Energetics

Question 1

Definition: Enthalpy Change
Change in heat energy at constant pressure

Question 2

The symbol for enthalpy change measured under standard conditions is ∆Hθ

Question 3

Standard conditions:
100kPa Pressure
298K (25°C)
1.0 mol dm-3 concentration for all solutions

Question 4

Endothermic reactions absorb energy from the surroundings
Temperature of surroundings decreases

Question 5

Endothermic reactions:
Products are higher in energy than Reactants
∆H is positive +

Question 6

Exothermic reactions release energy to the surroundings
Temperature of surroundings increases

Question 7

Exothermic reactions:
Products lower in energy than reactants
∆H is negative -

Question 8

BENMEX
Breaking bonds is endo, making bonds is exo

Question 9

Enthalpy change = Total energy to break bonds - total energy released forming bonds

Reactants - Products

Question 10

Definition: Activation Energy
Minimum needed energy to start a reaction

Question 11

Definition: Mean bond enthalpy
Energy required to break one mole of a covalent bond into gaseous atoms averaged over a range of different compounds

Question 12

Mean bond enthalpies are averaged over a range of compounds, because the exact amount of energy required to break a covalent bond depends upon the molecule that bond is in. For example, the energy required to break a C—H bond in methane (CH4) is different to the energy required to break a a C—H bond in methanol (CH3OH)

Question 13

Mean bond enthalpies are always endothermic processes and, therefore, have a positive sign. This is because energy is required to break bonds. The more positive the bond enthalpy the larger the amount of energy needed to break the bond and so the stronger the bond

Question 14

Examples of mean bond enthalpies:
CH4(g) → CH3(g) + H(g)
H2(g) → 2H(g)
CO2(g) → CO(g) + O(g)
HCl(g) → H(g) + Cl(g)

Question 15

The standard enthalpy of formation Δf Hθ is:

· The enthalpy change when one mole of substance is formed
· from its constituent elements under standard conditions
· with all reactants and products being in their standard states.

Question 16

Examples of standard enthalpy formation:
H2(g) + ½O2(g) → H2O(l)
N2(g) + 2H2(g) + 3/2O2(g) → NH4NO3(s)

Question 17

Question: Why does the following reaction NOT show the standard enthalpy of formation of ammonia?
N2 + 3H2 → 2NH3
ΔHθ = −92 kJ mol−1

Answer 17

Answer: This reaction produces TWO moles of ammonia

Question 18

The standard enthalpy of formation of all elements in their standard states is zero

Question 19

Question: State why the enthalpy of formation of Na(s) is zero

Answer 19

Answer: Na is an element

Question 20

Question: State why the enthalpy of formation of liquid Na is not zero

Answer 20

Answer: Liquid is not the standard state of Na

Question 21

Definition: Standard enthalpy of combustion

· The enthalpy change when one mole of a substance
· is completely burnt in excess oxygen
· under standard conditions and all reactants and products being in their standard states

Question 22

Combustion involves burning a substance in oxygen. This is almost always exothermic, so standard enthalpies of combustion almost always have negative values

Question 23

Examples of standard enthalpy of combustion:
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
C2H5OH(l) + 3O2(g) → 2CO2(g) + 3H2O(l)

Question 24

Substances that cannot be combusted like water, carbon dioxide and most other oxides, have zero enthalpy of combustion

Question 25

Question: Why may the enthalpy of combustion of a compound be difficult to measure?

Answer 25

Answer: Incomplete combustion may occur

Question 26

Enthalpy of Reaction (ΔrH):

Any reaction that doesn’t match one of the above definitions is referred to as an enthalpy of reaction

Question 27

There are some types of enthalpy change that are very difficult to measure experimentally.
This could be due to:

· Reaction may have a very high activation energy
· Reaction rate may be too slow
· Possibility of competing reactions

Question 28

Definition: Hess Law

Hess’ Law states that the enthalpy change for a chemical reaction is the same, regardless of whatever route is taken from reactants to products.

Question 29

In Hess cycles, if a route goes in the opposite direction to an arrow, you must change the sign of the enthalpy change for the arrow

Question 30

Question: Give one reason why the bond enthalpy that you calculated above is different from the mean bond enthalpy quoted in a data book

Answer 30

Answer: The data book value is averaged over a range of different compounds

Question 31

Hess Cycles: Bond Enthalpy
The reaction arrows always point towards central product

A → B
↓ ↓
C

Question 32

Hess Cycles: Formation Enthalpy:
The reaction arrows point out from the central product as both A and B are formed from the elements at C

A → B
↑ ↑
C

Question 33

Hess Cycles: Combustion Enthalpy
The reaction arrows always point towards central product which is always H2O and CO2

A → B
↓ ↓
H2O + CO2

Question 34

Equation: ∆H for bond enthalpy
Σ bond energies broken - Σ bond energies made

Question 35

Equation: ∆H for formation enthalpy
Σ∆fH products - Σ∆fH reactants

Question 36

Equation: ∆H for combustion enthalpy
Σ∆cH reactants - Σ∆cH products

Question 37

Calorimetry measures changes in temperature of the surroundings and can determine the enthalpy change of a reaction

Question 38

Q = mc∆t

Q - heat energy lost or gained (J)
m - mass of water/solution (g)
c - specific heat capacity of water (4.18J g-¹ K-¹)
t - temperature change (K)

Question 39

∆H = Q ÷ n

∆H - enthalpy change (kJ mol-1)
Q - energy taken in/released (kJ)
n - number of moles reacted/formed

Question 40

Definition: Specific heat capacity

The energy required to raise 1g of substance by 1K without a change of state

Question 41

Key Point for combustion reaction calculations:

-1g of water = 1cm3 of water
-temperature change is the same in both K and °C
-Q is calculated in J using Q = mcΔT
-Q must be in kJ when using ΔH = Q/m

-Sign of ΔH
If temperature of water increases it is exothermic so −ΔH
If temperature of water decreases it is endothermic so +ΔH

Question 42

Key Point for 2 Solutions Calorimetry:

-Use the total water volume when calculating the mass of water
-Use concentration and volume of solutions to find moles

Question 43

Key Point for Solid to a Solution calculations:

-Find moles of solution using concentration and volume (do not find moles of solid in excess)
-Use limiting moles in calculation

Question 44

Assumption made in calorimetry:

The energy transferred to the water from the reacting chemicals (or vice versa for endothermic reactions) is equal to the energy released (or taken in) by the reaction.

Question 45

Sources of error: Combustion calorimetry

-Heat loss to the surroundings
-Incomplete combustion of the fuel
-Heat energy transferred to the metal calorimeter
-Some fuel evaporates

Question 46

Improvements to Minimise Sources of Error: Combustion calorimetry

· Add a lid – reduces heat loss
· Insulate sides of calorimeter – reduces heat loss
· Reduce distance between flame and beaker – reduces heat loss
· Put sleeve around flame to protect it from draughts

Question 47

Sources of Error: Solution Calorimetry

-Heat loss to surroundings

Question 48

Improvements to Minimise Sources of Error: Solution Calorimetry

· Add a lid – reduces heat loss
· Insulate calorimeter – reduces heat loss

Question 49

Steps to Measure an Enthalpy Change Using a Cooling Curve

· Record the temperature for a suitable time (3 minutes) before adding reactants together
· To establish an accurate initial temperature
· Mix reactants then record temperature every minute until a trend is seen

· Plot a graph of temperature against time
· Extrapolate the cooling curve back to the point of addition
· To establish a theoretical temperature change accounting for heat loss