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A Level Chemistry > Module 3: Chapter 9 - Enthalpy > Flashcards

Module 3: Chapter 9 - Enthalpy Flashcards

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1
Q

What is enthalpy?

A

The heat content that is stored in a chemical system

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2
Q

What is enthalpy change?

A

The difference between the heat energy of the reactants and the heat energy of the products

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3
Q

What type of process is breaking bonds?

A

Energy is required to break forces of attraction, therefore bond breaking is an endothermic process

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4
Q

What type of process is making bonds?

A

Energy is released when new forces of attraction form, therefore making bonds is an exothermic process

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5
Q

Explain why combustion products are thermodynamically stable with respect to their reactants

A

As combustion is an exothermic process, the products are at a lower enthalpy than the reactants. Therefore the products are more thermodynamically stable with respect to the reactants as thermodynamic stability occurs when a system is in its lowest energy state.

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6
Q

Explain why combustion reactants are kinetically stable with respect to their products

A

As combustion is an exothermic process, the products are at a lower enthalpy than the reactants. Therefore the reactants are kinetically stable with respect to the products as kinetic stability occurs when a system is in its highest energy state

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7
Q

What is thermodynamic stability?

A

Thermodynamic stability is the stability of the lowest energy state of a system, it describes an equilibrium state of the system

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8
Q

What is kinetic stability?

A

Kinetic stability is the stability of the highest energy state of a system, it describes the reactivity of the system

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9
Q

What is the bond (dissociation) enthalpy?

A

The energy required to break one mole of the bond to give separated atoms in a gaseous molecule

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10
Q

What are surroundings?

A

Everything that is not the chemical system

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11
Q

What is a chemical system?

A

The atoms, molecules, opr ions that make up the chemicals involved in a reaction

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12
Q

What is an endothermic reaction?

A

A reaction in which the enthalpy of the products is greater than the enthalpy of the reactants, resulting in heat being taking in from the surroundings (ΔH is postive)

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13
Q

What is an exothermic reaction?

A

A reaction in which the enthalpy of the products is smaller than the enthalpy of the reactants, resulting in heat loss to the surroundings (ΔH is negative)

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14
Q

What is activation energy?

A

The minimum energy required to start a reaction by the breaking of bonds

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15
Q

What is the enthalpy change of a neutralisation reaction?

A

-57 kJ mol⁻¹

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16
Q

Why do all neutralisation reactions have an enthalpy change of -57 kJ mol⁻¹?

A

All neutralisation reactions should have an enthalpy change of -57 kJ mol⁻¹. This is because the overall ionic equation for all neutralisation reactions is “H⁺(aq) + OH⁻(aq) -> H₂O(l)”, all the other ions are spectator ions and therefore do not affect the enthalpy change.

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17
Q

How can you find the enthalpy change of a reaction experimentally?

A

By measuring the temperature change in a reaction. The energy given out can be used to heat a known mass of water, you can then calculate the enthalpy change

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18
Q

What is the equation for the heat energy given out (or absorbed) by a reaction?

A

q = m c ΔT
q = heat energy (J)
m = mass of substance heated (g)
c = specific heat capacity ( J g⁻¹ K⁻¹)
T = temperature rise (K)

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19
Q

What is the equation for enthalpy change per mole (obtained from experimental results)?

A

Enthalpy change per mole = q / number of moles reacting

q = m c ΔT

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20
Q

What is the specific heat capacity of water?

A

4.18 J g⁻¹ K⁻¹

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21
Q

100 cm³ of 0.200 mol dm⁻³ copper sulphate solution was put in a calorimeter and 2.00 g of magnesium powder added. The temperature of the solution rose by 25.1°C. Work out which reagent was in excess and then calculate the enthalpy change for the reaction. Assume that the density of the solution is 1.00 g cm⁻³, the specific heat capacity of the solution is 4.18 J g⁻¹ K⁻¹. Ignore the heat capacity of the metals

HARD QUESTION

A

-525 kJ mol⁻¹

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22
Q

25 cm³ of 1.00 mol dm⁻³ copper sulphate solution was put in a calorimeter and 6.0 g of zinc powder added. The temperature of the solution rose by 50.6°C. Work out which reagent was in excess and then calculate the enthalpy change for the reaction. Assume that the density of the solution is 1.00 g cm⁻³, the specific heat capacity of the solution is 4.18 J g⁻¹ K⁻¹. Ignore the heat capacity of the metals

HARD QUESTION

A

-212 kJ mol⁻¹

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23
Q

What is the standard enthalpy change of reaction (ΔᵣH°)

° = θ

A

The enthalpy change for a reaction with the quantities shown in the chemical equation

24
Q

What is standard pressure?

A

100 kPa

25
Q

What is standard temperature?

A

298 K

26
Q

What is standard concentration?

A

1 mol dm⁻³

27
Q

What is standard state?

A

The physical state of a substance under standard conditions

28
Q

What is the standard enthalpy change of formation (ΔfH°)?

