1. Chemical Energy and Calorimetry
Paul Walsh
Devil’s Doorway II
American, 2011

2. Thermodynamics
Thermodynamics is the study of energy and its transformations, where energy is defined as the capacity to do work or transfer heat.
The First Law of Thermodynamics states that energy can neither be created nor destroyed, but is always conserved.
In other words, energy can only be converted from one form to another or from one place to another. Therefore, any energy lost by a system is gained by the surroundings, and any energy gained by the system must come from the surroundings.

3. Enthalpy of Reaction, ΔH
Enthalpy, ΔH, is equal to the heat gained or lost during a physical or chemical change at constant pressure. It is measured in kJ/mol.
As shown in the following energy diagrams, enthalpy is a state function. A state function is any process that depends only on the initial and final conditions, and not the pathway taken, when changes occur.

4. Chemical Reactions and Energy
In terms of physical and chemical changes, energy is absorbed to break chemical bonds and overcome attractive forces that hold atoms together.
Likewise, energy is released when atoms rearrange and bonds/attractions are formed between atoms and molecules.
All types of chemical reactions fall into two categories: Endothermic or Exothermic.

5. Energy Diagrams EXOTHERMIC(ΔH<0) ENDOTHERMIC (ΔH>0)
Exothermic reactions release more energy when products form than is required to break reactants apart.
Endothermic reactions require more energy to break apart reactants than is released when products form.

6. Calorimetry
Calorimetry is experimental method of measuring heat transfer, and is one way to calculate the change in enthalpy for a physical or chemical change.
All substances change temperature when they are heated, but the magnitude of this change varies from one substance to another.
Calorimeters are devices used to measure this change in temperature when objects absorb or release energy.

7. Heat Capacity
The temperature change experienced by an object when it absorbs a certain amount of heat is determined by its specific heat capacity, C.
Specific heat capacity is defined as the amount of heat required to raise the temperature of a one gram sample one degree Celsius (or Kelvin).
The most common heat capacity is that of water, 4.18 J/(g °C). This heat capacity is significantly higher than most substances, and changes negligibly when substances are dissolved in water.

8. Energy Calculations
The following equations can be used to determine the energy change that occurs during physical and chemical changes:
Q = m . C . ΔT
Q is heat (J), m is the mass (g), C is the specific heat (J/(g °C)), and ΔT the change in temperature (°C or K).
ΔH = -Q/n
where ΔH is the change in enthalpy (J/mol), Q is the heat gained or lost (J), and n is the number of moles (mol).

9. Examples 3.00x103 J The specific heat of iron metal is 0.450 J/(g °C).
How much heat is required to raise the temperature of a 105 g block of iron from 25.0°C to 88.5 °C?
3.00x103 J

10. Examples
2. If 861 J of energy is added to a sample of water and the temperature raises from 25oC to 35oC, what was the mass of the water sample?
21 g

11. Examples
3. A peanut is completely burned under a calorimeter containing 25 g of water. The water in the calorimeter was initially 19.00oC. After the burning of the peanut, the water’s temperature raised to oC. Assuming no energy was lost to the atmosphere, how much energy (in Joules and Calories) is in a peanut?
46 J or 11 cal