Energy can be defined as the ability to do work.
There exist various forms of energy, these include, kinetic energy, potential energy, light energy, nuclear energy, heat energy, solar energy e.t.c.
Energy can neither be created nor destroyed but can be converted from one form to another. (law of conservation of mass)
A body/substance at rest possess potential energy. Potential energy is the energy possessed by a body by virtue of its position and
When chemical reactions occur, bonds are broken in the reactants and new bonds are formed in the product and the energy involved is also a form of potential energy.
Kinetic energy on the other hand is the energy possessed when a body is in motion. The atoms and molecules in a substance possess kinetic energy because they are always in motion
Both the kinetic energy and the potential energy of a system, make up the Internal energy (U) of the body / system.
Heat energy and Temperature
When a body is heated, the temperature will rise, this rise in temperature depends on the heat capacity (C) of the body
∆T= ∆Q (∆= delta)
C
∆T = is the rise in temperature Q(K)
∆Q = heat absorbed (J)
C= heat capacity (J/K)
Specific heat capacity (c): - The specific heat capacity of a substance is the heat capacity per unit mass of the substance.
c= C
m
C= mc
Where c= specific heat capacity (J/gK)
m = mass in grammes (g)
∆T = ∆Q/C substituting "cm" for C
∆T= ∆Q mc
Hence
∆Q =mc∆T
ENTHALPY (H)
This is the total heat content of a body/ system.
Every substance possess its own characteristics enthalpy.
When a substance undergoes a chemical reaction, then there will be a change in the enthalpy of the system.
The change in enthalpy ∆H is equal to the enthalpy of the product Hp minus the enthalpy of the reactant Hr
∆H= Hp - Hr
ENDOTHERMIC AND EXOTHERMIC REACTIONS
Chemical reactions are grouped into two as regards heat energy evolved during chemical reactions they are exothermic and endothermic reactions
EXOTHERMIC REACTION: - This is a reaction during which heat is given off to the surrounding.
In this reaction the heat content of the reactants is greater than the heat energy of the product.
Example of exothermic reactions
1. Combustion reactions: - all combustion reactions are exothermic reactions
a) Mg(s) + O2(g) → MgO(s)
b) C(s) + O2(g) → CO2(g)
2). Neutralization reaction
a). HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
b). H2SO4(aq) + Na2CO3(aq) → Na2SO4(aq) + H2O(l)
3) Solubility: - when water is added to some compounds they dissolve releasing a lot of heat in the process, this is observed as the container becomes hot. Examples include
a). dissolving pellets of sodium hydroxide in water
NaOH(s) → NaOH(aq)
b).
ENDOTHERMIC REACTION:- This is a reaction during which heat is absorbed from the surrounding.
Such reactions, the heat content of the product is greater than the heat content of the reactant.
example of endothermic reactions include
1. Decomposition reactions:- most compounds undergo decomposition when heated strongly
a). CaCO3(s) → CaO(S) + CO2(g)
HEAT OF REACTION AND CHEMICAL BONDS
When chemical reactions occur, bonds are broken, atoms rearrange themselves and new
bonds are formed to give new substances (products) Bond breaking requires energy while bond forming evolves
energy.
The minimum amount of energy that is required for a reaction to occur (bond breaking) is called
activation energy. Activation energy is a characteristic of a reaction, that is no two types of reactions possess the same activation energy.
Bond breaking is endothermic (requires/ absorbs energy) while bond forming is
exothermic (gives off energy). Thus, the heat of reaction comes from breaking and forming of chemical bonds.
Heat of reaction is negative [exothermic] when the energy required to break a bond is less
than the energy liberated when a bond is formed.
Heat of reaction is positive [endothermic] when the energy required to break a bond is greater than the energy given off when a bond is formed.
Heat of reaction - This is the amount of heat evolved or absorbed when reactants combine to form products.
HEAT CHANGES IN CHEMICAL REACTIONS
** HEAT OF FORMATION
This is the amount of heat evolved or absorbed when one mole of a substance is formed from
its elements. it is also known as enthalpy of formation
The
standard heat of formation of a substance(∆Hfθ) is the
heat evolved or absorbed, when one mole of the substance is formed from its
elements under standard conditions.
