1. Energy Transformations
Thermochemistry
Energy Transformations

2. Calorimetry
Calorimetry: The measurement of the heat into or out of a system for chemical and physical processes.
Based on the fact that the heat absorbed = the heat released
The device used to measure the absorption or release of heat is called a Calorimeter

3. Endothermic Reaction Endothermic reactions –
chemical reaction that absorbs energy to break existing bonds
Heat goes into the reaction (system) from the surroundings
The surroundings will feel colder
Temperature of endothermic reactions goes down
The sign for the heat change (enthalpy) will be positive
Ex: Boiling Water or Melting Ice (absorbing energy)

4. Endothermic Reaction

5. Exothermic Reactions Exothermic reactions –
chemical reaction in which energy is released
Heat goes out of the reaction (system) into the surroundings
The surroundings will feel hotter
Temperature of exothermic reactions goes up
The sign for the heat change(enthalpy) will be negative
EX: Condensation (gas to liquid) releasing energy!

6. Exothermic Reaction

7. Endo- and Exothermic ENDOTHERMIC EXOTHERMIC Heat Heat qsys < 0
Surroundings
Heat
qsys > 0
System
Surroundings
System
Heat
qsys < 0
DEMONSTRATION
example of an ENDOTHERMIC reaction
- hydration of ammonium nitrate
(NH4)+(NO3)- (s) + H2O -> NH4 (aq) + NO3 (aq)
(cools surroundings enough to freeze water)
WHY SPONTANEOUS ?
ENDOTHERMIC
EXOTHERMIC

8. Endothermic & Exothermic Reactions
Are the following reactions endothermic or exothermic?
CO + 3H2  CH4 + H2O H= -206kJ
I add magnesium metal to some hydrochloric acid. The temperature goes from 23C to 27 C
I mix together some vinegar & baking soda. The temperature goes from 28C to 23C

9. Chemistry Happens in Moles
An equation that includes energy is called a thermochemical equation
CH4 + O2 → CO2 + H2O ΔH = kJ
1 mole of CH4 releases kJ of energy
STEPS: Convert to moles and then multiply by the energy given over the number of moles of the compound!
If 10.3 g of CH4 are burned completely, how much heat will be produced?

10. The Work

11. Thermochemical equations
S + O2  SO2 H = – kJ
If we change the equation, then the H also changes …
SO2  S + O2 H = kJ
If the reaction is reversed the sign is reversed
Also, if numbers in the equation change, so will the amount of energy produced/absorbed:
2S + 2O2  2SO2 H = – kJ

12. Heat and Changes of State
Heat of combustion (∆H)= the heat of reaction for the complete burning of one mole of a substance
Molar heat of fusion (∆Hfus)= the heat absorbed by one mole of a substance in melting from a solid to a liquid at a constant temperature
Molar heat of solidification (∆Hsolid)= heat lost when one mole of a liquid freezes to a solid at a constant temperature (equal to the negative heat of fusion)
Molar heat of vaporization (∆Hvap)= the heat absorbed by one mole of a substance in vaporizing from liquid to a gas
Molar heat of condensation (∆Hcond)= heat released by one mole of a vapor as it condenses

13. Example (Heat of combustion)
The standard heat of combustion (∆H°rxn) for glucose (C6H12O6) is kJ/mol. If you eat and burn 71 g of glucose in one day, how much energy are you getting from the glucose?
C6H12O6 + 6 O2 → 6 CO2 + 6 H2O ΔH = kJ
Step one: convert g of glucose to moles
Step two: Use (∆H°rxn) to find amount of kJ gained