5.2 ENERGY

Question 1

Standard lattice enthalpy Δ(le)H°

Answer 1

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

Question 2

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

Answer 2

• Stronger ionic bonds, higher melting/boiling point

Question 3

Standard enthalpy of formation Δ(f)H°

Answer 3

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

Question 4

Standard enthalpy of atomisation Δ(at)H°

Answer 4

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

Question 5

First ionisation energy Δ(ie1)H°

Answer 5

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

Question 6

Second ionisation energy Δ(ie2)H°

Answer 6

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

Question 7

First electron affinity Δ(ea1)H°

Answer 7

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

Question 8

Second electron affinity Δ(ea2)H°

Answer 8

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

Question 9

Standard enthalpy of solution Δ(sol)H°

Answer 9

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

Question 10

Standard enthalpy of hydration Δ(hyd)H°

Answer 10

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

Question 11

Born-Haber cycle for NaCl

Answer 11

Question 12

Born-Haber cycle for MgCl2

Answer 12

Question 13

Born-Haber cycle for Na2O

Answer 13

Question 14

Born-Haber cycle for MgO

Answer 14

Question 15

Enthalpy cycle for KCl

Answer 15

Question 16

State the two factors effect the exothermic value of lattice enthalpy

Answer 16

• Size of ions in the ionic lattice
• The charges of those ions

Question 17

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

Answer 17

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

Question 18

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

Answer 18

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

Question 19

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

Answer 19

• Size of ions in the ionic lattice
• The charges of those ions

Question 20

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

Answer 20

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

Question 21

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

Answer 21

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

Question 22

Explain why exothermic reactions are spontaneous.

Answer 22

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

Question 23

Define entropy, S

Answer 23

• the measure disorder in a system

Question 24

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

Answer 24

• 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

Question 25

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

Answer 25

- If there is an increase in number of gas molecules, there is an increase in entropy (more disorder)

Question 26

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

Answer 26

- If there is a decrease in number of gas molecules, theres a decrease in entropy (less disorder)

Question 27

Standard entropy change of reaction **ΔS°**

Answer 27

- The entropy change that occurs when a **reaction takes place** under standard conditons

Question 28

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

Answer 28

- there is always tendancy towards higher entropy

Question 29

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

Answer 29

- Changes in the reaction have caused the system to become **more disordered**

Question 30

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

Answer 30

- Changes in the reaction have caused the system to become **more ordered**

Question 31

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

Answer 31

ΔS° = ΣS°(products) - ΣS°(reactants) - Standard entropy change of a reaction = sum of standard entropies of products - sum of standard entropies of reactants

Question 32

Units of standard entropy change **ΔS°**

Answer 32

J K-1 mol-1

Question 33

Free energy equation

Answer 33

ΔG = ΔH - TΔS

Question 34

What is the free energy equation used for?

Answer 34

- To determine if a reaction is feasible or not

Question 35

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

Answer 35

- Enthalpy change (ΔH) (kJ mol-1) - Temperature (T) (K) - Entropy change (ΔS) (J K-1 mol-1)

Question 36

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

Answer 36

- If **ΔG** is **0** or **negative** value the reaction is **feasible**

Question 37

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

Answer 37

- 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

Question 38

State how **minimum temperature** for **ΔG feasibility** is calculated

Answer 38

Minimum temperature = **ΔH/ΔS**

Question 39

State two reasons a feasible reaction may not occur.

Answer 39

- Activation energy may be too high - Rate of reaction may be too slow

Question 40

Define "oxidising agent"

Answer 40

- a species that causes another species to be oxidised (electron loss) by being reduced itself (gaining those electrons)

Question 41

Define "reducing agent"

Answer 41

- a species that causes another species to be reduced (electron gain) by being oxidised itself (losing those electrons)

Question 42

State the **4** rules of **constructing half-equations**

Answer 42

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-**

Question 43

State the 2 rules of combining half-equations

Answer 43

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

Question 44

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

Answer 44

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+

Question 45

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

Answer 45

- **Acidified potassium permanganate** (KMnO4) - **Purple** with **7+** oxidation state (**MnO4-**) to **colourless** with **2+** oxidation state (**Mn2+**)

Question 46

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

Answer 46

Question 47

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

Answer 47

- Can be used to find out unknown concentrations of oxidising agents (e.g KIO3)

Question 48

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

Answer 48

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

Question 49

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

Answer 49

- Blue/black in solutions with iodine (I2) - Colourless in solutions with iodine ions (I-)

Question 50

State the **two** types of **half cells**

Answer 50

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**

Question 51

Define **standard electrode potential, E°**

Answer 51

- the **electrode potential** of a **half cell** relative to **hydrogen**

Question 52

State **standard conditions**

Answer 52

- **298K** - **100kPa** - **1M**

Question 53

State the role of the **salt bridge**

Answer 53

- Allows **ions** to **flow** through

Question 54

State what form **electrochemical equations** for **E°** are written in

Answer 54

- the **reduced** form (as a **reduction** reaction)

Question 55

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

Answer 55

- **NO PR**oblem - the **more negative E° value** is **oxidised** ( so goes the **other way**) - the **more positive E° value** is **reduced** ( so stays the **same**)

Question 56

Draw the setup of a **standard hydrogen electrode** (**SHE**)

Answer 56

Question 57

State the formula for **standard cell potential, E°cell**

Answer 57

Question 58

State how the **feasibility** of an **electrochemical cell** is determined

Answer 58

- If the **standard cell potential, E°cell** is a **positive value** the reaction is **feasible**

Question 59

State and explain the **two** limitations of calculating **electrochemical cell feasibility**

Answer 59

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**

Question 60

State the **three** types of **cells**

Answer 60

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

Question 61

State the **reactants** required for **fuel cells**

Answer 61

- **Hydrogen gas** from a **fuel** - **Oxygen**

Question 62

State the **general** reaction of **fuel cells**

Answer 62

Question 63

Explain why **fuel cells** dont need to be **recharged**

Answer 63

- Because they can operate continuously as long as **fuel** and **oxygen** flow into the cell