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Chemistry > Unit 2.3 Enthalpy > Flashcards

Unit 2.3 Enthalpy Flashcards

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0
Q

What is enthalpy?

A

The heat content that is stored in a chemical system

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1
Q

What is chemical energy?

A

A special form of potential energy that lies within chemical bonds

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2
Q

How can you measure enthalpy?

A

By measuring the heat exchange with the surroundings

  • Heat loss in a chemical system = heat gain by surroundings
  • Heat gain in a chemical system = heat loss by surroundings
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3
Q

What is the law of conservation of energy?

A

Energy is never lost or created, it is simply transferred from 1 place to another

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4
Q

What is the chemical system?

A

The reactants and products

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5
Q

How can you calculate enthalpy change?

A

Enthalpy of products - enthalpy of reactants

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6
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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7
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 taken in from the surroundings
  • ΔH is positive
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8
Q

Definition of activation energy

A

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

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9
Q

Why is activation energy important for exothermic reactions?

A

Without it, they would take place spontaneously

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10
Q

Definition of ‘the standard enthalpy change of reaction’

A

The enthalpy change that accompanies a reaction in the molar quantities expressed in a chemical equation under standard conditions, all reactants and products being in their standard states.

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11
Q

What are standard conditions?

A
  • A pressure of 100 kPa (1 atmosphere)
  • A stated temperature, usually 298K (25°C)
  • A conc of 1 mol dm ^-3 (for reactions with aqueous solutions)
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12
Q

Definition of standard state

A

The physical state of a substance under standard conditions

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13
Q

Definition of ‘the standard enthalpy change of combustion’

A

The enthalpy change that takes place when 1 mole of a substance reacts completely with oxygen under standard conditions, all reactants and products being in their standard states

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14
Q

Definition of ‘the standard enthalpy change of formation’

A

The enthalpy change that takes place when 1 mole of a compound is formed from its constituent elements in their standard states under standard conditions

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15
Q

What is the enthalpy change of formation of an element?

A

0 kJ mol^-1

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16
Q

How can you calculate heat exchange, Q?

A

m - the mass of the surroundings involved in the heat exchange
c - the specific heat capacity of the surroundings
ΔT - the temperature change of the surroundings

Q = mcΔT

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17
Q

Definition of standard heat capacity

A

The energy required to raise the temperature of 1g of a substance by 1°C

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18
Q

When using an equation to represent the enthalpy change of combustion, what must you be careful not to do?

A

DO NOT put a balancing number in front of the substance being burnt
(If you do, more than 1 mol would have been combusted)

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19
Q

How can you determine the enthalpy change of combustion by carrying out an experiment?

A
  • Burn a known mass of a substance in air
  • Heat a known mass of water
  • Measure the temp change in the water
  • Use Q = mcΔT to calculate the heat gained by the water
  • The heat loss from chemicals is negative Q
  • You can then use moles = mass/molar mass to work out the moles of the substance
  • Divide Q by the moles to find how much heat is lost by 1 mol of the substance
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20
Q

Why might the experimental value of the standard enthalpy change of combustion be different to the data book value?

A
  • There may have been incomplete combustion

- There may have been heat loss to the surroundings

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21
Q

Definition of average bond enthalpy

A

The average enthalpy change that takes place when breaking by homolytic fission 1 mol of a given type of bond in the molecules of a gaseous species

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22
Q

Is bond breaking an exothermic or an endothermic process?

A

Endothermic - energy is needed to break bonds in the reactants

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23
Q

Is bond making an exothermic or an endothermic process?

A

Exothermic - energy is released as new bonds are formed in the products

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24
Q

How can you calculate enthalpy change using bond enthalpies?

A

ΔH = Σ(bond enthalpies of bonds broken) - Σ(bond enthalpies of bonds made)

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25
Q

In an exothermic reaction, are the bonds that are formed stronger or weaker than the bonds that are broken?

A

Stronger

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26
Q

Why is it not always possible to measure the enthalpy change of a reaction directly?

