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OCR Chemistry > 5.2 ENERGY > Flashcards

5.2 ENERGY Flashcards

1
Q

Standard lattice enthalpy Δ(le)H°

A

The enthalpy change that occurs when one mole of an ionic lattice is formed from its gaseous ions under standard conditions

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2
Q

Explain what a more exothermic (negative) lattice enthalpy Δ(LE)H° says about the ionic bonds and MP/BP

A
  • Stronger ionic bonds, higher melting/boiling point
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3
Q

Standard enthalpy of formation Δ(f)H°

A

The enthalpy change that occurs when one mole of a compound is formed from its elements under standard conditions

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4
Q

Standard enthalpy of atomisation Δ(at)H°

A

The enthalpy change that occurs when one mole of gaseous atoms is formed from its element under standard conditions

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5
Q

First ionisation energy Δ(ie1)H°

A

The enthalpy change that occurs when one mole of gaseous 1+ ions are formed from one mole of gaseous atoms under standard conditions

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6
Q

Second ionisation energy Δ(ie2)H°

A

The enthalpy change that occurs when one mole of gaseous 2+ ions is formed from one mole of gaseous 1+ ions under standard conditions

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7
Q

First electron affinity Δ(ea1)H°

A

The enthalpy change that occurs when one mole of gaseous 1- ions are formed from one mole of gaseous atoms under standard conditions

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8
Q

Second electron affinity Δ(ea2)H°

A

The enthalpy change that occurs when one mole of gaseous 2- ions are formed from one mole of gaseous 1- ions under standard conditions

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9
Q

Standard enthalpy of solution Δ(sol)H°

A

The enthalpy change that occurs when one mole of an ionic lattice is completely dissolved in water under standard conditions

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10
Q

Standard enthalpy of hydration Δ(hyd)H°

A

The enthalpy change that occurs when one mole of gaseous ions dissolve to form one mole of aqueous ions under standard conditions

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11
Q

Born-Haber cycle for NaCl

A
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12
Q

Born-Haber cycle for MgCl2

A
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13
Q

Born-Haber cycle for Na2O

A
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14
Q

Born-Haber cycle for MgO

A
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15
Q

Enthalpy cycle for KCl

A
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16
Q

State the two factors effect the exothermic value of lattice enthalpy

A
  • Size of ions in the ionic lattice
  • The charges of those ions
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17
Q

Explain how the size of ions in the ionic lattice effect the value of lattice enthalpy

A
  • Smaller ions get closer together so will have greater attraction so give more exothermic (negative) lattice enthalpy value
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18
Q

Explain how the charge of ions in the ionic lattice effect the value of lattice enthalpy

A
  • Ions with higher charge cause greater electrostatic attraction so give more exothermic (negative) lattice enthalpy value
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19
Q

State the two factors effect the exothermic value of the standard enthalpy change of hydration

A
  • Size of ions in the ionic lattice
  • The charges of those ions
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20
Q

Explain how the size of ions effect the value of the standard enthalpy change of hydration

A
  • Smaller ions can get closer to water molecules so greater attraction to water molecules so give more exothermic (negative) enthalpy change of hydration
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21
Q

Explain how the charge on ions effect the value of the standard enthalpy change of hydration

A
  • Higher charge on ion means greater attraction to water molecules so more exothermic (negative) enthalpy change of hydration
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22
Q

Explain why exothermic reactions are spontaneous.

A
  • Exothermic reactions result in products that are more energetically stable than reactants which allows it to proceed without external influence
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23
Q

Define entropy, S

A
  • the measure disorder in a system
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24
Q

Explain the difference in entropy, S of solids, liquids and gases

A
  • solids < liquids < gases
  • Gases have the highest entropy because their particles have most kinetic energy so have highest disorder
  • Solids have the lowest entropy because their particles have lowest kinetic energyso have the lowest disorder
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25
Q

What happens to entropy, S when there is an increase in gas molecules in a reaction?

A
  • If there is an increase in number of gas molecules, there is an increase in entropy (more disorder)
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26
Q

What happens to entropy, S when there is a decrease in gas molecules in a reaction?

A
  • If there is a decrease in number of gas molecules, theres a decrease in entropy (less disorder)
27
Q

Standard entropy change of reaction ΔS°

A
  • The entropy change that occurs when a reaction takes place under standard conditons
28
Q

What does the second law of thermodynamics say about standard entropy change of reaction ΔS°?

A
  • there is always tendancy towards higher entropy
29
Q

What does a positive standard entropy change of reaction ΔS° mean?

A
  • Changes in the reaction have caused the system to become more disordered
30
Q

What does a negative standard entropy change of reaction ΔS° mean?

A
  • Changes in the reaction have caused the system to become more ordered
31
Q

Equation for standard entropy change of reaction ΔS° (given entropies of reactants & products)

A

ΔS° = ΣS°(products) - ΣS°(reactants)

  • Standard entropy change of a reaction = sum of standard entropies of products - sum of standard entropies of reactants
32
Q

Units of standard entropy change ΔS°

A

J K-1 mol-1

33
Q

Free energy equation

A

ΔG = ΔH - TΔS

34
Q

What is the free energy equation used for?

