SSS 2: ENERGY AND CHEMICAL REACTION
Energy can be defined as the capacity to do work.
Forms of Energy
The different forms of Energy are:
1. Mechanical energy
2. Chemical energy
3. Electrical energy
4. Heat energy
5. Nuclear energy
6. Light (Photon) energy
7. Sound energy
8. Magnetic energy
All matter possess energy in one form or the other. When energy changes from one form to another, the total amount of energy before and after the change are always the same. This is the law of conservation of energy.
The law of conservation of energy states that energy can neither be created nor destroyed but it can be changed from one form to another.
Energy changes in chemical reactions
Energy changes occur during chemical reactions because the reactants and products possess different amounts of chemical energy. The forms of energy changes that may accompany a chemical reaction include:
I. Heat: For example, when a strong acid reacts with a strong base.
II. Light and Heat: For example, when magnesium burns in air or in an oxy-ethyne flame.
III. Sound and Heat: For example, when a mixture of hydrogen and oxygen is ignited.
Heat content
Heat content is also called enthalpy (H). It is the energy possessed by a chemical substance.
Heat of reaction
Heat of reaction is also called enthalpy change (ΔH). It is the amount of heat energy absorbed or released during a chemical reaction at constant pressure.
It represents the difference in energy between the products and the reactants:
The unit used for measuring heat change is the joule(J). The common unit is kilojoule(KJ).
The two types of energy changes that accompany a chemical reaction are:
I. Exothermic reaction
II. Endothermic reaction
Exothermic reaction
Exothermic reaction is the type of reaction in which heat is liberated to the surrounding. The heat content of the products (Hₚ) is less than the heat content of the reactants (Hᵣ). The enthalpy change is negative. ΔH < 0
I. Reaction between calcium oxide and water
CaO + H₂O → Ca(OH)₂ : ΔH = negative
II. Reaction between an acid and a base
III. Combustion reaction
IV. Reaction between ammonia and hydrochloric acid
NH₃ + HCl → NH₄Cl: ΔH = -ve
Endothermic reaction
Exothermic reaction is the type of reaction in which heat is absorbed from the surrounding. The heat content of the products (Hₚ) is greater than the heat content of the reactants (Hᵣ). The enthalpy change is positive. ΔH > 0
Examples of endothermic reactions include:
I. Thermal decomposition of calcium trioxocarbonate IV.
CaCO₃ → CaO + CO₂ : ΔH = positive
II. Dissolution of potassium trioxonitrate (V) in water
KNO₃ + H₂O → K⁺ + NO₃⁻ : ΔH = +ve
KNO₃ + H₂O → K⁺ + NO₃⁻ : ΔH = +ve
III. Evaporation of methylated spirit
C₂H₅OH(l) → C₂H₅OH(g) : ΔH = +ve
IV. Thermal dissociation of ammonium chloride
NH₄Cl → NH₃ + HCl : ΔH = +ve
Note:
•The forward reaction is endothermic
•The backward reaction is exothermic
Neutral Heat Change (No Heat Change)
There are some processes that do not involve heat exchange with surroundings.
ΔH = 0
Examples:
I. Dissolving sodium chloride (NaCl) in water (approximately no heat change).
II. Mixing solutions that do not react chemically.
Measurement of heat of reaction
The instrument for measuring the accurate values of ΔH is Calorimeter. The unit used for measuring heat is calorie or kilocalorie.
Calorie
A calorie is the amount of heat energy required to raise the temperature of 1 gram of water by 1°C at standard pressure (usually 1 atm).
Uses of calorie
I. For measuring small amounts of heat in chemical or physical processes.
II. It is often used in food science for energy content of foods (but kilocalorie is more common).
Kilocalorie (kcal)
A kilocalorie is the amount of heat energy required to raise the temperature of 1 kilogram of water by 1°C.
Uses of kilocalorie
It is commonly used to express energy content of food. For example, a slice of bread may contain about 80 kcal.
Types of Calorimeter
The types of calorimeters among others include:
I. Bomb Calorimeter
II. Glass Calorimeter
Bomb Calorimeter
A bomb calorimeter is used to measure the heat of combustion of substances, especially fuels and food samples.
Glass Calorimeter
A glass calorimeter is a simple laboratory calorimeter made of glass or polystyrene used to measure small heat changes in chemical reactions such as neutralization or dissolution.
Standard conditions for heat changes
The value of heat changes, ΔH, is affected by a change in the condition of the reaction which includes:
1. Change in temperature
2. Change in concentration
3. Change in pressure (if reactions involved are gases)
4. Change in physical state
Categories of heat changes
The different categories of heat changes are:
1. Heat of formation
2. Heat of combustion
3. Heat of neutralization
4. Heat of solution
Heat of formation
Heat of formation (ΔHf) is the amount of heat absorbed or evolved when one mole of a compound is formed from its elements in their standard states under standard conditions (25°C, 1 atm).
It is measured in kJ/mol.
Note
Standard elements in their most stable form e.g., O₂(g), H₂(g), N₂(g), C(graphite), S(rhombic) etc have ΔHf = 0.
ΔHf < 0 → exothermic (heat released)
ΔHf > 0 → endothermic (heat absorbed)
Heat of combustion
The heat of combustion (ΔHc) is the amount of heat released when 1 mole of a substance is completely burned in excess oxygen under standard conditions (25 °C, 1 atm).
It is measured in kJ/mol.
Combustion always involves exothermic reactions. ΔHc = -ve
Heat of neutralization
The heat of neutralization (ΔHn) is the amount of heat released or absorbed when 1 mole of water is formed by the reaction of an acid and a base under standard conditions (25 °C, 1 atm).
