Thermodynamics Flashcards
What does Hess’ 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
0 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- moles
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 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 completely 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 an example equation for:
Standard enthalpy of formation
Standard enthalpy of combustion
Standard enthalpy of atomisation
- Mg (s) + 1/2 O2 (g) —> MgO (s)
- CH4 (g) + 2O2 (g) —> CO2 (g) + 2H2O (g)
- 1/2 I2 (g) —> I (g)
Write an example equation for:
First ionisation energy
Second ionisation energy
First electron affinity
- Li (g) —> Li+ (g) + e-
- Mg+ (g) —> Mg2+ (g) + e-
- Cl (g) + e- —> Cl- (g)
Write an example equation for
Second electron affinity
Lattice enthalpy of formation
Lattice enthalpy of dissociation
- O- (g) + e- —> O*2- (g)
- Na+ (g) + Cl- (g) —> NaCl (s)
- NaCl (s) —> Na+ (g) + Cl- (g)
Write an example equation for
Enthalpy of hydration
Enthalpy of solution
Mean bond dissociation enthalpy
- Na+ (g) —> Na+ (aq)
- NaCl (s) —> Na+ (aq) + Cl- (aq)
- Br2 (g) —> 2Br (g)
What is a Born-Haber cycle
Thermochemical cycle showing all the enthalpy changes involved in the formation of ionic compound. Start with elements in their standard states (enthalpy of 0)
What factors affect the lattice enthalpy of an ionic compound
Size of ions
Charge of 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 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 size 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 happens when a solid is dissolved in terms of interactions of the ions with water ions
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 symmetrical 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 gain 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?
J/K/mol or JK^-1mol^-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 a 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
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 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 doesnt take into account the rate of reaction (Kinetics of the reaction)
In reality many reactions that are feasible at certain temperature have a rate of reaction that is so slow that effectively no reaction is occuring
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 the entropy decreases what is the value of G and what does this mean
G is 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 value 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 two things to look out for to decide if entropy increase/decreases/stays relatively constant
Number of moles - more moles made —> increase in entropy
Going from to 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 given out escapes into the surroundings
