

(I) CuO → Cu2+ + 2e

Copper atoms (Cu) have been oxidized to copper (II) ions (Cu2+) because the oxidation number has increased from 0 in
Cu to +2 in Cu2+.

(II) Cl− → Clo + e

A Chloride ion (Cl-) has been oxidized to chlorine atom (Cl) since the oxidation number has increased from −1 in Cl− to 0
in Cl.

Oxidizing agents/Oxidants

An oxidizing agent is a chemical substance which brings about oxidation of another substance but end up being
reduced.

Examples of oxidizing agents.

 Oxygen, O2
 Chlorine, Cl2
 Hydrogen peroxide, H2O2
 Potassium permanganate, KMnO4
 Manganese dioxide (manganese (IV) oxide), MnO2
 Concentrated Sulphuric acid, H2SO4

Characteristics of an oxidizing agent.

An oxidizing agent: -

(a) Supplies or donates oxygen to another substance

Example 1

PbO(s)+ H2(g) → Pb(s) + H2O(g)

Lead (II) oxide (PbO) is an oxidizing agent because it has donated oxygen to hydrogen.

(b) Removes hydrogen from another substance

Example 2

H2S(g) + Cl2(g) → 2HCl(g) + S(s)

Chlorine (Cl2) is an oxidizing agent because it has removed hydrogen from hydrogen sulphide (H2S)
(c) Accept electrons from another substance.
(d) Increases the oxidation number of another substance.

Example 3

Zno + Cu2+ → Zn2+ + Cuo

Copper (II) ion (Cu2+) is an oxidizing agent because it has caused the increase in the oxidation number of Zinc from 0 in
Zn to +2 in Zn2+.

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Test for oxidizing agents.

1. When aqueous chlorine is added to aqueous potassium iodide containing starch, a dark blue colorization occurs
indicating the presence of iodine (I2) in solution.

Cl2(aq) + 2KI(aq) → 2KCl(aq) + I2(g)

2. When aqueous copper (II) sulphate is added to aqueous potassium iodide, a dark blue colorization is seen indicating
the presence of iodine (I2) in solution.

2CuSO4(aq) + 4KI(aq) → 2K2SO4(aq) + 2CuI(aq) + I2(s)

REDUCTION

Reducing agent: is a substance which loses hydrogen. A reducing agent is always oxidized

1) Reduction in terms of oxygen: Reduction is the removal of oxygen from a chemical substance.

Examples

(a) CO2(g) + C(s) → 2CO(g)

Carbon dioxide (CO2) has been reduced to carbon monoxide (CO) since oxygen has been removed.

(b) PbO(s) + H2(g) → Pb(s) + H2O(g)

Lead (II) oxide (PbO) has been reduced to lead (Pb) by the removal of oxygen.

(c) 3CO(g) + Fe2O3(s) → 2Fe(s) + 3CO2(g)

Iron (III) oxide (Fe2O3) has been reduced to iron (Fe) by the removal of oxygen.

2) Reduction in terms of hydrogen: Reduction is the addition of hydrogen to a chemical substance.

Examples

(a) N2(g) + 3H2(g) → 2NH3(g)

Nitrogen (N2) has been reduced to ammonia (NH3) since hydrogen has been added.

(b) C2H4(g) + H2(g) → C2H6(g)

Ethene (C2H4) has been reduced to ethane (C2H6) by the addition of hydrogen

(c) 2C(s) + H2(g) → C2H2(g)

Carbon (C) has been reduced to ethyne (C2H2) by the addition of hydrogen.

3) Reduction in terms of electrons: Reduction is the gain of electrons.

Examples

(a) Cu2+ + 2e → Cu

Copper (II) ions (Cu2+) have been reduced to copper atoms (Cu) by gaining two electrons.

(b) Cl2 + 2e → 2Cl-

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 Chlorine molecules (Cl2) have been reduced to chloride ions (Cl-) by gaining two electrons.

4) Reduction in terms of oxidation number: Reduction is the decrease in the oxidation number of a chemical
substance.
Examples
(a) Zn2+ + 2 e → Zno
Zinc ions (Zn2+) have been reduced to Zinc atoms (Zn) since the oxidation number has decreased from +2 in Zn2+
to 0 in Zn.
Reducing agents/ Reductants

A reducing agent is a chemical substance which brings about reduction but end up being oxidized.