A
  • High activation energy
  • Slow reaction rate
  • More than 1 reaction taking place
27
Q

What does Hess’ law state?

A

If a reaction can take place by more than 1 route and the initial and final conditions are the same, the total enthalpy change is the same for each route

28
Q

If you want to work out the enthalpy change for a particular route:
Reactants —-(A)—-> product
And you are given a different route which involves an intermediate:
Reactants —-(B)—-> intermediate —-(C)—-> product
How can you work out A?

A

A = B + C (because of Hess’ law)

29
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’ law.

30
Q

How can you use Hess’ law to work out enthalpy change, using enthalpy change of formation values?

A

Set up an enthalpy cycle with the reactants, products and the elements
- Reactants —(A)—> products
- Elements —(B)—> reactants
- Elements —(C)–> products
So by following the arrows, you can find that C = A + B
Work out the values for B and C to find A
Don’t forget to include the number of moles, and that the standard enthalpy change of formation for an element is 0!

31
Q

How can you calculate the rate of reaction?

A

Change in concentration/time

Units: mol dm^-3 s^-1

33
Q

What does the collision theory state?

A

A chemical reaction can only take place when the reacting molecules collide
When 2 molecules collide, a reaction might take place if:
- The molecules have sufficient energy to overcome the activation energy of the reaction
- The molecules collide in the correct orientation

34
Q

What effect does concentration have on the reaction rate?

A

If the concentrations of the reactants are increased, the rate of reaction also increases:

  • Increased conc gives more molecules in the same volume
  • The molecules will be closer together and there is a greater chance of the molecules colliding
  • Collisions will be more frequent
35
Q

What effect does pressure have on the reaction rate?

A

If the pressure of a gas is increased, the rate of reaction also increases:

  • The molecules are pushed closer together
  • So the same number of molecules occupy a smaller volume
  • There is a greater chance of the molecules colliding
  • Collisions will be more frequent
36
Q

What is a catalyst?

A

A substance that increases the rate of a chemical reaction without being used up in the process

37
Q

How does a catalyst work?

A
  • It lowers the activation energy of the reaction by providing an alternative route with lower energy for the reaction to follow
  • The catalyst may react to form an intermediate, but it is later regenerated so that it does not undergo any permanent change
38
Q

Why are catalysts useful for many industrial processes?

A

They reduce costs:
- They speed up the process by lowering the activation energy
- So less energy is needed for the molecules to react
- Less energy reduces costs
Less energy also means that less fossil fuel is burnt:
- less CO2 will be released into the atmosphere
They can make a process more effective:
- By improving the percentage yield of an industrial preparation

39
Q

Definition of heterogeneous catalysis

A

Catalysis of a reaction in which the catalyst has a different physical state from the reactants

40
Q

Definition of homogeneous catalysis

A

Catalysis of a reaction in which the catalyst and reactants are in the same physical state

41
Q

What is biocatalysis?

A

Any process in which the catalyst is an enzyme

42
Q

What are enzymes?

A
  • Large protein molecules that act as biological catalysts

- They are able to catalyse the reactions of large quantities of biological molecules in very short periods of time

43
Q

What conditions do enzymes operate under?

A
  • Low temps
  • Atmospheric pressure
  • At an optimum pH value
44
Q

Why are enzymes useful in industry?

A
  • Lower temps and pressure can be used than with conventional inorganic catalysts, saving energy and costs
  • They often allow a reaction to take place which forms pure products, with no side reactions. This reduces the need for complex separation techniques, so reduces costs
  • Conventional catalysts are often poisonous and difficult to dispose of, but enzymes are biodegradable
45
Q

What happens in the Haber process?

A

Ammonia is made by reacting together nitrogen and hydrogen:

  • The triple bond in nitrogen has to be broken
  • This requires a large input of energy, contributing to a high activation energy
  • Iron is used to catalyse this reaction, weakening the triple bond in nitrogen, and lowering the activation energy
46
Q

What is the Boltzmann distribution?

