1. Calorimetry and Enthalpy
Section 5.2

2. Specific Heat Capacity (c)
The amount of thermal energy that is needed to raise the temperature of one gram of a substance 1oC (or 1 K) (0 oC = K)
Symbol is “c”, Units of J/g●oC
Refer to Table 1 for Specific Heat Capacities of selected substances
Larger values indicate that more energy is required to increase its temperature
Substance
Concrete
Aluminum
Wood
Water (liquid)
Specific Heat Capacity (J/ g●oC)
0.88
0.900
1.76
4.184

3. Calorimetry
The experimental process of measuring thermal energy change in a chemical or physical change.
Calorimeter – a device that is used to measure thermal energy changes in a chemical or physical change.

4. Calorimetry
With coffee-cup calorimeters, the following assumptions must be made
Any thermal energy transferred from the calorimeter to the outside environment is negligible
Any thermal energy absorbed by the calorimeter itself is negligible
All dilute, aqueous solutions have the same density (1.00g/mL) and specific heat capacity (4.18J/(g●oC)) as water.

5. Calorimetry
In calorimetry, the total amount of thermal energy absorbed or released by a chemical system is given the symbol q.
The value of q can be calculated using the equation below:

6. Calorimetry The value of q has 2 parts:
The magnitude of q tells you how much energy is involved
Most experiments will have you calculating the q value of the surrounding environment (ie. water) since it is nearly impossible to calculate the temperature change of the substance.

7. Calorimetry
When you calculate the q value of the surrounding the sign tells you the direction of energy transfer.
If q has a negative value, this states that the temperature change is decreasing (remember Tfinal – Tinitial) in order for the temperature to decrease the system needs to be absorbing energy from the surrounding. Since the system is gaining energy from the surrounding this is considered an endothermic reaction.
If q has a positive value, this states that the temperature change is increasing (remember Tfinal – Tinitial) in order for the temperature to increase the system needs to be releasing energy to the surrounding. Since the system is losing energy to the surrounding this is considered an exothermic reaction.

8. Calorimetry
Since energy cannot be destroyed, the total thermal energy of the system and its surroundings remains constant.
qsystem + qsurroundings = 0 or
qsystem = - qsurroundings

9. Enthalpy the total amount of thermal energy in a substance
enthalpy change ∆H - is the energy released to or absorbed from the surroundings during a chemical or physical change.
∆Hsystem = Hproducts - ∆Hreactants
if pressure remains constant, the enthalpy change of the chemical system is equal to the flow of thermal energy in and out of the system
∆Hsystem = |qsystem|
*remember qsystem = - qsurroundings and you always calculate the qsurroundings *

10. Molar Enthalpy Molar Enthalpy Change (∆Hr)
the enthalpy change associated with a chemical or physical change involving 1 mol of a substance (J/mol)
∆Hsystem = n x ∆Hr (where n = moles of ∆Hr = ∆Hsystem / n the substance)
When you write the balanced equation for the molar enthalpy change of formation of a product, the coefficient of that product must always be 1.

11. Representing Exothermic and Endothermic Reactions
There are three different ways to represent the enthalpy change of exothermic and endothermic reactions.
Thermochemical Equations
Enthalpy as a Separate Expression
Potential Energy Diagrams

12. Thermochemical Equation
When writing an exothermic balanced chemical equation the heat term will be located on the product side as a positive value to indicate that heat is produced.
Ex. Exothermic : H2(g) + ½ O2(g)  H2O(l) kJ
When writing an endothermic balanced chemical equation the heat term will be located on the reactant side as a positive value to indicate that heat is absorbed.
Ex. Endothermic : kJ + MgCO3(s)  MgO(s) + CO2(g)

13. Enthalpy as a Separate Expression
Indicate the enthalpy beside the chemical equation.
∆H is always negative for an exothermic reaction and positive for an endothermic reaction.
Ex. Exothermic:
H2(g) + ½ O2(g)  H2O(l) ∆Hrxn = kJ
Ex. Endothermic:
MgCO3(s)  MgO(s) + CO2(g) ∆Hrxn = 117.3kJ

14. Potential Energy Diagrams

15. Example problem
A student places 50.0mL of liquid water at 21.00°C into a coffee-cup calorimeter. She places a sample of gold at °C into the calorimeter. The final temperature of the water is 21.33°C. The specific heat capacity of water is 4.18J/g°C and the density of water, d, is 1.00g/mL. Calculate the quantity of thermal energy, q, absorbed by the water in the calorimeter.

16. Example problem
Ethanol CH3CH2OH(l) is used to disinfect the skin prior to an injection. If a 1.00g sample of ethanol is spread across the skin and evaporated, what is the expected enthalpy change? The molar enthalpy of vaporization of ethanol is 38.6kJ/mol.

17. Example problem
The combustion of methane gas, CH4(g) is an exothermic reaction. When 1 mol of methane burns, 802.3kJ of energy is released. Write the thermochemical equation using both by representing the energy change as a ∆H value, and by representing the energy change as an energy term in the equation.

18. Example problem
Draw potential energy diagrams for the following reactions:

19. Example problem
Draw potential energy diagrams for the following reactions:

20. Example problem
A 50.0ml sample of 1.0mol/L aqueous solution of hydrochloric acid, HCl(aq), was mixed with 50.0ml of a 1.0mol/L aqueous solution of sodium hydroxide, NaOH(aq), at 25.0°C in a calorimeter. After the solutions were mixed by stirring, the temperature was 31.9°C. Determine the quantity of thermal energy transferred by the reaction to the water, q, and state whether the reaction was endothermic or exothermic.

21. Homework