Examples of reducing agents.

 Hydrogen, H2
 Carbon monoxide, CO
 Carbon, C
 Ammonia, NH3
 Hydrogen sulphide, H2S
 Sulphur dioxide, SO2

Characteristics of a reducing agent.

A reducing agent:

(a) Removes oxygen from another substance

Example
3CO(g) + Fe2O3(s) → 2Fe(s) + 3CO2(g)

Carbon monoxide (CO) is a reducing agent because it has removed oxygen from iron (III) oxide (Fe2O3)

(b) Donates or supplies hydrogen to another sustenance

Example
H2S(g) + Cl2(g) → 2HCl(g) + S(s)
Hydrogen sulphide (H2S) is a reducing agent because it has donated hydrogen to chlorine.
(c) Donates electrons to another substance
(d) Decreases the oxidation number of another substance
Example
Zno + Cu2+ → Zn2+ + Cuo
Zinc atom (Zn) is a reducing agent because it has caused the decrease in the oxidation number of copper (II) ions
from +2 in Cu2+ to 0 in Cu.
Test for reducing agents.
Reducing agents can be tested by using acidified potassium per manganate or acidified potassium dichromate
(VI)
1. Reducing agents change the colour of the solution of potassium permanganate (VII) from purple to colourless.
MnO4-(aq) + 8H+(aq) + 5e → Mn2+(aq) + 4H2O(l)
(Purple) (Colourless)
2. Reducing agents change a solution of acidified potassium dichromate (VI) from orange to green.
Cr2O72-(aq) + 14H+(aq) + 6e → 2Cr3+ + 7H2O(l)
(Orange) (Green)
In both tests, the electrons come from the reducing agent.
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Redox reactions Reduction and oxidation reactions are called redox reactions because they occur simultaneously.
When one substance is reduced, the other one is oxidized.

Examples of redox reactions

1. Fe2O3 (s) + 2Al (s)  2Al (s)+ Al2O3(s) + 2Fe (s)

Iron (III) oxide (Fe2O3) has been reduced to iron (Fe) while aluminium (Al) has been oxidized to aluminium oxide, Al2O3.

2. PbO (s) + H2 (g)  Pb (s) + H2O (l)

Lead (II) Oxide (PbO) has been reduced to lead (Pb) while hydrogen (H2) has been oxidized to water (H2O)

3. Zn + Cu2+  Zn2+ + Cu

Zinc (Zn) has been oxidized to zinc ion (Zn2+) while copper (II) ion (Cu2+) has been reducedd to copper (Cu)

4. H2S (g) + Cl2 (g)  2HCl (g) + S (s)

Hydrogen sulphide (H2S) has been oxidized to Sulphur (S) while chlorine molecules (Cl2) have been reduced to hydrogen
chloride (HCl)

ENERGY FROM CHEMICAL REACTIONS

All chemicals have energy stored in the form of chemical energy. During a chemical reaction, changes to the stored
chemical energy occurs, this changes chemical energy into other forms of energy such as heat, light or electrical energy.
Reactions are described as either exothermic or endothermic, depending on whether the energy is absorbed or given
out to the surroundings.

Endothermic reaction

This is a reaction which absorbs heat energy from the surroundings resulting in a temperature drop in the surrounding.
The temperature of the surrounding decreases and the container becomes colder. In an endothermic reaction the total
energy of the products is greater than that of the reactants. Bond breaking is therefore an endothermic reaction.

Examples of endothermic reactions

(a) Photosynthesis
6CO2 + 6H2O → C6H12O6 + 6O2
(b) Photography in which light energy helps to decompose silver salt to silver on the photographic plate. The
essential reaction in photography is the reduction of silver ions to metallic silver. With exposure to light
energy, the silver salts decompose into ions as follows:
AgBr  Ag+ + Br-
Ag+ + e  Ag
Light is usually absorbed to dissociate silver salts
(c) Thermal decomposition of carbonate salts
CaCO3 (s)  CaO (s) + CO2(g)
CuCO3 (s)  CuO (s) + CO2 (g)
(d) Thermal decomposition of nitrate salts
2Pb(NO3)2 (s)  2Pb (s) + 4NO2 (g) + O2 (g)
Mg(NO3)2 (s)  2MgO (s) + 4NO2 (g) + O2 (g)

Exothermic reaction

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This is a reaction in which energy is released to the surroundings resulting in an increase in the surrounding
temperature. In an exothermic reaction the total energy of the products is lower than that of the reactants due to the
loss of energy to the surrounding.

