12.0     WHAT IS ENERGY?

Energy is often defined as the ability to do work. It exists in various forms which are interconvertible. When energy changes from one form to another, the total amount of energy before and after the change are always the same (Law of Conservation of Energy).

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         12.1     ENERGY CHANGES IN REACTIONS

Energy changes occur in chemical reactions as reactants changes to products. These changes can be observed as heat, light or sound, and are usually measured in thousands of joules. Therefore, the common unit for measuring heat energy is the kilojoule (kJ) where 1kJ = 1000J.

         12.2     EXOTHERMIC AND ENDOTHERMIC REACTIONS

An exothermic reaction is one during which heat is liberated to the surroundings. Examples include:

• Reaction between an acid and a base.
• Reaction between calcium oxide and water etc.
• Combustion of fuels (e.g., burning of wood, gasoline).
• Dissolution of NaOH in water.
• Reaction of Na with water.

An endothermic reaction is one during which heat is absorbed from the surroundings. Examples are:

• Thermal decomposition of calcium trioxocarbonates(IV).
• Thermal dissociation of ammonium chloride.
• Dissolution of NH₄Cl in water.
• Melting of ice.

         12.3     THERMODYNAMIC PHENOMENA

ENTHALPY – Enthalpy is defined as the heat content of a system, accounting for the energy absorbed or released during physical or chemical changes (e.g., dissolution, phase changes, or reactions). It is represented by the letter, H, then a change in the heat content of the reaction becomes ΔH (change in enthalpy).

Mathematically:

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ΔH = Heat content of products – Heat content of reactants

ΔH is negative for an endothermic reaction (i.e ΔH = – x kJ mol ^{– 1}), meaning the system lost energy. and positive for an exothermic reaction (i.e ΔH = + x kJ mol ^{– 1}), indicating that the system gains energy. A calorimeter is used to determine the accurate value of ΔH.

ENTROPY – Entropy is a measure of the degree of disorder or randomness of a system. The greater the disorder, the higher the entropy i.e processes that increase disorder (like gas expansion or salt dissolution) lead to an increase in entropy. For instance:

• Mixing of Gases – When two different gases mix, they spread out and form a more disordered system. For example, when nitrogen and oxygen are mixed, the entropy increases.
• Dissolution of Salts – When a salt like NaCl dissolves in water, its orderly crystal structure breaks down into individual ions (Na⁺ and Cl⁻), leading to an increase in disorder (increased entropy).

In terms of physical state of matter, the order of increasing entropy (degree of disorder) is as follows:

Solids ˂ Liquids ˂ Gases

• Solids have the lowest entropy.
• Liquids have greater entropy than solids, but lesser than gases.
• Gases have the highest entropy.

Entropy is represented as S, while change in entropy is represented as ΔS. In a reversible reaction:

 ΔS = ΔH / Temperature (T)

If reaction is endothermic, ΔS is positive (increase in entropy), while if reaction is exothermic, ΔS is negative (decrease in entropy). ΔS is measured in J K ^–1 mol ^{– 1}.

GIBB’S FREE ENERGY – Gibbs free energy, G, is that energy which is available for doing work. This is the driving force that brings about a chemical change. Mathematically, change in ΔG is given by ΔG = ΔH – TΔS. For all spontaneous processes, ΔG must be negative.

         12.4     LAWS OF THERMODYNAMICS

THE FIRST LAW OF THERMODYNAMICS – It states that the energy can be converted from one form to another, but it cannot be created nor destroyed. Thus, during a chemical process:

Change in internal energy of a system (ΔU) = Heat absorbed by system (q) + Work done by system (w)

THE SECOND LAW OF THERMODYNAMICS – It states that a spontaneous process occurs only if there is an increase in the entropy of a system and its surroundings, i.e. the change in the total entropy of the system must be positive.

         12.5     SPONTANEITY OF REACTION

A spontaneous process occurs when there is an increase in the entropy of a system and its surrounding.

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CONDITIONS NECESSARY FOR SPONTANEITY OF REACTION

• ΔH must be negative (ΔH < 0).
• ΔS must be positive (ΔS > 0).
• ΔG must be negative (ΔG < 0).

Spontaneity of a reaction as determined by the sign of ΔG:

• ΔG < 0: The reaction is spontaneous.
• ΔG > 0: The reaction is non-spontaneous.
• ΔG = 0: The reaction is at equilibrium.

12.6     ENERGY PROFILE DIAGRAM

An energy profile diagram is a graphical representation of the energy changes that occur during a chemical reaction. It shows the theoretical “energy pathway” of a reaction as it progresses from reactants to products. An energy pathway is a road that a reaction travels to get from one place (the reactants) to its destination (the products). The type of energy that this pathway tracks is potential energy.

EXAMPLE – Given that in an exothermic reaction, ΔH° = −100 kJ mol ^{– 1}, ΔS° = 50 J K ^{– 1} mol ^{– 1} and T = 300K.

• Convert ΔS° to kJ K ^{– 1} mol ^{– 1}.
• Calculate ΔG°.
• Determine if the reaction is spontaneous or not.

SOLUTION

• ΔS° = 50 J K ^{– 1} mol ^{– 1} ÷ 1000 = 0.050 kJ K ^{– 1} mol ^{– 1}.
• ΔG° = ΔH^o – TΔS^o

= −100 kJ mol ^{– 1} − (300K × 0. 050 kJ K ^{– 1} mol ^{– 1})

= −100 kJ mol ^{– 1} – 15 kJ mol ^{– 1}

= −115 kJ mol ^{– 1}

• Since ΔG° is negative, the reaction is spontaneous at 300 K.