° = θ

A

The enthalpy change when 1 mole of a substance is formed from its elemetns in their standard states under standard conditions of 298K and 100kPa

29
Q

What is the equation for the standard enthalpy change of formation of ammonia?

A

1/2 N₂(g) + 3/2 H₂(g) -> 1NH₃(g)

30
Q

What is the equation for the standard enthalpy change of formation of bromomethane?

A

C(s) + 3/2 H₂(g) + 1/2 Br₂(g) -> 1CH₃Br(l)

31
Q

What is the standard enthalpy change of combustion (ΔcH°)?

° = θ

A

The enthalpy change when one mole of a substance is completely burned in excess oxygen under standard conditions of 298K and 100kPa

32
Q

What is the equation for the standard enthalpy change of combustion of methane?

A

1CH₄(g) + 2O₂(g) -> 2H₂O(l) + CO₂(g)

33
Q

What is the equation for the standard enthalpy change of combustion of hexane?

A

1C₆H₁₄(l) + 19/2 O₂(g) -> 7H₂O(l) + 6CO₂(g)

34
Q

What is the equation for the standard enthalpy change of combustion of sodium?

A

1Na(s) + 1/4 O₂(g) -> 1/2 Na₂O(s)

35
Q

What is the standard enthalpy change of neutralisation (ΔₙₑᵤₜH°)

° = θ

A

The enthalpy change when one mole of H₂O(l) is formed from the reaction of an acid with a base under standard conditions of 298K and 100kPa

36
Q

What is the equation for all standard enthalpy changes of neutralisation?

A

H⁺(aq) + OH⁻(aq) -> H₂O(l)

37
Q

What is Hess’s law?

A

If a reaction can take place by 2 routes, and the starting and finishing conditions are the same, the total enthalpy change is the same for each route

38
Q

What is the standard enthalpy change of formation (ΔfH°)of an element?

A

zero (0)

39
Q

What does a hess’s law cycle involving enthalpy of formation data look like?

A

note directions of arrows

40
Q
A

-70 kJ mol⁻¹

41
Q

What does a hess’s law cycle involving enthalpy of combustion data look like?

A

Note direction of arrows

42
Q
A

-1075 kJ mol⁻¹

43
Q

HARD QUESTION

A

-235 kJ mol⁻¹

44
Q

What happens if you are given both data for enthalpy of formation and enthalpy of combustion?

A

Sometimes the enthalpy of combustion and enthalpy of formation can be the same value, for example “ΔcH of S = ΔfH of SO₂” and “ΔcH of C = ΔfH of CO₂”

45
Q

Why are you given the average bond enthalpy value and not just the bond enthalpy value?

A

the actual bond enthalpy can vary depending on the chemical environment of the bond, therefore and average value is used

46
Q

When would you be given the exact bond enthalpy value and not the average bond enthalpy value?

A

When only 1 molecule can contain that specific bond, for example H-H, Cl-Cl, H-Cl, O=O, etc

47
Q

What does a hess’s law cycle involving bond enthalpies look like?

A
48
Q

How do you determine the ΔT from experimental data (when a solution is heated and temperature is recorded periodically)

A

Draw a line of best fit for the cooling curve and extrapolate the line back to the time at which the reaction began

49
Q

In a hess’s law question involving bond enthalpies, what should you do if one of the reactants is in liquid or solid form?

A

You must include the enthalpy of vaporisation/and or enthalpy of fusion going from the reactants to the gaseous atoms

50
Q

HARD QUESTION

A

+314 kJ mol⁻¹

51
Q

Can enthalpy be directly measured?

A

No, enthalpy itself cannot be measured however enthalpy changes can be measured as they involve taking/giving energy out to the surrounding environment

52
Q

What are energy profile diagrams?

A

Diagram showing the relative enthalpies of the reactants and products, the enthalpy change, and the activation energy

53
Q

How does activation energy affect the rate of reaction?

A

If there is a low activation energy, the reaction will take palce rapidly as the energy required to break reactant bonds is readily available from the surroundings. However, very large activation energies may present such a large energy barrier the reaction may take place very slowly or not at all (without an external source of energy such as a bunsen burner)

54
Q

Wbhat is specific heat capacity?

A

The energy required to raise the temperature of 1g of a subtance by 1K

55
Q

Why may the experimental value achieved for the ΔcH be diffferent from the actual value?

A
  • Heat loss to the surroundings - energy may be transferred to the beaker and the surrounding air rather than to the solution
  • Not a 100% yield - incomplete combustion may occur, this can affect the energy released
  • Evaporation from the wick of the spirit burner - alcohols such as methanol are very volatile and may evaporate from the wick when the flame is extinguished. This can cause you to record a greater mass of alcohol reacted and therefore affect you value for the enthalpy cahnge
  • Non-standard conditions
  • Errors in measuring equipment
  • specific heat capacity of solution may differ
56
Q

What is an enthalpy cycle?

A

A diagram showing alternative routes between reactants and products which allows the indirect determination of an enthalpy change from other known enthalpy changes using hess’s law

A Level Chemistry (31 decks)

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