Chemistry - 3.1.4: Energetics

What happens to energy when bonds are:
1. made?
2. broken?

1. Energy is given out to the surroundings (exothermic).
2. Energy is taken in from the surroundings (endothermic).

2 types of reactions + what happens to the temperature of the surroundings with each one

- Endothermic: a reaction where energy was taken in from the surroundings (surroundings temperature decreases)
- Exothermic: a reaction where energy was given out to the surroundings (surroundings temperature increases)

If a reaction is endothermic in one reaction, what is the reaction in the opposite reaction?

Exothermic (and vice versa)

Standard conditions for measuring enthalpy changes

Pressure: 100kPa (approx normal atmospheric pressure)
Temperature: 298K (approx normal room temperature, 25°C)

When is a reaction thought to be over?

When the temperature of the products have cooled back/warmed up to 298K

Axis of enthalpy level diagrams

x-axis: progress of reaction
y-axis: enthalpy

In an (1. exothermic / 2. endothermic) reaction, do the reactants or products have more energy?

1. Exothermic: the reactants have more energy
2. Endothermic: the products have more energy

Units for enthalpy change?

kJ mol-1

Standard enthalpy of formation

The enthalpy change when one mole of a compound is formed from its constituent elements under standard conditions with all substances in standard states.

Standard enthalpy of combustion

The enthalpy change when one mole of substance is completely burnt in oxygen under standard conditions with all substances in standard states.

How to measure enthalpy change (6)

1. Measure mass of the water
2. Measure the initial and final temperature of the water (calculate change in temperature)
3. Use specific heat capacity of water
4. Calculate energy change (q = m c ΔT)
5. Calculate moles of fuel (limiting reactant)
6. Calculate energy change per mol

Specific heat capacity

The amount of energy needed to increase the temperature of 1 g of a substance by 1 K

2 types of calorimeters - features and which one is better

1. Simple calorimeter: spirit burner with fuel lit below copper can
2. Flame calorimeter: spiral chimney made of copper, enclosed flame, fuel burns in pure oxygen (not air) - reduces energy loss

How to reduce heat loss during calorimetry

- Leave all apparatus and solutions to stand in the lab for some time so can reach same temperature (temperature of the lab)
- Use a polystyrene cup and lid, put it in glass beaker as they are thermal insulators (do NOT use for combustion - need heat energy from fuel to reach water in copper can)

How to plot a cooling curve to work out the temperature reached when a reactant was added during calorimetry

- Plot points of the temperature every minute
- When you mix the reactants, the temperature should rise/fall rapidly
- Join all points together, extrapolate backwards from points after biggest rise/fall to time where the reactants were mixed to find estimate of temperature at that time

Hess' Law

The enthalpy change of a reaction is independent of the route taken from reactants to products.
Therefore: Reactants -> Products = Reactants -> Elements -> Products / A = B + C

If there are the same moles of an element on either side of an equation, what is their effect on the enthalpy change?

No effect - they take no part in the reaction

If a substance is of a lower energy than another, what is it?

Energetically more stable

The enthalpies of all elements in their standards states are...

Taken as 0

Bond disassociation enthalpy

The enthalpy change required to break a covalent bond with all species in the gaseous state

Mean enthalpy change

The heat energy needed to break 1 mol of a covalent bond averaged over a range of different compounds

Pros and cons of mean bond enthalpies?

Pros: useful, quick, easy to use, can calculate ΔH for a compound that has never been made
Cons: will only give specific compounds approximate answers

In a combustion cycle, the arrows face...

downwards.

In a formation cycle, the arrows face...

upwards.

Symbol for enthalpy change of
1. Combustion
2. Formation

1. ΔcH
2.ΔfH

When measuring enthalpy change, why might the measured value be different to the calculated value?

Given bond energies are averages taken from a range of different compounds + experimental error may occur

Activation energy

The minimum amount of energy needed to start a reaction and break the reactants' bonds.

Why stir the reaction mixture regularly during calorimetry experiment (if doing enthalpy change of formation)?

Ensure that heat is evenly distributed in the reaction mixture.

Why might the enthalpy change of a reaction not be measured accurately? (3)

- Heat energy may be lost to surroundings (not insulated enough)
- Not all of the reactants reacted (should stir properly / add one reactant in excess)
- Incomplete combustion

Why might you not be able to measure the enthalpy change of a reaction directly? (5)

- Impossible to react precise amount of water
- Difficult to ensure that the elements react to produce the correct products (e.g. might make CO2 instead of CO)
- Very difficult to measure the temperature rise of a solid
- Difficult to prevent solid from dissolving
- Too high energy change (dangerous - does NOT apply if question is about an experiment)

Enthalpy change

Heat energy change at constant pressure

Why might the enthalpy change value given in a question not be an average value?

It is the only substance with that bond

4 ways to improve calorimetry (exothermic) practical

1. Insulate beaker / use polystyrene cup and lid: reduce heat loss to surroundings
2. Record temp for suitable time before adding metal: establish accurate initial temperature
3. Record temp at regular intervals: plot temperature results against time on graph
4. Extrapolate cooling back to point of addition: establish theoretical max temperature change
(NOTE: sometimes, only 2 OR 3 are accepted)

Why would the mean enthalpy change be similar to the experimental value of the enthalpy change?

The enthalpy change is for the same reaction.

Give a way of improving a measurement of temperature change without changing the apparatus?

Using a greater mass / number of moles of substances

Temperature change = 38°C, Uncertainty = ±0.5°C
Why is the thermometer adequate for the experiment?

The temperature change is much bigger than the uncertainty

If the original question to calculate the enthalpy of combustion uses bond enthalpies, why might the result for the same alkane in a liquid form be different?

Bond enthalpies use gaseous alkanes - the liquid alkane would need energy to convert it into a gas

When measuring and calculating the enthalpy change of an endothermic reaction (e.g. enthalpy of hydration), why might the calculated enthalpy change be lower than the actual one? (2)

Heat gain from the surroundings (as reduced temperature change, final temperature higher) or incomplete dissolving