It is measured in kJ/mol.
For strong acid + strong base, it is typically exothermic and almost constant (~ -57 kJ/mol). ΔHc = -ve
Heat of solution
The heat of solution (ΔHs) is the amount of heat absorbed or released when 1 mole of a substance is dissolved in a Solvent, usually water, to form a solution of infinite dilution.
It is measured in kJ/mol.
It can be exothermic or endothermic, depending on the solute-solvent interaction.
Simple calculations
Example 1: The heat of fusion of sulphur is 64.4 KJmol⁻¹. Calculate the quantity of heat required to fuse:
(a) 4 moles of solid sulphur
(b) 2g of solid sulphur
(c) Determine the mass of sulphur that would required 32KJ of heat for fusion.
[S = 32g]
Solution
(a) 1 mole of sulphur when fused gives 64.4 KJmol⁻¹ of heat.
1 mole of sulphur → 64.4 KJmol⁻¹
4 moles of sulphur →4 × 64.4 KJmol⁻¹
= 257.6 KJmol⁻¹
(b) 1 mole of sulphur → 64.4 KJmol⁻¹
32g of sulphur →64.4 KJmol⁻¹
2g of sulphur →(2/32) × 64.4 KJmol⁻¹
= 4.025 KJmol⁻¹
(c) 64.4 KJmol⁻¹ of heat requires 32g of sulphur for fusion
64.4 KJmol⁻¹ of heat → 32g of sulphur
32KJ of heat → (32/64.4) × 32
→ 15.9g of sulphur
Example 2:
When 5g of liquid water was formed by burning hydrogen gas, 65KJ of heat was absorbed. Calculate the standard heat of formation of water.
Solution
H₂ + ½O₂ → H₂O: ΔH = xKJmol⁻¹
5g of water → 65KJ of heat
18g of water → (18/5) × 65
→ 234KJmol⁻¹
Example 3:
The amount of energy required to change 15g of water into steam is 45KJ. Calculate the standard heat of vapourization.
Solution
H₂O(l) → H₂O(g): ΔH° = xKJmol⁻¹
15g of water was vaporized by 45KJ
15g of water → 45KJ
18g of water → (18/15) × 45KJ
→54 KJmol⁻¹
Example 4:
Calculate the standard heat of formation when carbon is burnt to form 11g of carbon(IV)oxide gas.
C(s) + O₂(g) → CO₂(g) : ΔH = -393 KJmol⁻¹
Solution
C(s) + O₂(g) → CO₂(g) : ΔH = -393 KJmol⁻¹
44g of CO₂(g) → -393 KJmol⁻¹
11g of CO₂(g) → (11/44) × -393 KJmol⁻¹
→ -98.25 KJmol⁻¹
Example 5:
The heat of solution of ammonium trioxonitrate(V) is 25 KJmol⁻¹. Calculate the quantity of heat absorbed when:
(a) 2 moles
(b) 5g of ammonium trioxonitrate (V) solid is dissolved with a large volume of water
[NH₄NO₃ = 80gmol⁻¹]
Solution
(a) 1 mole of NH₄NO₃ →25 KJmol⁻¹
2 moles of NH₄NO₃ →(2/1)×25 KJmol⁻¹
→50 KJmol⁻¹
(b) 80g of NH₄NO₃ →25 KJmol⁻¹
5g of NH₄NO₃ →(5/80)×25 KJmol⁻¹
→1.56 KJmol⁻¹
Example 6:
-57.4KJ of heat was given off when 1 mole of hydrogen ion (H⁺) of a strong acid neutralizes 1 mole of an hydroxyl ion (OH⁻) of a strong base. Calculate the heat evolved when:
(a) 2 moles of H⁺
(b) 0.5 mole of H⁺(aq) reacts with excess alkaline in dilute solution.
(c) Calculate the number of moles of hydrogen ion (H⁺) that would be produced by the liberation of -5.14KJmol⁻¹
Solution
H⁺(aq) + OH⁻(aq)→H₂O(l) : ΔH° = -57.4 KJmol⁻¹
(a) 1 mole of H⁺ → -57.4 KJmol⁻¹
2 moles of H⁺ →2 × -57.4 KJmol⁻¹
= -114.8 KJmol⁻¹
(b) 1 mole of H⁺ → -57.4 KJmol⁻¹
0.5 moles of H⁺ →0.5 × -57.4 KJmol⁻¹
= -28.7 KJmol⁻¹
(c) 1 mole of H⁺ → -57.4 KJmol⁻¹
-57.4 KJmol⁻¹→1 mole of H⁺
-5.14 KJmol⁻¹ →(-5.14/-57.4) × 1 mole
→ 0.1 mole of H⁺ ion
Hess's Law of Heat Summation
The total enthalpy change of a reaction is constant regardless of the route by which the chemical change occurs, provided the conditions at the start of the reaction are the same as the final conditions.
B = A + C
A = B - C
ΔH = ΔH₁ + ΔH₂ + ΔH₃
Examples:
1. Calculate the standard enthalpy of formation of methane given that the complete combustion of one mole of methane liberate 891KJmol⁻¹ and the standard enthalpies for formation of carbon(IV)oxide and water are 393KJmol⁻¹ and 285KJmol⁻¹ respectively all exothermic
Solution
C + O₂→CO₂ : ΔH= -393KJmol⁻¹
H₂ + ½O₂→H₂O : ΔH= -285KJmol⁻¹
CH₄ + 2O₂→CO₂ + 2H₂O:ΔH= -891KJmol⁻¹
Born-Haber Cycle
Bond Energy
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