1. Energy and Chemical Reactions
Chapter 6
Energy and Chemical Reactions
James P. Joule Discovered
mechanical equivalent of heat, which led
to the First Law of Thermodynamics.
Germain Henri Hess
Hess’s Law. 1 1 1 1

2. Kinetic and Potential Energy
The Nature of Energy
Kinetic and Potential Energy
Kinetic energy is the energy of motion.
Potential energy is energy which is available, such as an object that can fall a distance, potentially releasing kinetic energy when it hits the ground.

3. The Nature of Energy Energy Units SI Unit for energy is the joule, J:
We sometimes use the calorie instead of the joule:
1 cal = J (exactly)
A nutritional Calorie:
1 Cal = 1000 cal = 1 kcal
SI Unit for energy is the joule, J:
For example, energy of motion (kinetic energy):
(in J)

4. The Nature of Energy
All objects and substances have an internal energy, E
(this is a form of potential energy, and in this course is
commonly called “chemical energy”).
In chemistry, we’re interested in the internal energies of reactants and products and how it changes (ΔE) in going from:
reactants products
We usually can only measure ΔE, not E
“Δ” means change in

5. First Law of Thermodynamics
Total energy is conserved
Energy cannot be created or destroyed.
In chemistry, we are concerned with heat and
chemical energy.
The first law of thermodynamics tells us that the
sum of heat and chemical energy must be
constant (at constant pressure).
Heat at constant pressure is called enthalpy.

6. DH = Hfinal - Hinitial = qP
Enthalpy
Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure.
Can only measure the change in enthalpy:
DH = Hfinal - Hinitial = qP
We will use enthalpy (H) and energy (E) interchangeably

7. Endothermic and Exothermic Processes
Endothermic: system absorbs heat from surroundings.
(ΔH = +)
Exothermic: system transfers heat to the surroundings.
(ΔH = -)
An endothermic reaction feels cold.
An exothermic reaction feels hot.

8. Enthalpies of Reaction
For a reaction
DHrxn = H(products) - H (reactants)
Enthalpy is an extensive property (magnitude DH is directly proportional to amount):
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) DH = -802 kJ
2CH4(g) + 4O2(g)  2CO2(g) + 4H2O(g) DH = kJ
When we reverse a reaction, we change the sign of DH:
CO2(g) + 2H2O(g)  CH4(g) + 2O2(g) DH = +802 kJ
Change in enthalpy depends on state (gas, liquid, solid):
H2O(g)  H2O(l) DH = -44 kJ

9. q = (specific heat)  (grams of substance)  T
Calorimetry
Heat Capacity and Specific Heat
Calorimetry = measurement of heat flow.
Calorimeter = apparatus that measures heat flow.
Heat capacity = the amount of energy required to raise the temperature of a substance (by 1o).
Molar heat capacity = heat capacity of 1 mol of a substance (energy required to raise temp. of one mole
of a substance (by 1o)
Specific heat = specific heat capacity = heat capacity of 1 g (energy required to raise temp of one gram by 1o.
Heat absorbed by system, q, is given by:
q = (specific heat)  (grams of substance)  T
(Be careful of the sign of q)

10. Calorimetry Some specific heats (in J/g-K)… N2(g): 1.04 Al(s): 0.90
Fe(s): 0.45
CH4(g): 2.20
CaCO3(s): 0.82
H2O(l): 4.18
NOTE the very high specific
heat for liquid water.
This has a moderating influence
on the weather and the
temperature of living organisms.

11. Calorimetry Constant-Pressure Calorimetry
Atmospheric pressure is constant!
DH = qP
Heat emitted by reaction (qrxn)
is absorbed by solution (qsoln)
So, qrxn = -qsoln
qrxn =
-qsoln = -(specific heat of soln)  (grams of solution)  DT.

12. Calorimetry
9.55g sample of solid NaOH dissolves in g water in coffee-cup calorimeter. Temp rises from 23.6 to 47.4oC. Calculate ΔH (kJ/mol) for the soln process:
NaOH(s) Na+ (aq) + OH- (aq)
Assume specific heat of soln is same as that of pure water (which is 4.18 J/g-K).

13. Calorimetry qrxn = -Ccal ΔT
Bomb Calorimetry (Constant-Volume Calorimetry)
Used to study combustion processes
qrxn = -Ccal ΔT
where Ccal is the heat capacity of the calorimeter (to be determined experimentally)

14. Hess’ Law
Hess’ law: if a reaction is carried out in a number of steps, H for the overall reaction is the sum of H for each individual step.
Note that: H1 = H2 + H3

15. Hess’ Law
A consequence of Hess’ Law is that reactions can be added just like algebraic equations.
For example:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) H = -802 kJ
2H2O(g)  2H2O(l) H = -88 kJ
CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H = -890 kJ

16. Enthalpies of Formation
If 1 mol of compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation, Hof .
This is also called the heat of formation.
Standard conditions (standard state): 1 atm and 25 oC (298 K).

17. Enthalpies of Formation
Standard enthalpy of formation of the most stable form of an element is zero. For example, it is zero for C(s), H2(g), Pb(s), O2(g), etc

18. Enthalpies of Formation
Using Enthalpies of Formation to Calculate Enthalpies of Reaction.
We use Hess’ Law to calculate enthalpies of a reaction from enthalpies of formation
Example:
Combustion of 1 mol of propane gas, C3H8 (g)
Can construct alternate path of three steps shown as 1, 2 and 3
Hrxn =H1 + H2 +H3

19. Enthalpies of Formation
Using Enthalpies of Formation to Calculate Enthalpies of Reaction
For a reaction:
(a) 2 SO2(g) + O2(g) SO3 (g) -- what is ΔH?
(c) 4 FeO (s) + O2 (g) Fe2O3 (s) -- what is ΔH?

20. Foods and Fuels
1 Nutritional Calorie = kcal
Foods

21. Foods and Fuels Fuels U.S.: 1.0 x 106 kJ of fuel per day.
Most from petroleum and natural gas.
Remainder from coal, nuclear, and hydroelectric.
Fossil fuels are not renewable.