Unit 1: Section 4- Energetics + RP2

Question 1

What is the definition of enthalpy change?

Answer 1

The heat energy transferred in a reaction at constant pressure.

Question 2

What are the units for enthalpy change?

Answer 2

KJ/mol-1

Question 3

What does ΔH^∅ mean?

Answer 3

The substances were in their standard states and the measurement was made under standard conditions.

Question 4

What are standard conditions of temperature and pressure?

Answer 4

298K
100KPa

Question 5

Is ΔH for an exothermic reaction positive or negative?

Answer 5

negative

Question 6

Is ΔH for an endothermic reaction positive or negative?

Answer 6

positive

Question 7

What is the definition of bond enthalpy?

Answer 7

The energy needed to break a bond.

Question 8

How do you calculate enthalpy change?

Answer 8

Enthalpy change of reaction = total energy absorbed - total energy released

Question 9

Why is using Hess’s law more exact for calculating bond enthalpies than using the enthalpy change equation?

Answer 9

Because when using this equation, you are using average values from the mean bond enthalpies, so they are less accurate.

Question 10

What are the conditions needed for standard enthalpy of formation?

Answer 10

• 1 mole of compound formed
• formed from ‘its elements’
• in standard conditions

Question 11

What are the conditions needed for standard enthalpy of combustion?

Answer 11

• 1 mole of substance
• complete combustion
• in standard conditions (even tho not accurate as burning is not at room temp)

Question 12

What is the enthalpy change equation?

Answer 12

q=mcΔT
q- enthalpy change in joules
m- mass of substance being heated in grams
c- specific heat capacity (4.18)
T- change in temp in K

Question 13

How do you convert from °C to K?

Answer 13

+273

Question 14

Which way do the arrows point in a Hess’s cycle for enthalpy of formation?

Answer 14

away from the elements, pointing up

Question 15

Which way do the arrows point in a Hess’s cycle for enthalpy of combustion?

Answer 15

towards the products of combustion, pointing down

Question 16

What does Hess’s Law state?

Answer 16

Hess’s Law states that overall enthalpy change for a reaction is independent of the route it takes.

Question 17

Give an example of a exothermic reaction

Answer 17

combustion of fuels
oxidation of carbohydrates e.g. glucose in respiration

Question 18

Give an example of an endothermic reaction

Answer 18

thermal decomposition of calcium carbonate

Question 19

What is the enthalpy change of formation of an element?

Answer 19

0

Question 20

Which is more exothermic- complete or incomplete combustion?

Answer 20

complete combustion

Question 21

How can we know if enthalpy change is measured at standard conditions?

Answer 21

the symbol is used

Question 22

RP2- What is Hess’s law?

Answer 22

the enthalpy change for a chemical reaction is always the same, regardless of the route from reactants to products

Question 23

RP2- How can you reduce the uncertainty in the mass measurement?

Answer 23

Use a balance with a greater resolution
Use a larger mass

Question 24

RP2- How do you calculate percentage uncertainty?

Answer 24

100x absolute uncertainty/calculated value

Question 25

RP2- How can you calculate enthalpy change of reaction experimentally?

Answer 25

q=mcΔT
q- enthalpy change in joules
m- mass of substance being heated in grams
c- specific heat capacity (4.18)
T- change in temp in K

divide this number (q) by the number of moles of the reactant used

make sure to add a sign to show whether enthalpy change is exo or endo thermic

Question 26

RP2- How do you convert from degrees Celsius to Kelvin?

Answer 26

add 273

Question 27

RP2- Why may an experimental value for enthalpy change be different to the theoretical value?

Answer 27

1. heat loss to apparatus/surrounding
2. incomplete combustion
3. non-standard conditions
4. evaporation of alcohol/ water

Question 28

RP2- How do you prevent heat loss to surroundings/apparatus?

Answer 28

• insulate the beaker by placing it in a polystyrene cup with a lid
• avoid large temperature differences between surroundings and calorimeter

Question 29

RP2- Other than preventing heat loss, how can the accuracy of this experiment be improved?

Answer 29

• read the thermometer at eye level to avoid parallax errors
• stir the solution so the temperature is evenly distributed
• use a digital thermometer for more accurate and faster readings
• use greater concentrations and masses, leading to a greater temperature change and thus smaller uncertainty

Question 30

RP2- How can you calculate the number of moles of something?

Answer 30

conc x volume

Question 31

Define Hess’s law

Answer 31

the total enthalpy change for a reaction is
independent of the route by which the chemical change takes place

Question 32

Which way do arrows point when calculating ΔH for a formation reaction?

Answer 32

upwards

Question 33

How do you calculate enthalpy change for a reaction?

Answer 33

ΔH reaction = Σ ΔH products - Σ ΔH reactants

Question 34

Which way do arrows point when calculating ΔH for a combustion reaction?

Answer 34

downwards

Question 35

Define mean bond energies

Answer 35

the mean bond energy is the enthalpy needed to break the covalent bond into gaseous atoms, averaged over different molecules
this only applies when the substances start AND end in gaseous state

Question 36

Why do we use mean values of bond energies?

Answer 36

because every single bond in
a compound has a slightly different bond energy

Question 37

How do you calculate ΔH if all the substances are gases?

Answer 37

ΔH = Σ bond energies broken - Σ bond energies made

Question 38

What direction do the arrows point in for gaseous atoms of elements?

Answer 38

the arrows point downwards

Question 39

Why are ΔH values calculated via mean bond energies less accurate than using formation or combustion data?

Answer 39

because the mean bond energies are not exact

Question 40

What happens to the enthalpies of combustion for successive members of a homologous series?

Answer 40

there is a constant rise in the size of the enthalpies of combustion as the number of
carbon atoms increases.