1. Energy & Chemistry 2H2(g) + O2(g) → 2H2O(g) + heat and light
This can be set up to provide ELECTRIC ENERGY in a fuel cell.
Oxidation:
2 H2 → 4 H e-
Reduction:
4 e- + O H2O → 4 OH-
H2/O2 Fuel Cell
Energy, page 288

2. Energy & Chemistry ENERGY is the capacity to do work or transfer heat.
HEAT is the form of energy that flows between 2 objects because of their difference in temperature.
Other forms of energy —
light
electrical
kinetic and potential
Positive and negative particles (ions) attract one another.
Two atoms can bond
As the particles attract they have a lower potential energy
NaCl — composed of Na+ and Cl- ions.

3. Potential & Kinetic Energy
Kinetic energy — energy of motion.

5. Internal Energy (E) The higher the T the higher the internal energy
PE + KE = Internal energy (E or U)
Internal Energy of a chemical system depends on
number of particles
type of particles
temperature
The higher the T the higher the internal energy
So, use changes in T (∆T) to monitor changes in E (∆E).

6. Thermodynamics
Thermodynamics is the science of heat (energy) transfer.
Heat transfers until thermal equilibrium is established.
∆T measures energy transferred.
SYSTEM
The object under study
SURROUNDINGS
Everything outside the system

7. Directionality of Heat Transfer
Heat always transfer from hotter object to cooler one.
EXOthermic: heat transfers from SYSTEM to SURROUNDINGS.
T(system) goes down
T(surr) goes up

8. Directionality of Heat Transfer
Heat always transfers from hotter object to cooler one.
ENDOthermic: heat transfers from SURROUNDINGS to the SYSTEM.
T(system) goes up
T (surr) goes down

9. Energy Change in Chemical Processes
Energy & Chemistry
All of thermodynamics depends on the law of
CONSERVATION OF ENERGY.
The total energy is unchanged in a chemical reaction.
If PE of products is less than reactants, the difference must be released as KE.
Energy Change in Chemical Processes
Potential Energy of system dropped. Kinetic energy increased. Therefore, you often feel a Temperature increase.

10. HEAT CAPACITY Which has the larger heat capacity?
The heat required to raise an object’s T by 1 ˚C.
Which has the larger heat capacity?

11. Specific Heat Capacity
How much energy is transferred due to Temperature difference?
The heat (q) “lost” or “gained” is related to
a) sample mass
b) change in T and
c) specific heat capacity
Specific heat capacity =
heat lost or gained by substance (J)
(mass,
g) (T change, K)

12. Table of specific heat capacities
Substance
Phase
cp J g-1 K-1
Cp J mol-1 K-1
Air (typical room conditionsA)
gas
1.012
29.19
Aluminium
solid
0.897
24.2
Argon
0.5203
Copper
0.385
24.47
Diamond
0.5091
6.115
Ethanol
liquid
2.44
112
Gold
0.1291
25.42
Graphite
0.710
8.53
Helium
5.1932
Hydrogen
14.30
28.82
Iron
0.450
25.1
Lithium
3.58
24.8
Mercury
0.1395
27.98
Nitrogen
1.040
29.12
Neon
1.0301
Oxygen
0.918
29.38
Uranium
0.116
27.7
Water
gas (100 °C)
2.080
37.47
liquid (25 °C)
4.1813
75.327
solid (0 °C)
2.114
38.09
All measurements are at 25 °C unless noted. Notable minimums and maximums are shown in maroon text.
Aluminum
A Assuming an altitude of 194 meters above mean sea level (the world–wide median altitude of human habitation), an indoor temperature of 23 °C, a dewpoint of 9 °C (40.85% relative humidity), and 760 mm–Hg sea level–corrected barometric pressure (molar water vapor content = 1.16%).

13. Specific Heat Capacity
If 25.0 g of Al cool from 310 oC to 37 oC, how many joules of heat energy are lost by the Al?

14. Heat/Energy Transfer No Change in State
q transferred = (sp. ht.)(mass)(∆T)

15. Heat Transfer
Use heat transfer as a way to find specific heat capacity, Cp
55.0 g Fe at 99.8 ˚C
Drop into 225 g water at 21.0 ˚C
Water and metal come to 23.1 ˚C
What is the specific heat capacity of the metal?

16. Heating/Cooling Curve for Water
Note that T is constant as ice melts or water boils

17. Chemical Reactivity
But energy transfer also allows us to predict reactivity.
In general, reactions that transfer energy to their surroundings are product-favored.
So, let us consider heat transfer in chemical processes.

18. FIRST LAW OF THERMODYNAMICS
heat energy transferred
∆E = q + w
work done
by the
system
energy
change
Energy is conserved!

19. The First Law of Thermodynamics
Exothermic reactions generate specific amounts of heat.
This is because the potential energies of the products are lower than the potential energies of the reactants.

20. The First Law of Thermodynamics
There are two basic ideas of importance for thermodynamic systems.
Chemical systems tend toward a state of minimum potential energy.
Chemical systems tend toward a state of maximum disorder.
The first law is also known as the Law of Conservation of Energy.
Energy is neither created nor destroyed in chemical reactions and physical changes.