A
  • To determine if a reaction is feasible or not
35
Q

What 3 factors does the feasibility (ΔG) of a reaction depend on?

A
  • Enthalpy change (ΔH) (kJ mol-1)
  • Temperature (T) (K)
  • Entropy change (ΔS) (J K-1 mol-1)
36
Q

What value shows a reaction is feasible using free energy change,ΔG?

A
  • If ΔG is 0 or negative value the reaction is feasible
37
Q

Use Gibbs free energy equation to explain why some endothermic reactions can be spontaneous.

A
  • Increasing entropy is energetically favourable so even if a reaction is enthalpically unfavourable (endothermic) it can still be spontaneous if entropy change (ΔS) is positive enough and temperature (T) is high enough to be greater than enthalpy change (ΔH):

ΔG = ΔH - TΔS

38
Q

State how minimum temperature for ΔG feasibility is calculated

A

Minimum temperature = ΔH/ΔS

39
Q

State two reasons a feasible reaction may not occur.

A
  • Activation energy may be too high
  • Rate of reaction may be too slow
40
Q

Define “oxidising agent”

A
  • a species that causes another species to be oxidised (electron loss) by being reduced itself (gaining those electrons)
41
Q

Define “reducing agent”

A
  • a species that causes another species to be reduced (electron gain) by being oxidised itself (losing those electrons)
42
Q

State the 4 rules of constructing half-equations

A

1) Write down the species before and after the reaction
2) Balance any atoms, apart from hydrogens and oxygens
3) Balance any oxygens out with water molecules, balance any hydrogens out with H+ ions
5) Balance charges with electrons e-

43
Q

State the 2 rules of combining half-equations

A

1) Make sure electrons on both half equations are equal
2) Cancel the electrons out and combine the rest of each equation

44
Q

State the 4 rules for constructing balanced redox equations using oxidation states.

A

1) Write down the reaction and all oxidation numbers to determine what has been oxidised and reduced
2) Balance only the species oxidised and reduced
3) If oxidation and reduction numbers are not equal, equalise them using balancing numbers
4) Balance out any other atoms with H2O and H+

45
Q

State the oxidising agent and colour change used in Fe2+/MnO4- redox titration

A
  • Acidified potassium permanganate (KMnO4)
  • Purple with 7+ oxidation state (MnO4-) to colourless with 2+ oxidation state (Mn2+)
46
Q

State the overall redox equation in the Fe2+/MnO4- redox titration using (Fe2+ → Fe3+) and (MnO4- → Mn2+)

A
47
Q

Describe what the iodine-sodium thiosulfate redox titration can be used for.

A
  • Can be used to find out unknown concentrations of oxidising agents (e.g KIO3)
48
Q

Describe the steps of deducing unknown concentration of an oxidising agent using the iodine-sodium thiosulfate redox titration.

A

1) Mix known volume & unknown concentration of oxidising agent (e.g KIO3) with potassium iodide , KI, to oxidise iodide ions from KI into iodine atoms
2) Add solution mixture from step 1 into conical flask with starch to give blue/black colour
3) Add the sodium thiosulfate , Na2S2O3 , reducing agent into a burette and titre it dropwise until solution goes from blue/black to colourless
4) Use the volume of sodium thiosulfate , Na2S2O3 , used to work about the concentration of oxidising agent

49
Q

State the two colours of starch solution with different species of iodine.

A
  • Blue/black in solutions with iodine (I2)
  • Colourless in solutions with iodine ions (I-)
50
Q

State the two types of half cells

A

1) Metal electrode dipped in an aqueous solution of its ions
2) Platinum electrode dipped in an aqueous solution of ions of the same element in different oxidation states

51
Q

Define standard electrode potential, E°

A
  • the electrode potential of a half cell relative to hydrogen
52
Q

State standard conditions

A
  • 298K
  • 100kPa
  • 1M
53
Q

State the role of the salt bridge

A
  • Allows ions to flow through
54
Q

State what form electrochemical equations for E° are written in

A
  • the reduced form (as a reduction reaction)
55
Q

State how we determine which half cell is oxidised and reduced using E° values

A
  • NO PRoblem
  • the more negative E° value is oxidised ( so goes the other way)
  • the more positive E° value is reduced ( so stays the same)
56
Q

Draw the setup of a standard hydrogen electrode (SHE)

A
57
Q

State the formula for standard cell potential, E°cell

A
58
Q

State how the feasibility of an electrochemical cell is determined

A
  • If the standard cell potential, E°cell is a positive value the reaction is feasible
59
Q

State and explain the two limitations of calculating electrochemical cell feasibility

A

1) Reaction may still not be feasible as rate of reaction is too slow due to too high activation energy
2) Non-standard concentrations affect the standard cell potential, E°cell

60
Q

State the three types of cells

A

1) Fuel cell
2) Rechargeable cell
3) Non-rechargeable cell

61
Q

State the reactants required for fuel cells

A
  • Hydrogen gas from a fuel
  • Oxygen
62
Q

State the general reaction of fuel cells

A
63
Q

Explain why fuel cells dont need to be recharged

A
  • Because they can operate continuously as long as fuel and oxygen flow into the cell

OCR Chemistry (13 decks)

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