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Chemistry A2 > A2 Chemistry Term 1 > Flashcards

A2 Chemistry Term 1 Flashcards

1
Q

Define electron affinity and describe the type of enthalpy change that occurs

A

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
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2
Q

Define lattice energy

A

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
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3
Q

Define enthalpy change of atomisation

A

Enthalpy change when one mole of gaseous atoms is formed from its elements under standard conditions

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4
Q

Describe the factors affecting the value of lattice energy

A

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
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5
Q

Define ion polarisation and describe the factors affecting it

A
  • 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
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6
Q

Outline the Born-Haber cycle for the formation of lattices

A

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

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7
Q

Describe how an energy level diagram is drawn

A
  • arrows going upwards represent an increase in energy: endothermic
  • arrows goings downwards represent decrease in energy: exothermic
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8
Q

Define enthalpy change of solution and hydration

A

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

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9
Q

Describe the the calculation for the enthalpy change in solution

A

gaseous ions → ionic solid → ions in aqueous solution

H(latt) + H(solution) = H(hydration of cation and anion)

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10
Q

Define entropy

A
  • 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)
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11
Q

Describe when entropy changes occur

A

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

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12
Q

Describe how to predict whether entropy change is positive or negative

A

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.

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13
Q

Describe how total entropy change and entropy change of surroundings is calculated

A

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

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14
Q

Define Gibb’s free energy

A

ΔG = ΔH(reaction) - TΔS(system)

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15
Q

Describe what Gibb’s free energy shows

A
  • The value of ΔG must be negative for a reaction to be spontaneous
  • Not spontaneous if ΔG is positive
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16
Q

Describe the effect of ΔH and TΔS on the spontaneity of a reaction

A

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.
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17
Q

Define Kw

A
  • 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
18
Q

State the formula for pH

A

pH = -log[H+]

19
Q

Define Ka and pKa

A

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
20
Q

State the assumptions made when calculating the pH of weak acids

A
  • 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
21
Q

Describe how indicators work

A
  • 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
22
Q

Describe the changes in pH graphs that occur in acid-base titrations

A

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

23
Q

Define equivalence point

A

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.

24
Q

Define a buffer solution and describe how it works

A
  • 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.
25
Q

State the formulas for calculating the pH of a buffer solution

A

[H+] = Ka x [acid]/[salt]

pH = pKa + log [salt]/[acid]

26
Q

Describe the use of buffer solutions in the blood

A

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.
27
Q

Define the solubility product

A

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

28
Q

Describe how precipitation can be predicted using the solubility product

A

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.

29
Q

Define common ion effect

A

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.

30
Q

Define the partition coefficient

A

The equilibrium constant that relates the concentration of a solute partitioned between two immiscible solvents at a particular temperature

31
Q

Define the rate of reaction

A

Change in concentration / time taken for the change

  • decrease in conc. of reactant
  • increase in conc. of product
32
Q

Define rate constant

A

The proportionality constant of the rate to the concentration of substance. It is used to form the rate equation.

33
Q

Define the order of reaction

A

the power to which the concentration of a reactant is raised in the rate equation

34
Q

Describe the concentration vs rate graphs formed by different orders

A 
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
35
Q

Define half-life and describe its values for the different orders

A

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

36
Q

State the formula for half-life

A

t1/2 = 0.693/k

37
Q

Define heterogenous and homogenous catalysts

A

Heterogenous: Catalyst is in a different phase to reaction mixture

Homogenous: Catalyst is in same phase as reaction mixture

38
Q

Outline the iodine-peroxidisulfate reaction

A

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.

  1. 2Fe 3+ + 2I - → 2Fe 2+ + I2
  2. 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.

39
Q

Outline the formation of sulfur trioxide

A

SO2 + 1/2O2 → SO3

  1. SO2 + NO2 → SO3 + NO
  2. NO + 1/2O2 → NO2
    - The NO2 is in gaseous form meaning it is a homogenous catalyst
40
Q

Outline the catalysis of the Haber process

A
  1. Diffusion: Nitrogen gas and Hydrogen gas diffuse to surface of iron.
  2. 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.
  3. Reaction: Ammonia formed
  4. Desorption: Bonds between ammonia and iron weaken and break
  5. Diffusion away.
41
Q

Describe the catalysis that occurs in catalytic converters

A
  • 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

Chemistry A2 (3 decks)

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