When 1 mole of liquid water is formed from the elements hydrogen and Oxygen the equation of the reaction is
H2(g) + 1/2O2(g) →H2O(1) ∆Hfθ = - 285kJmol-1
Therefore,
∆Hfθ of water = - 285kJmol-1
HEAT OF NEUTRALIZATION
The standard heat of
neutralization ∆Hnθ
is the amount of heat evolved when 1 mole of hydrogen ions, H+,
from an acid reacts with 1 mole of hydroxide ions, OH-, from an
alkali to form 1 mole of water under standard conditions.
Heat of
neutralization is also known as heat of formation of one mole of water from its
ionic components. It is always exothermic
H+(aq) + OH-(aq) → H2O(l) ∆Hnθ = – 57.4kJmol-1
HEAT OF COMBUSTION
The standard heat of combustion of a substance, ∆HCθ;
is the heat evolved when one mole of the substance is burned completely in
oxygen under standard conditions.
A bomb
calorimeter is the apparatus used for the determination of the heat of combustion of a substance.
The following expression can be used to determine the Heat
of combustion of a substance
Heat
of combustion = Heat energy produced
x molar mass
Mass
burnt 1
When
the heat evolved by the burning substance is used to raise the temperature of a
known mass of water, then the expression for heat of combustion can be given
as:
Heat
of combustion = mc∆θ x molar
mass
Mass burnt 1
Where
m = mass of water
C = Specific heat capacity of water
∆θ = change in temperature, that is,
θ2 – θ1
HEAT OF SOLUTION
Standard
heat of solution, ∆Hsθ , is the amount of heat evolved or
absorbed when 1 mole of substance is dissolved in so much water that further
dilution results in no detectable or noticeable heat change at standard temperature and pressure.
Heat of solution can be exothermic or endothermic.
The heat of solution involves two energies I. Latice energy and II. hydration energy
I. Lattice energy is the energy released when one mole of an ionic solid is formed from its gaseous ions (or the energy needed to separate the solid into gaseous ions).
In simple terms, it shows how strongly the positive and negative ions attract each other in an ionic compound. It is endothermic
Example
For sodium chloride:
Na+(g) + Cl-(g) →NaCl(s)
The energy released when these ions come together to form solid NaCl is called its lattice energy.
II. Hydration energy is the energy released when one mole of gaseous ions is dissolved in water and becomes surrounded by water molecules.
In simple terms, it is the energy given out when ions mix with water.
Example
When sodium chloride dissolves in water:
Na+(g) + Cl-(g) →Na+(aq) + Cl-(aq)
The energy released when the ions become hydrated (surrounded by water molecules) is called hydration energy. It is exothermic.
THERMODYNAMICS
Thermodynamics is the study of relationship between heat and
other forms of energy.
A System in thermodynamics is any part of the universe chosen
for thermodynamics consideration, i.e. the physical and chemical phenomenon or
process occurring in a given environment. A system
can be isolated, closed or open.
A Surrounding is the environment in which a reaction or a
process occurs.
The first law of thermodynamics: - this law states that energy can neither be created nor destroyed but
may be converted from one form to another.
In thermodynamics, we represent heat by q and all other forms
of energy are referred to as work denoted by w.
The conditions or state of a chemical system changes when:
i. Heat is evolved or absorbed, and / or
ii. Work is
done on or by the system
In any case, the internal energy, U, of the system is
affected and it changes.
From first law, heat is changed into internal energy of the
system, and it may be represented by the expression
change
in internal energy = Heat absorbed by the system + Work done by the system
i.e. U = q + w
Work done by the system is negative since this lead to
decrease in internal energy, therefore:
U = q - w
For a gaseous system,
w = P V (substituting for w)
U = q - P V
U = H - P V
H = U - P V
SECOND
LAW OF THERMODYNAMIC
The second law of thermodynamic states that a spontaneous
process occurs only if there is an increase in the entropy of a system and its
surroundings.
A Spontaneous reaction is one which can occur by itself without any source of external energy.
Factors which determine the spontaneity (spontaneous) of a reaction are:
i.
enthalpy, H: The heat content of the substances involved
ii.
entropy, S: The
measure of degree of disorderliness or randomness of a substance
iii.
free energy G: The energy which is available for doing work.
