Chemistry Chapter 6: Enthalpy Changes

 CHEMISTRY KEYWORDS:

Exothermic Reaction: Heat energy is released during a reaction. The value of ΔH is negative.

Endothermic Reaction: Heat energy is absorbed during a reaction. The value of ΔH is positive,

Standard Conditions: Pressure of 101 kPa and temperature of 298 K.

Specific Heat Capacity c: Energy needed to raise the temperature of 1g of a substance by 1 K ( 1^oC).

Hess’s Law: Enthalpy change in a chemical reaction is independent of the route by which the chemical reaction takes place as long as the initial and final conditions and states of reactants and products are the same for each route.

Exact Bond Energy: Energy needed to break a specific covalent bond in a named molecule in the gaseous state. Also called the Bond Dissociation Energy or Bond Enthalpy.

Average Bond Energy: Average energy needed to break a specific covalent bond averaged from various molecules in the gaseous state.

 

CHEMISTRY DEFINITIONS:

Enthalpy Change ΔH: Heat energy transferred during a chemical reaction.

Reaction Pathway Diagram: Shows the relative enthalpies of the reactants (on the left) and the products (on the right) and the enthalpy change as an arrow. It may also include the activation energy.

Activation Energy E_A: Minimum energy colliding particles must possess to break bonds to start a chemical reaction.

Standard Enthalpy Change of Reaction ΔH_r: Enthalpy change when the amounts of reactants shown in the stoichiometric equation react to give products under standard conditions.

Standard Enthalpy Change of Formation ΔH_f: Enthalpy changes when one mole of a compound is formed from its elements under standard conditions.

Standard Enthalpy Change of Combustion ΔH_c: Enthalpy changes when one mole of a substance is burnt in excess oxygen under standard conditions.

Standard Enthalpy Change of Neutralisation ΔH_neut: Enthalpy changes when one mole of water is formed by the reaction of an acid with an alkali under standard conditions.

 

CHEMISTRY IMPORTANT NOTES:

Examples of exothermic reactions include:

·       Combustion of fuels

·       Oxidation of carbohydrates in the bodies of animals and plants (respiration)

·       Reaction of water with quicklime (Calcium Oxide)

Examples of endothermic responses include:

·       Decomposition of limestone by heating (all thermal decomposition reactions are endothermic)

·       Photosynthesis (in which energy is supplied by sunlight)

·       Dissolving certain ammonium salts in water

 

ΔH = H_p – H_r

(Enthalpy change = Enthalpy of products – Enthalpy of reactants)

 

Standard Conditions:

·       Pressure of 101 kPa (approximately normal atmosphere pressure)

·       Temperature of 298 K

·       Each substance involved in the reaction is in its normal physical state (solid, liquid, gas)

q = mcΔT

 

SUMMARY:

Exothermic enthalpy changes have negative ΔH values. Endothermic enthalpy changes have positive ΔH values.

Enthalpy changes can be calculated experimentally using the relationship:

Enthalpy change, q = mass of liquid x specific heat capacity x temperature change.

For a mole of a defined substance, this is written ΔH = -mcΔT.

Standard enthalpy change of formation relates to the enthalpy change when one mole of a compound is formed from its elements under standard conditions. Similar definitions can be written for standard enthalpy change of combustion and reaction.

Hess’s law can be used to calculate enthalpy changes for reactions that do not occur directly or cannot be found by experiment.

Bond breaking is endothermic; bond making is exothermic.

Average bond energies are often used because the strength of a bond between 2 particular types of atoms is slightly different in different compounds.