Energetics II Flashcards
Define lattice enthalpy of formation
The enthalpy change when 1 mole of a solid ionic compound is formed from its gaseous ions under standard conditions
define entropy
- measure of disorder
- number of ways that energy can be shared between particles
Define enthalpy change of atomisation
the enthalpy change when 1 mole of gaseous ions is made from an element in its standard state.
Define 1st electron affinity
the enthalpy change when 1 mole of gaseous 1- ions are formed from 1 mole of gaseous atoms.
Name 2 factors that affect strength of ionic bonding
- size of charge
- ionic radius
Explain how Size of charge effects strength of ionic bond and melting and boiling point (2)
- larger charge = stronger electrostatic attraction between ions
- therefore, more energy needed to overcome these forces, leading to a higher melting and boiling point.
Describe and explain how ionic radius affects the strength of ionic bonds (2)
- smaller ion = stronger electrostatic attraction between ions
- can also pack closer together, more energy required to overcome the stronger forces = higher boiling and melting points.
Describe the structure of a born haber cycle from bottom to top. (enthalpies) and whether the steps are exothermic or endothermic. (5)
- original ionic compound (formation of which is exothermic)
- enthalpy of atomisation (for all elements if needed) (endothermic)
- ionisation energy of Cation (endothermic)
- 1st electron affinity of Anion (endothermic)
- (going right now) enthalpy of lattice (exothermic) formation.
what are the assumptions that theoretical values of lattice energy make? (2)
- assume a perfect ionic model
- 100%, and attractions are purely electrostatic.
Why might theoretical and actual (born haber) values be different? (3)
- different if the compound shows covalent character - ie the cation distorts the anion slightly (kind of like they are sharing)
- The higher the covalent character, the higher the larger the difference would be in the theoretical and Born Haber value.
- the Born Haber value for a compound would be larger than the theoretical when the compound has some covalent character.
Define enthalpy of solution.
- the enthalpy change when 1 mole of an ionic solid dissolves in a large enough amount of water to ensure they don’t interact.
Define enthalpy of lattice formation
- the enthalpy change when 1 mole of an ionic crystal is formed from its constituent gaseous ions.
define enthalpy of hydration
- when one mole of gaseous ions become hydrated such that further dilution causes no heat change.
How do we calculate enthalpy change of solution?
Enthalpy change of solution = -Enthalpy of lattice formation + sum of enthalpy of hydration.
What can be concluded about a substance from its enthalpy change of solution?
- substance more likely soluble if exothermic
- substance more likely to be insoluble if endothermic (as lattice energy much larger than hydration enthalpy.)
what happens to the entropy when a solid dissolves into ions?
- entropy increases as there is more disorder, as solid changes to solution and number of particles increases.
- this can make the total ΔSsystem can make the ΔStotal positive even if ΔH solution is endothermic, esp at higher temps.
what is a system?
- your reactants and products
- does not change temp or pressure
- mass cannot be transferred to the surroundings
- energy can be transferred to the surroundings.
why do solids have lower entropies than liquids, which are lower than gasses?
- as a solid is heated up, and turns into liquid, its particles vibrate more, meaning that its entropy increases
- gasses are most disordered, so have the highest enthalpy.
What will an increase in disorder and entropy lead to
a positive entropy change ( ΔS’ system = +ve)
Give the equation to calculate ΔS system.
= ΣS products - ΣS reactants
What is the unit for entropy?
J/K/mol
Give the equation to calculate ΔS surroundings
what are some things to keep in mind when using the equation?
-ΔH reaction/T
- covert ΔH to J/mol (x1000)
- convert degrees to Kelvin (+273)
Give the equation to calculate total entropy change.
ΔS total = ΔS system + ΔS surrounding
What must be positive for a reaction to be spontaneous?
- ΔS total
What is Gibbs free energy?
- balance between entropy and enthalpy that determines the feasibility of a reaction.
What must be the condition for a reaction to be feasible?
the ΔG (gibbs free energy) must be LESS than 0.
Give the equation of Gibbs free energy.
ΔG = ΔH - TΔS system
(gibbs free energy = enthalpy change - (Temp x entropy change of system).
What will happen if the reaction involves a decrease in entropy? (ΔG is +ve)
- less likely that ΔG is negative, and less likely that the reaction will occur.
What happens if the reaction involves an increase in entropy?
- more likely the reaction tales place and more likely that ΔG is -ve.
When y=mx+c is applied to the Gibbs energy equation. give the equation that results.
c = ΔH
m=-ΔS
Which Kc value predicts that the reaction will not happen?
Kc < 10^-10
Which Kc value predicts that the reactants dominate in eq
Kc = 0.1
Which Kc value predicts that the reaction has equal amounts of reactants and products?
- Kc = 1
Which Kc value predicts that the products predominate in eq?
Kc = 10
Which Kc value predicts that the reaction goes to completion?
- Kc > 10^10
What can limit a reaction, causing it to not go to completion?
- if it has a high activation energy, kinetic factors limit the reaction.
What is the cycle that links enthalpy of hydration, solution, and Lattice enthalpy.
ΔsolutionH
LiCl —–> Li+ + Cl-
I ΔDissH I ΔhydH
V V
Li+ + Cl-
How to find the temp at what the reaction is feasible?
rearrange the Activation energy equation
- LnK = -ΔG/RT
Explain, in terms of entropy, why the combustion of prop-2-en-1-ol is always feasible in
the gaseous state.
ΔH is negative and ΔsSystem is positive
- so ΔG is always negative
