Thermodynamics

What does Hess's Law
state?

The enthalpy change for a reaction is
independent of the route taken

Define standard enthalpy of
formation.

The enthalpy change when one mole of a
compound is formed from its constituent
elements in standard conditions, with all products
and reactants in their standard states.

What is the standard
enthalpy of an element?

Zero, by definition.

Define standard enthalpy of
combustion

The enthalpy change when one mole of a
substance is completely burnt in (excess) oxygen

Define standard enthalpy of
atomisation

Enthalpy change when one mole of gaseous
atoms is formed from a compound in its standard
state in standard conditions.

Define first ionisation energy

Enthalpy change when one mole of
electrons is removed from one mole of
gaseous atoms to form one mole of
gaseous 1+ ions.

Define second ionisation
energy.

Enthalpy change when one mole of electrons is
removed from one mole of gaseous 1+ ions to
form one mole of gaseous 2+ ions

Define first electron affinity

Enthalpy change when one mole of gaseous
atoms gains one mole of electrons to form one
mole of gaseous 1- ions.

Define second electron
affinity.

Enthalpy change when one mole of gaseous 1-
ions gains one mole of electrons to form one
mole of gaseous 2- ions

Define lattice enthalpy of
formation

Enthalpy change when one mole of solid ionic
lattice is formed from its constituent gaseous
ions.

Define lattice enthalpy of
dissociation.

Enthalpy change when one mole of solid ionic
lattice is dissociated (broken into) into its
gaseous ions

Define enthalpy of
hydration

Enthalpy change when one mole of gaseous ions
become hydrated/dissolved in water to infinite
dilution [water molecules totally surround the ion]

Define enthalpy of solution

Enthalpy change when one mole of solute
dissolves completely in a solvent to infinite
dilution.

Define mean bond
dissociation enthalpy

Enthalpy change when one mole of (a certain
type of) covalent bonds is broken, with all
species in the gaseous state

Write example equations
for:
Standard enthalpy of formation
Standard enthalpy of combustion
Standard enthalpy of atomisation
First ionisation energy
Second ionisation energy
First electron affinity

Standard enthalpy of formation Mg (s) + ½ O 2 (g) → MgO (s)
Standard enthalpy of combustion CH 4 (g) + 2O 2 (g) → CO 2 (g) + 2H 2 O (g)
Standard enthalpy of atomisation 1/2I 2 (g) → I (g)
First ionisation energy Li (g) → Li + (g) + e -
Second ionisation energy Mg + (g) → Mg 2+ (g) + e -
First electron affinity Cl (g) + e - → Cl - (g)

Write example equations
for:
Second electron affinity
Lattice enthalpy of formation
Lattice enthalpy of dissociation
Enthalpy of hydration
Enthalpy of solution
Mean bond dissociation enthalpy

Second electron affinity O - (g) + e - → O 2- (g)
Lattice enthalpy of formation Na + (g) + Cl - (g) → NaCl (s)
Lattice enthalpy of dissociation NaCl (s) → Na + (g) + Cl - (g)
Enthalpy of hydration Na + (g) → Na + (aq)
Enthalpy of solution NaCl (s) → Na + (aq) + Cl - (aq)
Mean bond dissociation enthalpy Br 2 (g) → 2Br (g)

What is a Born-Haber
cycle?

Thermochemical cycle showing all the enthalpy
changes involved in the formation of an ionic
compound. Start with elements in their standard
states (enthalpy of 0)

What factors affect the
lattice enthalpy of an ionic
compound?

Size of the ions, charge on the ions

How can you increase the
lattice enthalpy of a
compound? Why does this
increase it?

Smaller ions, since the charge centres will be closer together.
Increased charge, since there will be a greater electrostatic
force of attraction between the oppositely charged ions. N.B.
Increasing the charge on the anion has a much smaller effect
than increasing the charge on the cation, since increasing
anion charge also has the effect of increasing ionic size.

