A2 Chemistry Term 1 Flashcards
Define electron affinity and describe the type of enthalpy change that occurs
Enthalpy change when one mole of electrons is added to 1 mole of gaseous atoms to form one mole of gaseous 1- anions under standard conditions.
- 1st is exothermic as bond forms between electron and atom
- 2nd is always endothermic due to repulsion between electrons
Define lattice energy
Enthalpy change when one mole of an ionic compound is formed from its gaseous ions under standard conditions
- Always exothermic as bonds are formed between ions
Define enthalpy change of atomisation
Enthalpy change when one mole of gaseous atoms is formed from its elements under standard conditions
Describe the factors affecting the value of lattice energy
Ion size:
- as size of ion increases lattice energy becomes less exothermic.
- if charge is same, increasing radius, charge density is lower.
- This results in weaker electrostatic forces of attraction in the ionic lattice
Charge on ion:
- lattice energy becomes more exothermic as ionic charge increases.
- if ions are same size, greater charge means higher charge density.
- Results in greater electrostatic forces of attraction
Define ion polarisation and describe the factors affecting it
- Positive ion in ionic lattice may attract electrons in anions towards it.
- Results in distortion of electron cloud of anion
Factors:
- Size of ions : small cation, large anion - increases polarisation
- Charge: Bigger charge results in greater polarisation
Outline the Born-Haber cycle for the formation of lattices
elements in standard states → ions in gaseous state → ionic compound
H(latt) + H1 = H(f)
H1 = atomisation of cation + ionisation energy of cation + atomisation of anion + electron affinity of anion
Describe how an energy level diagram is drawn
- arrows going upwards represent an increase in energy: endothermic
- arrows goings downwards represent decrease in energy: exothermic
Define enthalpy change of solution and hydration
Solution: Enthalpy change when one mole of ionic solid dissolves in sufficient water to form an infinitely dilute solution.
Hydration: Enthalpy change when one mole of gaseous ion dissolves in sufficient water to form infinitely dilute solution
Describe the the calculation for the enthalpy change in solution
gaseous ions → ionic solid → ions in aqueous solution
H(latt) + H(solution) = H(hydration of cation and anion)
Define entropy
- The measure of dispersal of energy at a specific temperature
- Measure of randomness or disorder of a system
- System becomes more energetically stable as entropy increases
- The unit of entropy is J / (K x mol)
Describe when entropy changes occur
Changes of state:
- gas has the most entropy
- solid has most ordered particles and so has lowest entropy
Temperature:
- Increasing temperature makes particles within a substance move more, hence the particles become less ordered
- Thus, entropy increases
Change in number of gaseous molecules:
- more gas molecules = more ways of arranging molecules = higher entropy
Describe how to predict whether entropy change is positive or negative
Exothermic reaction: energy released increases number of ways of arranging energy. Energy goes to rotation and translation of molecules in surroundings. Hence, increased entropy and increased probability of chemical change occurring spontaneously.
Endothermic reaction: energy absorbed from surroundings decreases ways of arranging energy. Likely to be a decrease in entropy and decreased probability of spontaneous chemical change.
Describe how total entropy change and entropy change of surroundings is calculated
ΔS(total) = ΔS(system) + ΔS(surroundings)
ΔS(system) = ΔS(products) - ΔS(reactants)
ΔS(surroundings) = -ΔH(reaction) / T
- ΔH(reaction) is standard enthalpy change of reaction
- T is temp in Kelvin
Define Gibb’s free energy
ΔG = ΔH(reaction) - TΔS(system)
Describe what Gibb’s free energy shows
- The value of ΔG must be negative for a reaction to be spontaneous
- Not spontaneous if ΔG is positive
Describe the effect of ΔH and TΔS on the spontaneity of a reaction
In exothermic reactions (ΔH is -ve):
- Spontaneous when ΔS is positive
- Spontaneous when ΔS is negative and temp is small. If temp is very high may not be spontaneous
In endothermic reactions (ΔH is +ve):
- Not spontaneous when ΔS is negative
- Not spontaneous when ΔS is positive and temp is small. If temp is high may be spontaneous.
Define Kw
- It is the ionic product of water. It is [H+][OH-]
- extent of ionisation of water is very low and so the concentration of water is considered constant.
