1. Thermochemistry Chapter 11 Unit 12
THE FLOW OF ENERGY – HEAT

2. In the 1500’s Sir Francis Bacon proposed the concept of heat.
Rub your hands vigorously together.
What is happening?
Why?

3. What would happen if you placed ice in your hand?
Why?

4. The transfer of energy is a major aspect of chemistry.
Thermochemistry: study of the transfer of heat energy in chemical reactions.

5. Energy Transformations:
The heat energy can be released or absorbed, depending on the chemical reaction that is being studied.
Energy: the capacity for doing work or supplying heat.

6. Heat: (q) is the energy transferred from one object to another because of a temperature difference between the two.

7. Heat itself cannot be detected only the change in temperature caused by the addition or loss of heat.
Heat always flows from a warmer object to a cooler object.

8. Exothermic and Endothermic Process
All energy involved in a process can be accounted for as work, stored energy or heat.

9. Endothermic Reaction absorbs heat from its surroundings
Endothermic Reaction absorbs heat from its surroundings. Surroundings get colder, q has a positive value.
A + B + heat → C

10. Exothermic reaction release heat into its surroundings
Exothermic reaction release heat into its surroundings. q has a negative value.
A + B → C + heat

11. Specific Heat * 1 Calorie = 1 kilocalorie = 1000 calories (c)
Specific Heat: the quantity of heat needed to raise the temperature of 1g of water 1 oC
* 1 Calorie = 1 kilocalorie = 1000 calories (c)

12. The statement “10g of sugar contain 41 Calories” means that 10 g of sugar releases 41 kilocalories of heat when completely burned.

13. The calorie is also related to the joule, the SI unit of heat and energy
* 1 J = cal or
* J = 1 cal

14. Heat affects the temperature of objects with a high specific heat much less than the temperature of those with lower specific heat.

15. Specific Heat ( C ) =
q = heat (J or cal)
m x T Mass (g) x change
in temperature (oC)

16. The units of specific heat can be expressed in Joules or calories/ (g x o C).

17. Example:
The temperature of a piece of copper with a mass of 95.4 g increases from 25 °C to 48 °C when the metal absorbs 849 J of heat. What is the specific heat of copper?
Answer: J/g x °C

18. MEASURING HEAT CHANGES

19. Calorimetry : the accurate and precise measurement of a heat change for chemical and physical processes.

20. Enthalpy (H) is the heat content in a chemical system.

21. Enthalpy and Reaction Tendency
In nature there is a tendency for processes to occur that lead to a lower energy state. Thus, there is a movement toward greater stability in a system.

22. Most reactions in nature are exothermic
Most reactions in nature are exothermic. These tend to occur spontaneously. This means the need for an external force or agent to assist the reaction is not needed.

23. Endothermic reactions require energy to occur
Endothermic reactions require energy to occur. These are generally not spontaneous in nature. These require the assistance of some agent.

24. The term spontaneous has no relationship with the speed at which the reactions occur. Rather, these are referencing to the lack of the need of an agent to assist the process.

25. The equation for Enthalpy is:
H = m x C x T,
Where T = Tf - Ti

26. The sign of H shows if the reaction was endothermic or exothermic { (-) for exo, (+) for endo} .

27. Calorimetry is an important tool in the food industry to determine the calorie content of foods.

28. Thermochemical Equations:
the amount of heat released is shown as a product of the reaction.
Chemical Reactions are accompanied by energy transformations.
Ex. CaO(s) + H2O(l)  Ca(OH)2(s) KJ.

29. Thermochemical Equations
When ΔH is negative, then that amount of heat is released and will be on the products side of the equation.
When ΔH is positive, then that amount of heat is added into the equation which means that it is on the reactants side of the equation.

30. Physical changes of state involve the transfer of energy.
Heat of Fusion: the heat absorbed by a substance in melting from a solid to a liquid at a constant temperature.
All solids absorb heat as they melt to become liquids.

31. Bond Energy and Reaction Heat
The change in the heat content of a reaction system is related to:
The changes in the number of bonds breaking and forming.
The strengths of these bonds as the reactants form products.