When two atoms are joined together to form a chemical bond, heat energy is given out. Bond making is therefore an
exothermic reaction.

Examples of exothermic reactions

(a) All combustion or burning processes e.g. burning of fuels such as coal, oils, wood.

(b) Tissue respiration in all living organisms

C6H12O6 + 6O2 → 6CO2 + 6H2O + energy

(c) Dissolving NaOH crystals in water

NaOH (s)  NaOH (aq)

(d) Neutralization reaction

NaOH (aq) + HCl (aq)  NaCl (aq) + H2O (L)

(e) Converting anhydrous salt to hydrated salt

CuSO4 (s) + 5H2O (L)  CuSO4.5H2O (s)

(f) Haber process- manufacture of ammonia from nitrogen

N2 (g) + H2 (g)  2NH3 (g)

(g) Contact process- production of H2SO4. Heat energy is given off when Sulphur dioxide is oxidized to Sulphur trioxide.

2SO2 + O2 (g)  2SO3 (g)

(h) Formation of hydrogen chloride from its elements

H2 (g) + Cl2 (g)  2HCl (g)

COLLISION THEORY

The collision theory states that: For a reaction to take place, the particles of the reacting substances must move and
collide with each other with a certain amount of kinetic energy. The number of collisions taking place per unit time
depends on the number of particles. If the particles are increased, the number of collisions also increases.

However, not all collisions will result in the formation of products. Therefore, effective collisions are required in a
chemical reaction. Effective collisions only occur when the reactant particles have enough energy to overcome the
activation energy of the reaction and when particles collide in the correct orientation. Activation energy is the
minimum amount of energy required to make the reaction take place.

The speed of any chemical reaction depends on the number of effective collisions between reactants. The greater the
number of effective collisions, the higher the rate of reaction

Rate of chemical reaction

The rate of reaction is the measure of how long (speed) the chemical reaction will take place. Or it is the
change in the concentration of reactants or products in a given period of time.
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Calculations on the rate of chemical reaction
Rate of chemical reaction = Change of amount of substance/ change in Time
Or
Reaction rate = Change in mass or volume or concentration/ Change in time
The speed of any chemical reaction depends on the number of effective collisions between reactants. The
greater the number of effective collisions, the higher the rate of reaction
Factors affecting the rate of reaction

1. Concentration
2. Temperature
3. Pressure
4. Surface area (size of particles)
5. Catalyst

1. Effects of a catalyst on the rate of reaction

A catalyst is a chemical substance which alters the rate of reaction but remains chemically unchanged at the
end of the reaction.
A catalyst usually speeds up the reaction by lowering the activation energy of the reaction. For example, the
volume of oxygen produced from the decomposition of hydrogen peroxide (H2O2) can be measured using a gas
syringe with a catalyst manganese (IV) oxide, a black solid. The addition of manganese (IV) oxide ( MnO2 ) speeds
up the reaction and increases the volume of oxygen formed within a short time.

ACTIVATION ENERGY SYMBOL: Ea

Activation energy is the minimum energy required to start a reaction. As a result, a catalyst allows a reaction to go by a
different pathway with lower activation energy allowing more collisions for a successful reaction.

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 Activation energy is usually the energy barrier because if any collision is not energetic enough, the reaction will be
futile.

