1. Thermochemistry: Chemical Energy
Chapter 8
Thermochemistry: Chemical Energy

2. Energy Energy – capacity to supply heat or do work
Energy = Heat + Work
E = q + w
2 types of Energy
Potential Energy
Kinetic Energy

3. Energy Two fundamental kinds of energy. Law of Conservation of Energy
Potential energy is stored energy.
Kinetic energy is the energy of motion.
Law of Conservation of Energy
Energy can be converted from one kind to another but never destroyed

4. Energy Units Conversions SI Unit – Joule (J) Additional units
Calorie (Cal) – food calorie
calorie (cal) – scientific calorie
Conversions
1 cal = J
1000 cal = 1 Cal

5. Energy and Chemical Bonds
Chapter 6
Kept a careful accounting of atoms as they rearranged themselves
Reactions also involve a transfer of energy

6. Energy and Chemical Bonds
A chemical
Potential - attractive forces in an ionic compound or sharing of electrons covalent compound
Kinetic – (often in form of heat) occurs when bonds are broken and particles allowed to move
To determine the energy of a reaction it is necessary to keep track of the energy changes that occur during the reaction

7. Internal Energy and State Functions
In an experiment: Reactants and products are the system; everything else is the surroundings.
Energy flow from the system to the surroundings has a negative sign (loss of energy).
Energy flow from the surroundings to the system has a positive sign (gain of energy).

8. Internal Energy and State Functions
Tracking energy changes
Energy changes are measured from the point of view of the system (Internal Energy - IE)
Change in Energy of the system – ΔE
ΔE = Efinal - Einitial

9. Internal Energy and State Functions
IE depends on
Chemical identity, sample size, temperature, etc.
Does not depend on the system’s history
Internal Energy is a state function
A function or property whose value depends only on the present state (condition) of the system, not on the path used to arrive at that condition

10. Expansion Work E = q + w In physics w = force (F) x distance (d)
Force – energy that produces movement of an object
In chemistry w = expansion work
Force - the pressure that the reaction exerts on its container against atmospheric pressure hence it is negative
Distance – change in volume of the reaction
w = -PΔV

11. Energy and Enthalpy q = DE + PDV ΔE = q – PΔV
The amount of heat exchanged between the system and the surroundings is given the symbol q.
q = DE + PDV
At constant volume (DV = 0): qv = DE
At constant pressure: Energy due to heat and work but work minimal compared to heat energy
qp = DE + PDV = DH
Enthalpy change (heat of reaction): DH = Hproducts – Hreactants

12. The Thermodynamic Standard State
ΔH = amount of energy absorbed or released in the form of heat
DH = Hproducts – Hreactants
Important factors
States of matter
Thermodynamic standard state – most stable form of a substance at 1 atm and at a specified temperature, usually 25oC; and 1 M concentration for all substances in solution
DH – valid for the reaction as written including exact # of moles of substances
N2H4(g) + H2(g)  2 NH3(g) + heat (188 kJ)

13. Enthalpies of Physical and Chemical Change

14. Enthalpies of Physical and Chemical Changes
Enthalpies of Chemical Change: Often called heats of reaction (DHreaction).
Endothermic: Heat flows into the system from the surroundings and DH has a positive sign. Unfavorable Process
Exothermic: Heat flows out of the system into the surroundings and DH has a negative sign. Favorable process

15. Enthalpies of Physical and Chemical Changes
Reversing a reaction changes the sign of DH for a reaction.
C3H8(g) + 5 O2(g)  3 CO2(g) + 4 H2O(l) DH = –2219 kJ
3 CO2(g) + 4 H2O(l)  C3H8(g) + 5 O2(g) DH = kJ
Multiplying a reaction increases DH by the same factor.
3 C3H8(g) + 15 O2(g)  9 CO2(g) + 12 H2O(l) DH = –6657 kJ

16. Problems
How much heat (in kilojoules) is evolved or absorbed in each of the following reactions?
Burning of 15.5 g of propane: C3H8(g) + 5 O2(g)  3 CO2(g) + 4 H2O(l) DH = –2219 kJ
Reaction of 4.88 g of barium hydroxide octahydrate with ammonium chloride: Ba(OH)2·8 H2O(s) + 2 NH4Cl(s)  BaCl2(aq) + 2 NH3(aq) + 10 H2O(l) DH = kJ

17. Determination of Heats of Reaction
Experimentally – calorimetry
Hess’s Law
Standard Heat’s of Formation
Bond Dissociation Energies

18. Calorimetry and Heat Capacity
Calorimetry is the science of measuring heat changes (q) for chemical reactions. There are two types of calorimeters:
Bomb Calorimetry: A bomb calorimeter measures the heat change at constant volume such that q = DE.
Constant Pressure Calorimetry: A constant pressure calorimeter measures the heat change at constant pressure such that q = DH.

19. Calorimetry and Heat Capacity

20. Calorimetry and Heat Capacity
Heat capacity (C) is the amount of heat required to raise the temperature of an object or substance a given amount.
Specific Heat: The amount of heat required to raise the temperature of 1.00 g of substance by 1.00°C.
Molar Heat: The amount of heat required to raise the temperature of 1.00 mole of substance by 1.00°C. C = q D T

21. Problems
What is the specific heat of lead if it takes 96 J to raise the temperature of a 75 g block by 10.0°C?
When 25.0 mL of 1.0 M H2SO4 is added to 50.0 mL of 1.0 M NaOH at 25.0°C in a calorimeter, the temperature of the solution increases to 33.9°C. Assume specific heat of solution is J/(g–1·°C–1), and the density is 1.00 g/mL–1, calculate the heat absorbed or released for this reaction.