How can Born-Haber cycles
be used to see if
compounds could
theoretically exist?

Use known data to predict certain values of
theoretical compounds, and then see if these
compounds would be thermodynamically stable.
Was used to predict the existence of the first
noble gas containing compound.

What actually happens
when a solid is dissolved in
terms of interactions of the
ions with water molecules?

Break lattice → gaseous ions

dissolve each gaseous ion in
water. The aqueous ions are surrounded by water molecules
(which have a permanent dipole due to polar O-H bond)

What is the perfect ionic
model?

Assumes that ions are perfectly spherical and
that there is an even charge distribution (100%
polar bonds). Act as point charges.

Why is the perfect ionic
model often not accurate?

Ions are not perfectly spherical. Polarisation often
occurs when small positive ions or large negative
ions are involved, so the ionic bond gains
covalent character. Some lattices are not regular
and the crystal structure can differ.

Which kind of bonds will be
the most ionic? Why?

Between large positive ions and small negative
ions e.g. CsF

Define the terms
spontaneous and feasible

If a reaction is spontaneous and feasible, it will
take place of its own accord

does not take
account of rate of reaction.

Is a reaction with a positive
or negative enthalpy change
more likely to be
spontaneous?

Negative - exothermic

Define entropy

Randomness/disorder of a system.
Higher value for entropy = more
disordered

What units is entropy
measured in?

JK -1 mol -1

What is the second law of
thermodynamics?

Entropy (of an isolated system) always
increases, as it is overwhelmingly more likely for
molecules to be disordered than ordered

Is a reaction with positive or
negative entropy change
more likely to be
spontaneous?

Is a reaction with positive or negative entropy
change more likely to be spontaneous?
Positive - reactions always try and increase the
amount of disorder

Compare the general
entropy values for solids,
liquids and gases

Solids < liquids < gases

How would you calculate
the entropy change for a
reaction?

Entropy change = sum of products' entropy - sum
of reactants' entropy
ΔS = ΣS(products) - ΣS(reactants)

Define Gibbs free energy
using an equation

ΔG = ΔH - TΔS (G = Gibbs free energy, H =
enthalpy change, S = entropy change, T =
temperature)

What does the value for
Gibbs free energy for a
reaction show?

If G < 0, reaction is feasible. If G = 0, reaction is JUST feasible. If G > 0, reaction is not feasible.

What is the significance of
the temperature at which G
= 0?

This is the temperature (in Kelvin) at which the
reaction becomes feasible.

How would you calculate
the temperature at which a
reaction becomes feasible?

Rearrange to T = (ΔH)/(ΔS) since G = 0

What are the limitations of
using G as an indicator of
whether a reaction will
occur?

Gibbs free energy only indicates if a reaction is feasible. It
does not take into account the rate of reaction (the kinetics of
the reaction). In reality, many reactions that are feasible at a
certain temperature have a rate of reaction that is so slow
that effectively no reaction is occurring.

If the reaction is exothermic
and entropy increases, what
is the value of G and what
does this mean?

G always negative, so reaction is always feasible

• product favoured

If the reaction is
endothermic and entropy
decreases, what is the value
of G and what does this
mean?

G always positive, so reaction is never feasible -
reactant favoured

If the reaction is exothermic
and entropy decreases,
what is the value of G and
what does this mean?

Temperature dependent

If the reaction is
endothermic and entropy
increases, what is the value
of G and what does this
mean?

Temperature dependent

Why is entropy zero at 0K?

No disorder - molecules/atoms are not moving or
vibrating and cannot be arranged in any other
way. Maximum possible state of order

What are the two key things
to look out for to decide if
entropy
increases/decreases/stays
relatively constant?

Number of moles - more moles made → increase in entropy
Going from solid → liquid/gas or liquid → gas

How is it possible for the
temperature of a substance
undergoing an endothermic
reaction to stay constant?

The heat that is given out escapes to the
surroundings