- It is used to calculate the pH of strong bases
- It is 1 x 10^-14 at 298K
State the formula for pH
pH = -log[H+]
Define Ka and pKa
Ka: acid dissociation constant
- indicates extent of dissociation of acid
- high value means almost completely ionised
- low value indicates partially ionised
pKa = -log Ka
- used to easily compare strengths of acid as Ka values are very low
- lower the pKa the stronger the acid
State the assumptions made when calculating the pH of weak acids
- Concentration of H+ produced by ionisation of water molecules is negligible and so is ignored
- Ionisation of weak acid is so small that the concentration of undissociated HA molecules at equilibrium is approximately same as original acid
Describe how indicators work
- weak acid which is a different colour to its conjugate base
- adding acid or alkali changes position of equilibrium meaning amount of substance changes and colour changes
Describe the changes in pH graphs that occur in acid-base titrations
Strong acids + Strong bases:
- sharp fall with midpoint around pH 7
- lines are flat
Strong acids + Weak bases:
- sharp fall with midpoint slightly acidic e.g. 5
- basic line is curvy, acidic line is flat
Weak acids + Strong bases:
- sharp fall with midpoint slightly basic e.g. 9
- basic line is flat, acidic line is curvy
Weak acids + Weak bases:
- no sharp fall
Define equivalence point
point at which H+ ions in acid have exactly reacted with OH- ions in alkali. This is also known as the end point of the titration.
Define a buffer solution and describe how it works
- solution in which pH does not change significantly when small amounts of acid or alkalis are added
- contains large reserve supplies of weak acid and conjugate base
- adding acid shifts position of equilibrium left as H+ combines with conjugate base to reform acid. Large reserves mean concentration of base and acid is fairly constant.
- adding base shifts equilibrium right as OH- reacts with H+ reducing H+ concentration. Large reserve supplies ensure concentrations do not change significantly.
State the formulas for calculating the pH of a buffer solution
[H+] = Ka x [acid]/[salt]
pH = pKa + log [salt]/[acid]
Describe the use of buffer solutions in the blood
CO2 + H2O ⇌ H+ + HCO3-
- if H+ concentration increases then position of equilibrium shifts to the left as H+ reacts with HCO3- until equilibrium is restored
- if H+ concentration decreases then equilibrium shifts to the right as CO2 and H2O react to increase concentration and restore equilibrium.
Define the solubility product
Product of concentrations of each ion in a saturated solution of a sparingly soluble salt at 298K, raised to the power of their relative concentrations
Describe how precipitation can be predicted using the solubility product
If the product of the concentrations of the ions is higher than Ksp then a precipitate will form. If not, then no precipitate will form.
Define common ion effect
The reduction of the solubility of a dissolved salt achieved by adding a solution of a compound which has an ion in common with the dissolved salt.
Define the partition coefficient
The equilibrium constant that relates the concentration of a solute partitioned between two immiscible solvents at a particular temperature
Define the rate of reaction
Change in concentration / time taken for the change
- decrease in conc. of reactant
- increase in conc. of product
Define rate constant
The proportionality constant of the rate to the concentration of substance. It is used to form the rate equation.
Define the order of reaction
the power to which the concentration of a reactant is raised in the rate equation
Describe the concentration vs rate graphs formed by different orders
Zero order: - horizontal straight line - rate does not change with concentration First order: - Straight line going through origin - doubling concentration doubles rate Second order: - Upwardly curving line - doubling concentration quadruples rate
Define half-life and describe its values for the different orders
Time taken for concentration of reactants to fall to half of original value.
Zero order: half-lives decrease
First order: half-life is constant
Second order: half-lives increase
State the formula for half-life
t1/2 = 0.693/k
Define heterogenous and homogenous catalysts
Heterogenous: Catalyst is in a different phase to reaction mixture
Homogenous: Catalyst is in same phase as reaction mixture
Outline the iodine-peroxidisulfate reaction
S2O8 - + 2I - → 2SO4 2- + I2
As both ions have a negative charge, considerable energy is required for the reaction to occur. This means that uncatalysed, it is a slow reaction.
- 2Fe 3+ + 2I - → 2Fe 2+ + I2
- 2Fe 2+ + S2O8 2- → 2Fe 3+ + 2SO4 2-
Fe 3+ is the catalyst meaning it is a homogenous catalyst as it is in the aqueous form.
Outline the formation of sulfur trioxide
SO2 + 1/2O2 → SO3
- SO2 + NO2 → SO3 + NO
- NO + 1/2O2 → NO2
- The NO2 is in gaseous form meaning it is a homogenous catalyst
Outline the catalysis of the Haber process
- Diffusion: Nitrogen gas and Hydrogen gas diffuse to surface of iron.
- Adsorption: Reactant molecules are chemically adsorbed onto surface of iron. The bonds that form between the molecules and the iron are strong enough to weaken previous intramolecular bond but weak enough to allow them to leave the surface.
- Reaction: Ammonia formed
- Desorption: Bonds between ammonia and iron weaken and break
- Diffusion away.
Describe the catalysis that occurs in catalytic converters
- Converts harmful nitrogen oxides and CO from exhaust gases into harmless gases.
- Honeycomb structure contains small beads of platinum, palladium or rhodium which are heterogenous catalysts.
- Adsorption → Weakening → Reaction → Desorption