22. Hess’s Law Allows the enthalpy to be determined for:
Reactions that occur too quickly or take too long to use calorimetry
Reactions that are too dangerous
Works like the Haber process in chapter 6
Take reactions for which the heat is known and manipulate them to give the desired reaction

23. Standard Heats of Formation
Standard Heats of Formation (DH°f): The enthalpy change for the formation of 1 mole of substance in its standard state from its constituent elements in their standard states.
The standard heat of formation for any element in its standard state is defined as being ZERO.
DH°f = 0 for an element in its standard state

24. Standard Heats of Formation
H2(g) + 1/2 O2(g)  H2O(l) DH°f = –286 kJ/mol
3/2 H2(g) + 1/2 N2(g)  NH3(g) DH°f = –46 kJ/mol
2 C(s) + H2(g)  C2H2(g) DH°f = +227 kJ/mol
2 C(s) + 3 H2(g) + 1/2 O2(g)  C2H5OH(g) DH°f = –235 kJ/mol

25. Standard Heats of Formation
Calculating DH° for a reaction:
DH° = Σ[DH°f (products) x moles] – Σ[DH°f (Reactants) x moles]
For a balanced equation, each heat of formation must be multiplied by the stoichiometric coefficient.
aA + bB cC + dD
DH° = [cDH°f (C) + dDH°f (D)] – [aDH°f (A) + bDH°f (B)]

26. Problems
Calculate DH° (in kilojoules) for the reaction of ammonia with O2 to yield nitric oxide (NO) and H2O(g), a step in the Ostwald process for the commercial production of nitric acid.
Calculate DH° (in kilojoules) for the photosynthesis of glucose from CO2 and liquid water, a reaction carried out by all green plants.

27. Energy Calculations Other methods for calculating enthalpies
Bond dissociation energies – measures the energy given off by the formation of bonds in the products and substracts the energy required to break bonds in the reactants

28. Why do chemical reactions occur?
A chemical reaction will move from less stability to greater stability.
Achieved by giving off more energy than is absorbed by the reactants
This indicates that exothermic reactions occur by why do endothermic reactions occur?
Gibb’s Free Energy
DG = DH – TDS
DH – enthalpy, T – temperature, DS - entropy

29. An Introduction to Entropy
Second Law of Thermodynamics: Reactions proceed in the direction that increases the entropy of the system plus surroundings. (increases the degree of disorder)
A spontaneous process is one that proceeds on its own without any continuous external influence.
A nonspontaneous process takes place only in the presence of a continuous external influence.

30. An Introduction to Entropy

31. An Introduction to Entropy

32. An Introduction to Entropy
The measure of molecular disorder in a system is called the system’s entropy; this is denoted S.
Entropy has units of J/K (Joules per Kelvin).
DS = Sfinal – Sinitial
Positive value of DS indicates increased disorder (favorable).
Negative value of DS indicates decreased disorder (unfavorable).

33. Problems
Predict whether DS° is likely to be positive or negative for each of the following reactions. Using tabulated values, calculate DS° for each:
a. 2 CO(g) + O2(g)  2 CO2(g) b. 2 NaHCO3(s)  Na2CO3(s) + H2O(l) + CO2(g) c. C2H4(g) + Br2(g)  CH2BrCH2Br(l) d. 2 C2H6(g) + 7 O2(g)  4 CO2(g) + 6 H2O(g)

34. An Introduction to Free Energy
To decide whether a process is spontaneous, both enthalpy and entropy changes must be considered:
Spontaneous process: Decrease in enthalpy (–DH). Increase in entropy (+DS).
Nonspontaneous process: Increase in enthalpy (+DH). Decrease in entropy (–DS).

35. An Introduction to Free Energy
Gibbs Free Energy Change (DG): Weighs the relative contributions of enthalpy and entropy to the overall spontaneity of a process.
DG = DH – TDS
DG < 0 Process is spontaneous (favorable)
DG = 0 Process is at equilibrium
DG > 0 Process is nonspontaneous (unfavorable)

36. Problems
Which of the following reactions are spontaneous under standard conditions at 25°C?
a. AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) DG° = –55.7 kJ
b. 2 C(s) + 2 H2(g)  C2H4(g) DG° = 68.1 kJ
c. N2(g) + 3 H2(g)  2 NH3(g) DH° = –92 kJ; DS° = –199 J/K

37. An Introduction to Free Energy
Equilibrium (DG° = 0): Estimate the temperature at which the following reaction will be at equilibrium. Is the reaction spontaneous at room temperature?
N2(g) + 3 H2(g)  2 NH3(g) DH° = –92.0 kJ DS° = –199 J/K
Equilibrium is the point where DG° = DH° – TDS° = 0

38. Problem
Benzene, C6H6, has an enthalpy of vaporization, DHvap, equal to 30.8 kJ/mol and boils at 80.1°C. What is the entropy of vaporization, DSvap, for benzene?

39. Optional Homework
Text , 8.32, 8.50, 8.52, 8.56, 8.58, 8.66, 8.70, 8.74, 8.82, 8.88, 8.90
Chapter 8 Homework from website

40. Required Homework
Chapter 8 Assignment