Some catalysts slow down the reactions; these are called inhibitors (negative catalysts)

Characteristics of a catalyst

1. It catalyzes both the forward and reverse reaction

2. It undergoes physical change

3. It remains chemically unchanged at the end of the reaction

4. It is only needed in very small amounts

5. It is poisoned or rendered useless in the presence of impurities

Examples of some catalysts used for important reactions

Catalyst Reaction catalyzed

Iron Haber process: production of ammonia
Manganese (IV) oxide Decomposition of hydrogen peroxide
Vanadium Contact process: Manufacture of sulphuric acid
Nickel Hydrogenation of vegetable oils
aluminosilicates Catalytic cracking of hydrocarbons to produce
ethylene and propene
Platinum + alumina Dehydrogenation of alkanes

2. Effects of temperature on the rate of reaction

Temperature is the measure of the average kinetic energy of the particles. When temperature is increased, the
rate of reaction also increases. This because the particle gain kinetic energy and move faster and collide
effectively, on the other hand, when temperature is reduced, the rate of reaction also reduces. This is because
the particles lose kinetic energy and move slower and do not collide effectively.
Note
Temperature increases the rate of reactions for endothermic reactions

3. Effects of pressure on the rate of reaction

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 Pressure becomes a dominant factor in reactions involving gases when pressure is increased, the rate of reaction
also increases. This is because the volume reduces forcing the gas particles closer together and collides
effectively on the other hand, when pressure is reduced, the rate of reaction also reduces. This is because the
volume increases and the gas particles are further apart and do not collide effectively

4. Effects of concentration on the rate of reaction

Concentration refers to the reactants in solution. When concentration is increased, the rate of reaction also
increases. This is because the number of particles in the solution increases and collides with each other
effectively. On the other hand, when concentration is reduced, the rate of reaction also reduces. This is because
the number of particles in the solution reduces and do not collide effectively.
Graphical representation of concentration

5. Effects of surface area (particle size) on the rate of reaction

Particle size usually refers to particles of a solid reactant. The rate of reaction is faster when the size of particles
is small. This is because a small sized particle has a large surface area.
Note
i. When a reactant is in solid state, the reaction takes place on the surface of the solid. By breaking up the
solid into smaller pieces, the surface area is increased giving a greater area for collisions to occur. This
results in an increase of the rate of reaction.
ii. This explains why mixtures of saw dust, fine products of flour mills and combustion of gases can cause
an explosion due to the large surface area.
Graphical representation of surface area
Example: Reaction of magnesium with dilute hydrochloric acid.

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 Note: The reaction stops when the reactant in fewer quantities is used up, and hence curve levels off.

6. Light Intensity
Light (when viewed as an electromagnetic wave) is considered to be an energy source and has sufficient
impact energy to break chemical bonds. This energy is more than enough to overcome the activation energy.
The greater the intensity or energy of light, the more reactant molecules are likely to gain kinetic energy, so
the faster the reaction should be. Methane reacts very slowly with chlorine in the dark, but the rate of
reaction is much faster in the presence of ultraviolet light.
Methods of controlling the rate of chemical reactions: by reducing the frequency of collisions between
reacting particles such as in explosions in flour mills or coal mines when ignited to surface area
Measuring the rate of reaction
The rate of reaction can be measured by measuring:
i. how quickly a product is obtained
ii. how quickly a reactant is used up
Example
Consider the reaction below;
CaCO3(s) + 2HCl(aq) → CaCl2(aq) + H2O(l) + CO2(g)
In the reaction above, the rate of reaction can be measured by measuring:
a. The volume of carbon dioxide over time
b. The decrease in mass of the system due loss of carbon dioxide
[A] Measuring the rate of reaction by measuring the volume of the gas produced
A graduated syringe is used to measure the volume of carbon dioxide gas formed over time.

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The total volume of carbon dioxide given off at one minute interval is recorded and the graph of total volume
of carbon dioxide against time is plotted.

Things to note from the graph

a. The gradient of the graph can be calculated, the gradient of the graph is equal to the rate of reaction.
The gradient of the graph at various points of the curve will give the rate of reaction. The reaction is
fastest at the start because the gradient of the graph is the highest. The value of the gradient
decreases with time and finally becomes zero. This means that as the reaction proceeds, the reaction
slows down and finally comes to a stop.
Gradient= Y2 –Y1/ X2-X1
Rate of reaction = change in reactants or products/ change in time
b. When the curve levels off, it means the reaction has stopped
[B] Measuring the rate of reaction by measuring the decrease in mass of a system
A mass balance is used to follow the loss in mass of a system

The mass readings will drop over time as the carbon dioxide gas formed escapes. The mass readings are taken
at intervals and plotted against time

Note: The cotton wool is used as a stopper. It will allow the escape of carbon dioxide into the atmosphere and
prevent the solution inside the conical flask from splashing out.

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