Heat of Combustion

Heat of Combustion

1. Heat of combustion – the heat change when one mole of a substance is completely burnt in oxygen under standard conditions.
2. Combustion – redox reaction between substance (fuel) reacts rapidly with oxygen with the production of heat energy.
3. Combustion reaction gives out heat and always an exothermic reaction.
4. Heat evolved in combustion of fuel = Heat absorbed by water.
5. Bomb calorimeter is used to determine the heat of combustion.
6. The more carbon and hydrogen atoms per molecules in a fuel, the more heat that is released when 1 mol of fuel combusts.

There are differences in heats of combustion:

Chemical equation	ΔH (kJ mol-1)
H2(g) + ½ O2(g) –> H2O(l)	-286
C(s) + O2(g) –> CO2(g)	-392
CH4(g) + 2O2(g) –> CO2(g) + 2H2O(l)	-890
CH3OH(l) + 3/2 O2(g) –> CO2(g) + 2H2O(l)	-728
C2H5OH(l) + 3O2(g) –> 2CO2(g) + 3H2O(l)	-1376
C3H7OH(l) + 9/2 O2(g) –> 3CO2(g) + 4H2O(l)	-2016

The selection of suitable fuel:

• Fuel value (the amount of heat energy given out when one gram of the fuel is completely burnt in excess of oxygen): The higher the fuel value, the more energy is released.
• Effect on the environment: Production of soot which caused air pollution. Hydrogen fuel is known as clean fuels (no soot or poisonous gases).
• Cost per gram of fuel.

Fuel values of common fuels

Substances	Fuel value (kJ g-1)
Fruits	2
Egg	6
Coal	14
Glucose	15.5
Dry cow dung	15.5
Sugars	17
Wood	18
Gasoline	34
Butanol	36.6
Kerosene	37
Biodiesel (Vegetable oil)	42.2
Diesel fuel	46
Natural gas	53.6
Hydrogen	143

Qualities of a fuel are based on the following:

• Easily available
• Cheap in cost
• High fuel value
• Do not pollute the environment
• Less storage space

The Existence of Various Energy Sources

1. The Sun: solar cell (still expensive and inefficient)
2. Fossil Fuels: relatively high fuel value and convenient to use (non-renewable and cause greenhouse effect and acid rain)
3. Water: hydroelectric power. It is clean, renewable, convenient and economical to use (high cost of construction and destruction of the surrounding environment)
4. Biomass: plants and droppings of animals – biodiesel (large areas of land to grow plants)
5. Radioactive substances: uranium and plutonium (non-renewable and very destructive if an accident occurs)

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Heat of Neutralisation

Heat of Neutralisation

1. Heat of neutralisation – the heat change when one mole of water is formed from the reaction between an acid and an alkali.
2. Neutralisation – a reaction between an acid reacts with a base (alkali) to form a salt and water.
3. Hydrogen ion from acid reacts with hydroxide ions from alkali to form water. H⁺(aq) + OH^-(aq) –> H₂O(l)
4. Neutralisation reaction gives out heat and always an exothermic reaction.

