5. energetics

Total chemical energy of a substance

Energy change from energy released when bonds are formed and absorbed when bonds are broken during a chemical reaction

The reaction absorbs heat from the surroundings and causes surrounding temperature to decrease.

Products have higher energy content than reactants.

The reaction releases heat to the surroundings, causing surrounding temperature to increase.

Products have lower energy content than reactants.

lower

Minimum energy that reacting particles must possess to overcome the energy barrier before becoming products

a balanced chemical equation that includes the physical states of all reactants and products and the energy change

Temperature: 298K (different from gaseous state)
Pressure: 1 bar
Concentration of solution: 1.0 moldm^-3

State in which substance is most stable at 298K and 1 bar

0

Heat change at constant pressure and under standard conditions

The enthalpy change when molar quantities of reactants as specified in the chemical equation react to form products at 1 bar and 298K.

The enthalpy change when 1 mole of a substance is formed from its constituent elements at 1 bar and 298K

The heat evolved when 1 mole of a substance is completely burnt in excess oxygen at 1 bar and 298K

The heat evolved when 1 mole of water is formed from the neutralisation between an acid and a base at 1 bar and 298K

Weak acidsases do not completely ionise in aqueous solutions, and energy has to be absorbed to ionise un-ionised weak acidase, and less energy is released.

q = mcΔT

q (J)

q = CΔT

q (J)

Amount of heat required to raise the temperature of 1g of a substance by 1K

Amount of heat needed to raise the temperature of a substance by 1 unit

ΔH = q/n

positive for endo, negative for exo
n = moles

an instrument used to measure changes in thermal energy

1. No heat loss or gain from surrounding air
2. Heat capacity of calorimeter is omitted
3. Assume density of all solutions to be the same as water (1.00 gcm^-1)

Enthalpy change of a reaction is determined by only the initial and final states and is independent of the reaction pathway taken.

ΔH = ΣmΔH(combustion of reactants) - ΣnΔH(combustion of products)

m, n being number of moles

ΔH = ΣmΔH(formation of products) - ΣnΔH(formation of reactants)

m, n being number of moles

Energy required to break 1 mole o f a covalent bond in gaseous state

ΔH = ΣΔH(bonds broken) - ΣΔH(bonds formed)

Bond energy data in data booklet are average values and might not be the same as the actual experiment.

Energy required to form 1 mole of gaseous atoms from the element at 298K and 1 bar.

Energy required to convert 1 mole of a compound into gaseous atoms at 298K and 1 bar.

Heat evolved when 1 mole of solid ionic compound is formed from its gaseous constituent ions.
Always negative !

Charge of ions and sizes of ions

LE ∝ (q+ x q-)/(r+ + r-)

Energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of singly charged gaseous cations

Enthalpy change when 1 mole of electrons is added to 1 mole of gasoues atoms to form 1 mole of singly charged gaseous anions

1st EA is usually negative, but 2nd and subsequent EAs are positive.

Energy is required to remove electrons

1st EA is usually negative as effective nuclear charge of the atom attracts the electron, and stronger electron results in greater energy released.

Subsequently, energy is required to overcome the electrostatic repulsion between the electron and anion, so it is positive.

The theoretical lattice energy is calculated based on the assumption that the compound is completely ionic, but the compound could have covalent character which causes a difference in results.

Cation with high charge density, anion with large electron cloud. The electron cloud of the anion would be distorted to a larger extent.

Heat evolved one 1 mole of free gaseous ions are dissolved in an infinite volume of water at 298K and 1 bar.

Na+ (g) + aq -> Na+(aq)

Heat is evolved in forming ion-dipole interactions between water molecules and ions

q

More charge density = stronger ion-dipole interactions = more exothermic

Enthalpy change when 1 mole of solute is completely dissolved in an infinite volume of solvent at 298K and 1 bar.

Positive: salt is likely insoluble
Negative: salt is likely soluble

Note that entropy change has to be taken into account too though

1. Break up ionic lattice to form gaseous ions
2. Hydration of gaseous ions

Enthalpy change of solution = - LE + enthalpy change of hydration
(298K 1 bar)

A reaction that takes place naturally in the direction stated, without continuous input of energy

Change in the freedom of motion of particles in a system

Measure of randomness or disorder in a system, reflected by the number of ways particles and energy can be distributed.

There are more ways to distribute particles and their energy

Change in temperature
Change in phase
Change in number of (usually gaseous) particles
Mixing of particles
Dissolution of an ionic solid

Increase in temperature = entropy increase
Average kinetic energy of the particles and range of energies of particles increase. There are more ways to distribute particles and their energy, so entropy increases.

Increase in particles = positive entropy change
Particles become more disordered and there are more ways that particles and energy can be distributed, so entropy change is positive

Solid to gas
Solid: particles can only vibrate in fixed positions; energy is the least dispersed and it has the least entropy.

As temperature rises, particles gain more kinetic energy and the range of kinetic energies also increase, causing there to be more ways to distribute energy, increasing entropy.

Liquid: when solid melts, particles move more freely in liquid state and become more disordered, so there are more ways to distribute particles and their energy, increasing entropy.
Increase temp = increase in KE and range of energies

Gas: when liquid vaporises, particles are able to move even more freely and there are more ways to distribute particles and energy. Also an increase in disorder as there is an increase in number of gaseous particles.

E.g., a container of 50cm^3 of H2 and another container of 50cm^3 of N2 are joined together to form 100cm^3 of a mixture (gases are at constant pressure). Each gas would expand to occupy the whole 100cm^3 container, increasing volume available for each gas. This increases the number of ways to distribute particles and their energy, increasing entropy.

Entropy increases as ions are able to move freely in water
Entropy decreases as water molecules that were originally free to move become restricted in motion as they have to position themselves around the ions.

Energy change that occurs in going from the reactants to the products

ΔG=ΔH-TΔS

Less than 0: takes place spontaneously. Exergonic
0: At equilibrium (e.g., boiling and melting)
More than 0: not spontaneous. Endergonic

Hess' law can be used !

1. ΔG can only be used to predict spontaneity at 298K and 1 bar
2. Kinetics consideration: the reaction might be happening at a rate too slow to be observed