Thermodynamics

Enthalpy change of formation

Enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions.

enthalpy formation of elements is 0

Enthalpy of hydration

Enthalpy change when 1 mol of gaseousions is converted into aqueous ions.

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

Enthalpy change of solution

The enthalpy change when 1 mole of an ionic substance dissolves in enough solvent to form an infinitely dilute solution.

NaCl (s) \---\> NaCl (aq)

Enthalpy change of atomisation

The enthalpy change when one mole of gaseous atoms is formed from an element/compound in standard state.

1/2 Cl2 (g) \---\> Cl (g)

Δformationꝋ is always endothermic as energy is always required to break any bonds between the atoms in the element, to break the element into its gaseous atoms

First electron affinity

The enthalpy change when 1 mole of gaseous 1- ions is made from 1 mole of gaseous atoms.

O (g) + e\- --\> O- (g)

The first electron affinity is always exothermic as energy is released when electrons are attracted to the atoms

Second electron affinity

The enthalpy change when 1 mole of gaseous 2- ions is made from 1 mole of gaseous 1- ions.

However, the second electron affinity of an element can be endothermic as illustrated by oxygen. This is because a large force of repulsion must be overcome between the negatively charged ion and second electron, so energy must be put in.

Bond disassociation enthalpy

The enthalpy change required to break a covalent bond with all species in the gaseous state.

Lattice enthalpy of decomposition

The enthalpy change required to convert one mole of an ionic solid into its gaseous ionic constitutions.

KCl (s) \---\> K+ (g) + Cl- (g)

Lattice enthalpy of formation

Enthalpy change when one mole of a solid ionic compound is formed from its constituent ions in the gas phase.

ΔH lattice enthalpy formation ꝋ is therefore exothermic, as when ions are combined to form an ionic solid lattice there is an extremely large release of energy

A large value suggests the ionic compound is more stable.

It cannot be determined directly so multiple experimental values must be needed to calculate it - Born Haber cycle

What affects lattice energy?

smaller ions \= greater energy

greater charge\= greater energy

TO WORK OUT LATTICE ENTHALPY ADD ALL THE POSITIVE VALUES AND ADD THE FORMATION NEGATIVE VALUE SO ITS POSITIVE ALSO. TAKE AWAY THE NEGATIVE ELECTRON AFFINITY VALUE.

Explain why there is a different between the hydration enthalpies of magnesium and sodium ions? (2)

Magnesium has a higher positive charge than sodium. It is also smaller due to its increased positive nuclear charge. This means it attracts water more strongly.

Suggest why the second electron affinity is endothermic and positive? (1)

The negative X- ion repels the electron

What can a scientists deduce from a comparison of theoretical and actual value of lattice enthalpy?

The model assumes all ions are spherical and are in a lattice calculated value is smaller than a cycle value.

Experimental lattice enthalpy allows covalent interactions and polarisation.

Theoretical lattice enthalpy values only assume ionic interactions no covalent. If there is a difference between the two it means the compound is not perfectly ionic.

Perfect Ionic Model

Ions can be regarded as perfect spheres with no polarisation or covalent interactions.

Only electrostatic attraction is present.

Why does a temperature higher than the actual value calculated get used in a reaction?

To increase the rate of reaction.

Explain why this reaction does not occur at a temperature?

-for a reaction to be feasible the gibbs free energy should be 0 or less than 0

-at the temperature the gibbs free energy value is positive and above 0 so the reaction is not feasible

How can the conditions of a reaction be changed to allow a reaction with an unfeasible reaction temperature (too high) to occur?

the temperature decreases

ΔG will become (more) negative because

TΔS is less positive

enthalpy of hydration equation

ΔH(solution) \= ΔH(lattice) + Σ(ΔH hydration)

why is hydration an exothermic process for chloride ions? (2)

water is polar - due to delta hydrogen +

chloride ion is attracted to the delta hydrogen +

force of attraction between the chloride ion and water

in enthalpy of hydration questions/ enthalpy of solution

if 2333k is above the boiling point of water the water would evaporate as the temperature is too high

equation for entropy

from gibbs

ΔS \= (ΔH - ΔG) / T

for free energy graph questions the reason there is not data for some temperatures is because there is a change in state

what are the two values needed to calculate enthalpy of solution

-lattice dissociation enthalpy

-enthalpy of hydration

what does enthalpy of solution measure

it is a measure of energy released when attraction between ions and water molecules form

Is enthalpy of solution endothermic or exothermic?

both

Is Enthalpy of hydration exothermic or endothermic?

exothermic

why is the enthalpy of hydration for chloride more negative than bromide? (3)

-chloride is a smaller ion

-force of attraction between chloride ions and water is stronger

- negative chloride ions is attracted to the delta positive hydrogen

Explain why magnesium ions attract water

Water is polar O on water has a delta negative charge

Magnesium ions are positive so are attracted to the delta negative charge on the oxygen

Suggest why a value for the enthalpy of solution of magnesium oxide is not found in any data books.

magnesium oxide dissolves in water to produce mg(oh)2

why does fluorine have a higher electronegativity than chlorine

it is smaller

the bonding pairs are more strongly attracted to the nucleus

why does the entropy value for dissolving silver fluoride in water have a positive value? (1)

produces 2 products which increases the amount of disorder

explain why the dissolving of silver fluoride in water is a spontaneous process

the entropy change is positive and enthalpy change negative

gibbs energy is always negative and below 0

Hess's Law

The total enthalpy change of a reaction is independent of the route taken.

Enthalpy change definition

The heat energy transferred in a reaction at constant pressure

equation for enthalpy of combustion

reactants - products

where do arrows point for combustion enthalpy

down

equation for enthalpy of formation

products - reactants

where do arrows point for formation enthalpy

up

equation for mean bond enthalpy (hess's law)

reactants - products

where do arrows point for mean bond enthalpy

down

Suggest why the entropy of water is zero at 0 K

Particles are in maximum state of order

what two things occur when a solid ionic lattice is dissolved in water?

-bonds between the ions break to give gaseous ions

(lattice enthalpy of dissociation) (endothermic)

-bonds between ions and water are made

(enthalpy change of hydration) (exothermic)

is lattice enthalpy of dissociation exo or endothermic?

endothermic

is enthalpy change of hydration exo or endothermic?

exothermic

is enthalpy change of solution exo or endothermic?

both

for enthalpy change of hydration add the positive ions and negative ions values together (there should be two values)

equation for enthalpy of solution

lattice dissociation enthalpy + enthalpies of hydration \= enthalpy change of solution

why can water form bonds in enthalpy of solution

it is polar

positive ions of the metal form bonds with the - of the oxygen

negative ions of the non metal form bonds with the \= of the hydrogen

Entropy

A measure of disorder.

The number of different ways particles can be rearranged and energy shared out by the particles.

What affects entropy?

-physical state

-dissolution

-number of particles

-more complexed molecules

entropy equation

ΔS total \= ΔS products -ΔS reactants

Units for entropy change

J/mol

entropy change of surroundings equation

ΔS surroundings \= -ΔH/T

What must happen for a reaction to be feasible

total entropy must be positive ΔS+

gibbs free energy value must be negative ΔG-