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Flashcards in Module 3.2 Deck (90)
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1

Define enthalpy

The thermal energy stored in a chemical system

2

Define enthalpy change

The heat energy transferred in a reaction at constant pressure

3

What are the standard conditions?

298K and 100KPa

4

Why do we use standard conditions?

- To make meaningful calculations

5

How do we convert from kPa to Pa?

Multiply by 1000

6

Describe the graph of an exothermic reaction?

- Goes from high to low (reactants have a higher enthalpy than products)
- Products are more stable than reactants
- Enthalpy change is negative

7

Describe the graph of an endothermic reaction?

- Goes from low to high (products have a higher enthalpy than reactants)
- Reactants are more stable than reactants
- Enthalpy change is positive

8

On an exothermic/endothermic enthalpy graph, how can you show enthalpy change?

It is the direct difference between reactants and products

9

Define activation energy?

The minimal amount of energy needed for a reaction to take place

10

On an exothermic/endothermic enthalpy graph, how can you show activation energy?

The difference between the HIGHEST point of the activation energy and the enthalpy of reactants

11

Give an example of an exothermic reaction

Oxidation

12

Give an example of an endothermic reaction

Photosynthesis

13

Define enthalpy change of formation

The enthalpy change when one mole of compound is formed from its elements in their standard states under standard conditions

14

Define enthalpy change of combustion

The enthalpy change when one mole of a substance is completely burnt in oxygen under standard conditions

15

Define enthalpy change of neutralisation

The enthalpy change when one mole of water is produced when an acid and alkali react under standard conditions

16

What is meant by standard states

Physical states under standard conditions

17

What formula do we use to determine enthalpy change directly from experimental results?

q = mcΔT

q = Heat loss OR gained (joules)
m = Mass of water in the colorimeter (grams)
c = Specific heat capacity of water (4.1.8J/grams/Kelvin)
ΔT = Change in temperature (Kelvin)

18

What does the q stand for in q = mcΔT?

Enthalpy change (measured in joules)

19

What does the m stand for in q = mcΔT?

Mass of water (measured in grams)

20

What does the c stand for in q = mcΔT?

Specific heat capacity ( 4.18 J/grams/Kelvin)

21

What does the ΔT stand for in q = mcΔT?

Change in temperature (measured in Kelvin)

22

When using q = mcΔT how can we find ΔH?

ΔH = Q/number of moles

23

When using ΔH = Q/number of moles, what must we be careful about?

Whether the reaction is exothermic or endothermic.

If its exothermic add a negative in front of Q, if it's endothermic leave it positive

24

How do we calculate enthalpy change using q = mcΔT?

1) Write out the formulas required
2) Write down the values you have been given
3) Convert anything if necessary
4) Input numbers into the formula
5) Divide by 1000 to get from J to kJ
6) Watch out for the type of reaction (Exo/endo)

25

Before using the ΔH = Q/number of moles formula, what must you do?

- Establish whether the reaction is exothermic or endothermic and add its sign to Q
- Convert J to kJ (divide by 1000)

26

Define average bond enthalpy

The mean energy needed for one mole of a given type of gaseous bond to undergo homolytic fission

27

Why are calculated bond enthalpies different from the actual bond enthalpies?

They are different because, bond enthalpies used in calculations are an average. The actual bond enthalpy will vary depending on the rest of the chemical.

28

What are the 3 steps of every reaction?

Step 1) Bonds break
Step 2) Atoms rearrange
Step 3) Rearranged atoms form bonds together

29

Is bond-breaking endothermic or exothermic?

Endothermic

30

Is bond-making exothermic or endothermic?

Exothermic