ENTROPY (S)
Entropy is defined as the measure of degree of disorderliness or
randomness of a system.
The standard entropy change (∆Sθ) of a system is a state (solid, liquid or gas) function because it depends on the initial and final state of the system. That
is:
∆Sθ
= Sθproducts - Sθreactants
The S.Iunit of is JK-1mol-1
Entropy increases from solid to liquid to gaseous state
because as a substance goes from solid to liquid to gaseous state, the randomness of its particles increases,
that is; ∆Sθ tends to positive.
For a reversible process at constant temperature,
S = H/T
When ∆S is positive, there is increase in entropy. When ∆S is negative there is decrease in the
entropy of a system.
GIBB’S
FREE ENERGY
this is the amount of energy set aside by a body for doing work. The free energy of a system is the energy which is available
for doing work in the system; that is, it is the driving force that brings about a
chemical change.
The standard free energy change (∆Gθ) is a state
function because it depends on the initial and final state of the system. That
is:
∆Gθ
= Gθproducts - Gθreactants
Free energy takes into account the effect of the enthalpy
and entropy factors as represented in the equation below: and so, the relationship between the three factors is shown below.
G = H-TS
For a change at constant temperature,
△G = △H - T△S
NOTE:
1. When △G is negative, the reaction is spontaneous
or feasible.
2. When △G is positive, the reaction is not
spontaneous, unless the resultant effect of both H and
S leads to a net decrease in G
3. When △G is zero, the system is in equilibrium
Example: The
reaction: C(s) + O2(g) → CO2(g)
is carried out at a temperature of 57oC. If the enthalpy change is -5000J and the
entropy change is +15J.Calculate the free energy change
Solution:
△G = △H - T △S
= -5000 - (57 + 273)
x (+15)
= -5000 - 330 x 15
= -5000 - (+4950)