A

The distribution of energies of molecules at a particular temperature, often shown as a graph
(Some molecules move fast and have high energy, but some move slowly and have low energy. Most have an average energy)

47
Q

Name 3 features of the Boltzmann distribution.

A
  • Area under the curve = total number of molecules in the sample
  • There are no molecules in the system with 0 energy
  • There is no maximum energy for a molecule
  • Only the molecules with an energy greater than the activation energy are able to react
48
Q

How does the Boltzmann distribution change when the temperature is increased?

A

It flattens and shifts to the right:

  • The peak moves to a higher energy with a lower height
  • A greater proportion of molecules exceeds the activation energy, so the rate of reaction increases
  • The number of molecules in the system doesn’t change, so the area under the curve remains the same
49
Q

How does the Boltzmann distribution change when a catalyst is added?

A
  • The curve doesn’t change at all
  • But the activation energy is lower, so a greater proportion of molecules exceeds the lower activation energy
  • Reaction rate increases
50
Q

What is dynamic equilibrium?

A

The equilibrium that exists in a closed system when:

  • The rate of the forward reaction is equal to the rate of the reverse reaction
  • The concentrations of the reactants + products remain the same
51
Q

What is the position of equilibrium?

A

The extent of a reaction at equilibrium

52
Q

What factors affect the position of equilibrium?

A
  • Concentrations of the reactants and products
  • Temperature
  • Pressure (in reactions involving gases)
53
Q

What does le Chatelier’s principle state?

A

When a system in dynamic equilibrium is subjected to a change, the position of equilibrium will shift to minimise the change

54
Q

What effect does increasing the concentration of a reactant have on the position of equilibrium?

A

It causes the position of equilibrium to move in the direction that decreases this increased reactant concentration:

  • The system opposes the change by decreasing the conc of the reactant by removing it
  • So the position of equilibrium moves to the right-hand side, forming more products
55
Q

What effect does increasing the concentration of a product have on the position of equilibrium?

A

It causes the position of equilibrium to move in the direction that decreases this increased product concentration:

  • The system opposes the change by decreasing the concentration of the product by removing it
  • The position of equilibrium will move to the left-hand side, forming more reactants
56
Q

How can you tell which side of a reaction has the highest pressure?

A

The side with the greater moles of gas has the highest pressure

57
Q

What effect does increasing the total pressure of the system have on the position of equilibrium?

A

It causes the position of equilibrium to move to the side with fewer gas molecules, because this will decrease the pressure

58
Q

What effect does decreasing the total pressure of the system have on the position of equilibrium?

A

It causes the position of equilibrium to move to the side with the greater number of gas molecules, because this will kncrease the pressure

59
Q

What effect does increasing the temperature of the system have on the position of equilibrium?

A

It causes the position of equilibrium to move in the direction that decreases the temperature (the endothermic direction)

60
Q

What effect does decreasing the temperature of the system have on the position of equilibrium?

A

It causes the position of equilibrium to move in the direction that increases the temperature (the exothermic direction)

61
Q

What effect does a catalyst have on the position of equilibrium?

A

It doesn’t affect the position of equilibrium:

  • It speeds up the rate of the forward and reverse reactions equally
  • It increases the rate at which equilibrium is established
62
Q

What conditions favour production of ammonia?

A

Ammonia is produced by the forward reaction:

  • The forward reaction produces fewer gas molecules, favoured by using high pressure
  • The forward reaction is exothermic, favoured by using low temp
63
Q

What are the drawbacks to using the ‘favourable’ conditions for ammonia production?

A
  • Although a low temp should produce a high equilibrium yield, the reaction would take place at a very low rate
  • A lot of energy is required to maintain a high pressure, so it would be very expensive
  • A high pressure also has a lot of safety implications
64
Q

What conditions are required for the Haber process?

A
  • Temperature of 400-500°C
  • Pressure of 200 atmospheres
  • An iron catalyst
66
Q

What factors affect the rate of a chemical reaction?

A
  • Temperature
  • Pressure (when the reactants are gases)
  • Concentration
  • Surface area
  • Adding a catalyst

Chemistry (4 decks)

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