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Flashcards in Energetics Deck (20):
1

temperature

measure of the average kinetic energy of the particles

2

What is the total energy in chemical reactions?

total energy conserved - energy cannot be created or destroyed

3

Can enthalpy directly be measured?

no but changes in enthalpy can be

4

How are enthalpy changes calculated?

from their effect on the temperature of the surroundings

5

equation for heat change of a substance

q = mc deltaT

6

what are the units of enthalpy changes?

kjmol-1

7

what do -ve enthalpy changes represent?

energy lost to surroundings (exothermic)

8

what do +ve enthalpy changes represent?

energy gained from surroundings (endothermic)

9

breaking bonds enthalpy change

endothermic (energy required)

10

making bonds enthalpy change

exothermic (energy released)

11

average bond enthalpy

energy needed to break one mol of a bond in a gaseous molecule

12

enthalpy of formation

enthalpy change when one mol of substance is formed from its elements in standard states

13

enthalpy of combustion

enthalpy change when one mol of substance is completely burnt in oxygen under standard conditions

14

standard enthalpy of reaction

enthalpy change for the given reaction at standard conditions

15

standard conditions

  • normal, most pure state
  • 100kPa
  • 298K

16

calorimetry method

  • measure heat transferred to surroundings from reactants
  • this is amount of energy that has been lost/gained by reaction
    • divide by number mol that reacted to give enthalpy of reaction

17

main sources of error in calorimetry experiments

  • heat loss to surroundings
  • specific heat capacity of equipment

18

How to construct a Hess cycle?

if 2mol are made multiply enthalpy change by 2

19

how do you calculate the enthalpy change of the reaction from bond enthalpies?

reaction enthalpy = sum bond enthalpies of reactants - sum bond enthalpies of products

20

reasons for differences between true reaction enthalpy and bond calculated reaction enthalpy?

  • bonds are not average
  • usually reactants and products are not in gaseous state as in definition