= -5000 - 4950
= -9950J
or -9.950kJ
OBJECTIVE QUESTIONS
1. Heat energy is best defined as
A. energy due to position
B. energy due to motion
C. energy transferred because of temperature difference
D. chemical energy
2. The SI unit of heat energy is
A. calorie
B. joule
C. kelvin
D. watt
3. Heat always flows from
A. colder body to hotter body
B. hotter body to colder body
C. solid to liquid
D. liquid to gas
4. Which of the following is an endothermic process?
A. Burning of wood
B. Respiration
C. Melting of ice
D. Neutralization reaction
5. A reaction that releases heat to the surroundings is called
A. exothermic
B. endothermic
C. reversible
D. equilibrium
6. During an exothermic reaction, the temperature of the surroundings
A. decreases
B. remains constant
C. increases
D. becomes zero
7. The heat required to raise the temperature of 1 kg of a substance by 1°C is called
A. latent heat
B. specific heat capacity
C. heat of reaction
D. enthalpy
8. Which of the following reactions is exothermic?
A. Decomposition of calcium carbonate
B. Photosynthesis
C. Burning of fuel
D. Melting of ice
9. The heat absorbed or released during a chemical reaction is called
A. thermal energy
B. heat of reaction
C. kinetic energy
D. bond energy
10. Which instrument is used to measure heat energy changes in reactions?
A. Thermometer
B. Barometer
C. Calorimeter
D. Hygrometer
11. When ammonium chloride dissolves in water and the solution becomes cold, the process is
A. exothermic
B. endothermic
C. neutral
D. reversible
12. Which of the following requires heat to proceed?
A. Combustion
B. Freezing of water
C. Decomposition of potassium chlorate
D. Neutralization
13. Heat is transferred mainly by all except
A. conduction
B. convection
C. radiation
D. condensation
14. In an endothermic reaction, energy is
A. given out
B. absorbed
C. destroyed
D. converted to mass
15. Which of the following increases the rate of a chemical reaction?
A. Decrease in temperature
B. Increase in temperature
C. Cooling the reactants
D. Removing heat
16. During photosynthesis, energy is
A. released
B. absorbed
C. destroyed
D. ignored
17. The breakdown of calcium carbonate using heat is an example of
A. exothermic reaction
B. endothermic reaction
. combustion
D. neutralization
18. The total heat content of a substance is known as
A. entropy
B. enthalpy
C. pressure
D. volume
19. Which of the following best describes heat?
A. A form of mass
B. A form of matter
C. A form of energy
D. A chemical
20. When heat is added to reactants, the reaction is more likely to
A. slow down
B. stop
C. speed up
D. reverse
SECTION B: Laws of Thermodynamics
21. The First Law of Thermodynamics is based on the principle of
A. conservation of mass
B. conservation of energy
C. entropy
D. heat flow
22. Which statement best describes the First Law of Thermodynamics?
A. Energy can be created
B. Energy can be destroyed
C. Energy cannot be created or destroyed but can be converted
D. Heat always flows from hot to cold
23. The Second Law of Thermodynamics states that heat
A. flows from cold to hot naturally
B. flows from hot to cold naturally
C. cannot be transferred
D. remains constant
24. A machine that converts all heat into work without loss is
A. efficient
B. possible
C. impossible
D. practical
25. The degree of disorder in a system is known as
A. enthalpy
B. entropy
C. energy
D. temperature
26. According to the Second Law of Thermodynamics, the entropy of the universe
A. decreases
B. increases
C. remains constant
D. becomes zero
27. Which law explains why heat engines are not 100% efficient?
A. First law
B. Second law
C. Third law
D. Boyle’s law
28. The Third Law of Thermodynamics states that the entropy of a perfect crystal at absolute zero is
A. maximum
B. minimum
C. zero
D. infinite
29. Which temperature is called absolute zero?
A. 0°C
B. 100°C
C. –273°C
D. 273°C
30. A spontaneous process is one that
A. requires no energy
B. occurs naturally
C. decreases entropy
D. stops heat flow
THEORY QUESTIONS
(a) Define heat energy.
(b) Distinguish between exothermic and endothermic reactions.
(c) Give two examples each of exothermic and endothermic reactions.
(a) What is enthalpy change of a reaction?
(b) Explain the meaning of positive and negative enthalpy change.
(c) Sketch and label an energy profile diagram for:
(i) an exothermic reaction
(ii) an endothermic reaction.
(a) Define activation energy.
(b) Explain why some reactions do not occur at room temperature.
(c) Describe the effect of a catalyst on activation energy.
(a) State Hess’s law.
(b) Explain Hess’s law using a suitable energy cycle.
(c) Give one practical application of Hess’s law.
(a) What is heat of neutralization?
(b) Write a balanced chemical equation for the neutralization of hydrochloric acid with sodium hydroxide.
(c) State the standard heat of neutralization for strong acids and bases and explain why it is nearly constant.
(a) Define heat of combustion.
(b) Write an equation for the combustion of methane.
(c) Explain why heat of combustion values are usually negative.
(a) What is an energy profile diagram?
(b) With the aid of a diagram, explain how a catalyst affects the energy profile of a reaction.
(a) Define bond energy.
(b) Explain how bond energy can be used to calculate the enthalpy change of a reaction.
(c) Calculate the enthalpy change for the reaction:
H₂(g) + Cl₂(g) → 2HCl(g)
(Given: H–H = 436 kJ mol⁻¹, Cl–Cl = 243 kJ mol⁻¹, H–Cl = 431 kJ mol⁻¹)
(a) What is the law of conservation of energy?
(b) Explain how this law applies to chemical reactions.
(a) Define calorimetry.
(b) Describe a simple experiment to determine the heat of reaction using a calorimeter.
(c) State two sources of error in calorimetric experiments.
(a) What is heat of solution?
(b) Explain why dissolving ammonium chloride in water causes a fall in temperature.
(c) Give one practical application of endothermic reactions.
(a) Differentiate between heat and temperature.
(b) State two units of heat energy.
(c) Explain why stirring increases the rate of heat transfer in a chemical reaction.
(a) Define standard conditions for thermochemical measurements.
(b) State two reasons why standard conditions are necessary.
(a) What is the effect of heat on the rate of chemical reaction?
(b) Explain your answer using the collision theory.
(a) Define thermochemistry.
(b) State three areas of application of thermochemistry in everyday life.
16. (a) State the first law of thermodynamic
(b) Calculate: (a) the heat adsorbed by a system when it does 72J of work and its internal energy decreases by 90J
(b) U for a gas that releases 35J of heat and has 128J of work done on it.
