# Chapter 9: Enthalpy (9.1-9.4) Flashcards

1
Q

exothermic reactions

A
• products have less enthalpy than the reactants (negative ΔH)
• chemicals lose energy, the surroundings gain this energy (hence increase in temp)
2
Q

endothermic reactions

A
• products have more enthalpy than the reactants (positive ΔH)
• the surroundings lose energy (hence decrease in temp) which goes to the chemicals
3
Q

standard enthalpy changes are measured using the standard conditions.
what are those conditions?

A

1 atmosphere pressure / 100 kPa
room temperature of 25 degrees / 298K
solutions must have a concentration of 1mol/dm3
all substances in their standard states (physical state under standard conditions)

4
Q

equation for the heat energy change

A

q = m X c X ΔT

q: heat energy change in joules
m: mass of surroundings (the thing that you measure the temp change of) in grams
ΔT: temperature change

5
Q

unit for ΔH

A

kJ/mol

make sure you work it out per mole!

6
Q

common errors when determining enthalpy change of combustion (experimentally) and how to minimise them

A

errors: heat loss to surroundings, incomplete combustion of reactant in spirit burner, evaporation of reactant from wick, non-standard conditions being used (heated up during combustion)
solutions: adding a lid to water beaker, using draft shields around apparatus, insulating system

7
Q

breaking bonds

A

endothermic (positive ΔH)

energy needed from the surroundings to break the bonds

8
Q

making bonds

A

exothermic (negative ΔH)

9
Q

how to work out the enthalpy change of a reaction from average bond enthalpies - gaseous molecules of covalent substances

A

ΔrH = ∑(bond enthalpies of reactants) - ∑(bond enthalpies of products)

10
Q

how activation energy effects reaction rate

A

small activation energy: rapid (energy needed to break bonds is readily available from surroundings)
Higher activation energies: slow (large energy barrier that reactions happen very slowly or not at all)

11
Q

what is average bond enthalpy

A

the mean amount of energy required to break 1 mole of a specified type of covalent bond in a gaseous molecule

12
Q

limitations of average bond enthalpies

A

The bond is in a different environment

Actual value for a bond in a certain molecule may be slightly higher/lower.

13
Q

state Hess’ law

A

Hess’ law states that, if a reaction can take place by more than one route, and the starting and finishing conditions are the same, the total enthalpy change is the same for each route

14
Q

Enthalpy change of neutralisation

A

The standard enthalpy change of neutralisation is the enthalpy change when solutions of an acid and an alkali react together under standard conditions to produce 1 mole of water

15
Q

Enthalpy change of formation

A

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

16
Q

Enthalpy change of combustion

A

the enthalpy change that takes place when one mole of a substance reacts completely with oxygen under standard condition, all reactants and products being in their standard states.

17
Q

what is Hess’ law used for?

A

to calculate enthalpy changes that are not easy to measure directly in experiments

18
Q

Working out enthalpy change (e.g delta h in diagrams)

A

H(products) - H(reactants)