Example 1: (Strong acids – monoprotic acid and strong alkalis)
Chemical equation: HCl(aq) + NaOH(aq) –> NaCl(aq) + H₂O(l)
Ionic equation: H⁺(aq) + OH^-(aq) –> H₂O(l)
Heat of neutralisation of strong acids and strong alkalis are the same (ΔH = -57.3 kJ mol^-1)
Example 2: (Strong acids – diprotic acid and strong alkalis)
Chemical equation: H₂SO₄(aq) + NaOH(aq) –> Na₂SO₄(aq) + 2H₂O(l)
Ionic equation: 2H⁺(aq) + 2OH^-(aq) –> 2H₂O(l)
Heat of neutralisation of strong acids and strong alkalis are the same (ΔH = -57.3 kJ mol^-1)
Example 3: (Weak acids and strong alkalis)
Chemical equation: CH₃COOH(aq) + NaOH(aq) –> CH₃COONa(aq) + H₂O(l)
Ionic equation: H⁺(aq) + OH^-(aq) –> H₂O(l)
Heat of neutralisation of weak acids and strong alkalis are lower (ΔH = -55.0 kJ mol^-1) than heat of neutralisation of strong acids and strong alkalis (ΔH = -57.3 kJ mol^-1).
Example 4: (Strong acids and weak alkalis)
Chemical equation: HCl(aq) + NH₄OH(aq) –> NH₄Cl(aq) + H₂O(l)
Ionic equation: H⁺(aq) + OH^-(aq) –> H₂O(l)
Heat of neutralisation of strong acids and weak alkalis are lower (ΔH = -51.5 kJ mol^-1) than heat of neutralisation of strong acids and strong alkalis (ΔH = -57.3 kJ mol^-1).
Example 5: (Weak acids and weak alkalis)
Chemical equation: CH₃COOH(aq) + NH₄OH(aq) –> NH₄Cl(aq) + H₂O(l)
Ionic equation: H⁺(aq) + OH^-(aq) –> H₂O(l)
Heat of neutralisation of strong acids and weak alkalis are lower (ΔH = -50.4 kJ mol^-1) than heat of neutralisation of strong acids and strong alkalis (ΔH = -57.3 kJ mol^-1).

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Energy Change during Formation and Breaking of Bonds

Energy Change during Formation and Breaking of Bonds
Bond breaking

• Usually chemical bonds of the reactant.
• Heat energy is absorbed.

Bond forming

• Usually new chemical bonds of the product.
• Heat energy is given out.

Relationship between energy change and the formation and breaking of bonds

• In a chemical reaction, if the heat energy absorbed in bond breaking is lower than the heat energy given out in bond forming, the reaction is an exothermic reaction.
• Example: ΔH (bond breaking) = +600 kJ, ΔH (bond forming) = -800 kJ, ΔH (heat of reaction) = [(+600) + (-800)] kJ= -200 kJ
• In a chemical reaction, if the heat energy absorbed in bond breaking is higher than the heat energy given out in bond forming, the reaction is an endothermic reaction.
• Example: ΔH (bond breaking) = +800 kJ, ΔH (bond forming) = -600 kJ, ΔH (heat of reaction) = [(+800) + (-600)] kJ= +200 kJ

Applications of Exothermic and Endothermic Reaction in Everyday Life
Hot pack

• Contains of anhydrous calcium chloride / anhydrous magnesium sulphate / wet iron powder and sodium chloride / calcium oxide.
• Uses: reduce swelling and mucles or joint sprain.

Cold pack or Ice pack

• Contains of ammonium nitrate / potassium nitrate / sodium thiosulphate.
• Uses: reduce swelling, muscles or joint sprain and reduce fever.

*Berry Important Notes
You must be able to

1. Calculate the number of moles of salt precipitated / metal displaced / water produced / fuel used;
2. Calculate the heat energy released (ΔH); and
3. Calculate the heat of precipitation / heat of displacement / heat of neutralization / heat of combustion.

Heat of Precipitation (Form 4, Chapter 8 Salts)

1. Heat of precipitation – the heat change when one mole of a precipitate is formed from their ions in aqueous solution.
2. Precipitation reaction = double decomposition which is used to prepare insoluble salts.
3. Heat change of a solution = mcθ Joule [m = mass of the solution (g), c = specific heat capacity of the solution (J g^-1˚C^-1), θ = temperature change in the solution (˚C)]
4. Heat change in a reaction, mcθ = n x ΔH
5. Heat of reaction / Heat of precipitation, ΔH = mcθ / n

Example 1:
Chemical reaction: Pb(NO₃)₂(aq) + 2KI(aq) –> PbI₂(s) + 2KNO₃(aq)
Ionic reaction: Pb²⁺(aq) + 2I^-(aq) –> PbI₂(s)
Heat of precipitation of PbI₂ = – Heat change / Number of moles of PbI₂
Example 2:
Chemical reaction: BaCl₂(aq) + Na₂SO₄(aq) –> BaSO₄(s) + 2NaCl(aq)
Ionic reaction: Ba²⁺(aq) + SO₄^2-(aq) –> BaSO₄(s)
Heat of precipitation of BaSO₄ = – Heat change / Number of moles of BaSO₄