21. ∆E = q + w SYSTEM heat transfer in (endothermic), +q heat transfer out
(exothermic), -q
SYSTEM
∆E = q + w
w transfer in
(+w)
w transfer out
(-w)

22. ENTHALPY ∆H = Hfinal - Hinitial Heat transferred at constant P = qp
Most chemical reactions occur at constant P, so
Heat transferred at constant P = qp
qp = ∆H where H = enthalpy
and so ∆E = ∆H + w (and w is usually small)
∆H = heat transferred at constant P ≈ ∆E
∆H = change in heat content of the system
∆H = Hfinal - Hinitial

23. ENTHALPY ∆H = Hfinal - Hinitial Process is ENDOTHERMIC
If Hfinal > Hinitial then ∆H is positive
Process is ENDOTHERMIC
If Hfinal < Hinitial then ∆H is negative
Process is EXOTHERMIC

24. Exothermic reaction — heat is a “product” and ∆H = – 241.8 kJ
USING ENTHALPY
Consider the formation of water
H2(g) + 1/2 O2(g) → H2O(g) kJ
Exothermic reaction — heat is a “product” and ∆H = – kJ

25. USING ENTHALPY
Making liquid H2O from H2 + O2 involves two exothermic steps.
H2 + O2 gas
H2O vapor
Liquid H2O
Making H2O from H2 involves two steps.
H2(g) + 1/2 O2(g) → H2O(g) kJ
H2O(g) → H2O(l) kJ
H2(g) + 1/2 O2(g) → H2O(l) kJ
Example of HESS’S LAW—
If a rxn. is the sum of 2 or more others, the net ∆H is the sum of the ∆H’s of the other rxns.

26. Hess’s Law & Energy Level Diagrams
Forming H2O can occur in a single step or in a two steps. ∆Htotal is the same no matter which path is followed.
Active Figure 6.18

27. Hess’s Law
Hess’s Law of Heat Summation, Hrxn = H1 +H2 +H , states that the enthalpy change for a reaction is the same whether it occurs by one step or by any (hypothetical) series of steps.
Hess’s Law is true because H is a state function.
If we know the following Ho’s

28. Hess’s Law
For example, we can calculate the Ho for reaction [1] by properly adding (or subtracting) the Ho’s for reactions [2] and [3].
Notice that reaction [1] has FeO and O2 as reactants and Fe2O3 as a product.
Arrange reactions [2] and [3] so that they also have FeO and O2 as reactants and Fe2O3 as a product.
Each reaction can be doubled, tripled, or multiplied by half, etc.
The Ho values are also doubled, tripled, etc.
If a reaction is reversed the sign of the Ho is changed.

29. Hess’s Law Hess’s Law in a more useful form.
For any chemical reaction at standard conditions, the standard enthalpy change is the sum of the standard molar enthalpies of formation of the products (each multiplied by its coefficient in the balanced chemical equation) minus the corresponding sum for the reactants.

30. Hess’s Law

31. Hess’s Law Given the following equations and Hovalues
calculate Ho for the reaction below.

32. Hess’s Law
Use a little algebra and Hess’s Law to get the appropriate Hovalues

33. Some Thermodynamic Terms
Notice that the energy change in moving from the top to the bottom is independent of pathway but the work required may not be!
Some examples of state functions are:
T (temperature), P (pressure), V (volume), E (change in energy), H (change in enthalpy – the transfer of heat), and S (entropy)
Examples of non-state functions are:
n (moles), q (heat), w (work)
∆H along one path = ∆H along another path
This equation is valid because ∆H is a STATE FUNCTION
These depend only on the state of the system and not how it got there.
V, T, P, energy — and your bank account!
Unlike V, T, and P, one cannot measure absolute H. Can only measure ∆H.

34. Some Thermodynamic Terms
The properties of a system that depend only on the state of the system are called state functions.
State functions are always written using capital letters.
The value of a state function is independent of pathway.
An analog to a state function is the energy required to climb a mountain taking two different paths.
E1 = energy at the bottom of the mountain
E1 = mgh1
E2 = energy at the top of the mountain
E2 = mgh2
E = E2-E1 = mgh2 – mgh1 = mg(h)

35. Standard States and Standard Enthalpy Changes
Thermochemical standard state conditions
The thermochemical standard T = K.
The thermochemical standard P = atm.
Be careful not to confuse these values with STP.
Thermochemical standard states of matter
For pure substances in their liquid or solid phase the standard state is the pure liquid or solid.
For gases the standard state is the gas at 1.00 atm of pressure.
For gaseous mixtures the partial pressure must be 1.00 atm.
For aqueous solutions the standard state is 1.00 M concentration.
∆Hfo = standard molar enthalpy of formation
the enthalpy change when 1 mol of compound is formed from elements under standard conditions.
See Table 6.2 and Appendix L

36. Enthalpy Values
Depend on how the reaction is written and on phases of reactants and products
H2(g) + 1/2 O2(g) → H2O(g) ∆H˚ = -242 kJ
2H2(g) + O2(g) → 2H2O(g) ∆H˚ = -484 kJ
H2O(g) → H2(g) + 1/2 O2(g) ∆H˚ = +242 kJ
H2(g) + 1/2 O2(g) → H2O(l) ∆H˚ = -286 kJ

37. ∆Hfo, standard molar enthalpy of formation
H2(g) + ½ O2(g) → H2O(g) ∆Hf˚ (H2O, g)= kJ/mol
C(s) + ½ O2(g) → CO(g) ∆Hf˚ of CO = kJ/mol
By definition, ∆Hfo = 0 for elements in their standard states.
Use ∆H˚’s to calculate enthalpy change for
H2O(g) + C(graphite) → H2(g) + CO(g)

38. Using Standard Enthalpy Values
In general, when ALL enthalpies of formation are known,
∆Horxn =  ∆Hfo (products) -  ∆Hfo (reactants)
Remember that ∆ always = final – initial

39. Using Standard Enthalpy Values
Calculate the heat of combustion of methanol, i.e., ∆Horxn for
CH3OH(g) + 3/2 O2(g) → CO2(g) + 2 H2O(g)
∆Horxn =  ∆Hfo (prod) -  ∆Hfo (react)