Heat of Displacement (Form 4, Chapter 6 Electrochemistry & Form 5, Chapter 3 Oxidation and Reduction)

1. Heat of displacement – the heat change when one mole of a metal is displaced from its salt solution by a more electropositive metal.
2. Heat change of the reaction mixture / Heat energy released / Heat given out in the reaction = mcθ Joule
3. Heat change in a reaction, mcθ = n x ΔH
4. Heat of reaction / Heat of displacement, ΔH = mcθ / n

Example 1:
Chemical equation: Mg(s) + FeCl₂(aq) –> MgCl₂(aq) + Fe(s)
Ionic equation: Mg(s) + Fe²⁺(aq) –> Mg²⁺(aq) + Fe(s)
Example 2:
Chemical equation: Zn(s) + CuSO₄(aq) –> ZnSO₄(aq) + Cu(s)
Ionic equation: Zn(s) + Cu²⁺(aq) –> Zn²⁺ (aq) + Cu(s)

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Thermochemistry

Thermochemistry:

• The Law of Conservation of Energy – energy can neither be created nor destroyed but it can be changed from one form to the other form.
• Exothermic reaction – a chemical reaction that gives out heat to the surrounding.
• Endothermic reaction – a chemical reaction that absorbs heat from the surrounding.
• Surroundings do not involve in the reactions. Example: water, container, the air, solvent and thermometer.
• Heat of reaction – the heat change when the number of moles of reactants in the chemical equation reacts to form products in standard conditions.
• Standard conditions: temperature (25˚C / 298 K), pressure (1 atm), concentration of solution (1.0 mol dm^-3), reactants and products are at their normal physical states.
• Heat of precipitation – the heat change when one mole of a precipitate is formed from their ions in aqueous solution.
• Heat of displacement – the heat change when one mole of a metal is displaced from its salt solution by a more electropositive metal.
• Heat of neutralisation – the heat change when one mole of water is formed from the reaction between an acid and an alkali.
• Heat of combustion – the heat change when one mole of a substance is completely burnt in oxygen under standard conditions.

Exothermic reaction

1. Chemical energy –> Heat energy
2. The heat energy is transferred to the surrounding.
3. Temperature of the surrounding increases.

Example of Chemical Reactions:

• Respiration
• Burning of metal
• Reaction of an alkaline metals (Group 1) with water
• Reaction of a reactive metal with acid
• Neutralisation reaction between acid and alkali
• Reaction of a carbonate with acid
• Combustion of carbon compound
• Displacement reaction of metals
• Rusting of iron

Example of Physical Processes:

• Freezing process
• Condensation process
• Dissolving an alkali in water
• Dissolving an concentrated acid in water

Endothermic reaction

1. Heat energy –> Chemical energy
2. The heat is absorbed from the surrounding.
3. Temperature of the surrounding decreases.

Example of Chemical Reactions:

• Photosynthesis
• Decomposition of nitrate salts
• Decompositon of carbonates salts
• Reaction between acid with hydrogen carbonates

Example of Physical Processes:

• Melting process
• Boiling process
• Sublimation process
• Dissolving of ammonium salts in water
• Dissolving of potassium salts in water
• Dissolving of thiosulphate  in water

Heat of Reaction

1. Enthalpy (H) – absolute energy content of a substance.
2. Change in energy content (ΔH) – absolute energy content cannot be determined, but ΔH can be determined when the reactants are converted to the products.
3. 1 kJ (kilojoule) = 1000 J
4. The unit ΔH  is kJ

Energy Level Diagram

1. ΔH = H _product – H _reactant = negative value.
2. Example: CH₄(g) + 2O₂(g) –> CO₂(g) + 2H₂O(l) ΔH = -890 kJ
3. The value of ΔH is negative = exothermic reaction.
4. ΔH = H _product – H _reactant = positive value.
5. Example: N₂(g) + 3H₂(g) –> 2NH₃(g) ΔH = +91.8kJ
6. The value of ΔH is positive = endothermic reaction.

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