1. Thermochemistry Chemistry that deals with Energy Change
First Law of Thermodynamics
Heat (q), work (w) and internal energy (E)
Calculation of heat gained or lost by system
Calorimeter for measuring heat exchange (between system and surrounding;
Enthalpy Change and the State Function
Enthalpy of Formation and Hess’s law of summation

2. Energy Basics
The energy changes that accompany chemical reactions and physical changes are very important to us.
The reactions involved in metabolism of food.
The reactions involved in generation of electricity.
The reactions involved in the powering of our automobiles.
Sliding a match head along a rough surface initiates a combustion reaction that produces energy in the form of heat and light. (credit: modification of work by Laszlo Ilyes)

3. Energy Energy is the capacity to supply heat or do work.
There are two forms of energy: Kinetic & Potential
Potential energy is the energy possessed by an object by virtue of its position, composition, or condition.
Kinetic energy (Ek) is the energy that an object posses because of its motion.

4. Potential & Kinetic Energy
Potential Energy: Ep = -mgh
Energy associated with the relative position of a substance in a force field, such as gravity, magnetic, chemical bonds, nuclear force, etc.
(m = mass; g = gravitational constant, and h = height)
Kinetic Energy: EK = ½mv2
Energy associated with the translational motion of an object. Kinetic energy depends on the mass and speed of the object. (m = mass; v = velocity or speed)

5. Potential Energy
Potential energy is the energy possessed by an object by virtue of its position, composition, or condition.
One important form of potential energy of interest to chemists is Chemical energy.
Chemical energy is energy stored within the structural units or chemical bonds in a substance.
If the water flows through generators at the bottom of a dam, such as the Hoover Dam shown here, its kinetic energy is converted into electrical energy. (credit a: modification of work by Steve Jurvetson; credit b: modification of work by “curimedia”/Wikimedia commons)

6. Energy and Energy Changes
Energy is interconvertible - One form can be converted to another form.
Law of conservation of energy – During a chemical or physical change, energy can be neither created nor destroyed, although its form can change.
Although energy can assume many forms, the total energy of the universe is constant.
Water that is higher in elevation, for example, at the top of Victoria Falls, has a higher potential energy than water at a lower elevation. As the water falls, some of its potential energy is converted into kinetic energy. (credit a: modification of work by Steve Jurvetson; credit b: modification of work by “curimedia”/Wikimedia commons)

7. Thermal Energy
One form of kinetic energy of particular interest to chemists is called thermal energy or heat;
This is the energy associated with the random motion of atoms and molecules in the sample.

8. Thermal Energy and Temperature
Thermal energy – Kinetic energy associated with the random motion of atoms and molecules.
Temperature – A quantitative measure of “Hot” and “Cold”.
Fast moving molecules  High thermal energy  “Hot”
Slow moving molecules  Low thermal energy  “Cold”
Heat (q) – The transfer of thermal energy between two bodies at different temperature.
Thermochemistry – The study of the heat flow during a chemical or physical change.

9. What do we know about heat?
Heat will flow from a hotter object to a colder object.
a) Substances H and L are initially at different temperatures, and their atoms have different average kinetic energies.
b) When they are put into contact with each other, collisions between the molecules result in the transfer of kinetic (thermal)energy from the hotter to the cooler matter.
c) The two objects reach “thermal equilibrium” when both substances are at the same temperature, and their molecules have the same average kinetic energy.

10. First Law of Thermodynamics
Energy is neither created nor destroyed during chemical or physical processes;
The energy change (DE) in a system depends only on heat (q) gained or lost and the work (w) done on or by the system:
DE = q + w
For processes involving gas expansion: w = -pDV
DE = q – pDV; q = DE + pDV = DH

11. Work Associated with Gas Expansion

12. Principles of Heat Flow
Definitions:
The System: The substance or substances undergoing the chemical or physical change.
The Surroundings: The rest of the universe.
We are usually only concerned with the surroundings that are in direct contact with the system.
We will discuss the flow of heat between the system and surroundings.

13. Heat Flow During Chemical Reactions
Direction of Heat flow
Matter undergoing a chemical or physical process can absorb or release heat.
Endothermic process – A reaction or change that absorbs heat.
Heat flows from the surroundings to the system.
The sign on q is positive for an endothermic process.
Exothermic process – A reaction or change that releases heat.
Heat flows from the the system to the surroundings.
The sign on q is negative for an exothermic process.

14. Calorimetry
A calorimeter is a device used to measure the amount of heat and direction of heat flow in a chemical or physical process.
The calorimeter contains water and/or other materials of known heat capacity.
The walls of the calorimeter are insulated so all of the heat flow is only between the process and the calorimeter components.
System = Substance(s) undergoing the chemical or physical change.
Surroundings = Calorimeter components (water, etc.)
When thermal equilibrium is reached the system and surroundings will be at the same final temperature.
The heat gained (or lost) by system is equal in magnitude and opposite in sign to that lost (or gained) by calorimeter.
qreaction = - qcalorimeter

15. Calorimeter for measuring Heat Exchanges
Coffee – cup calorimeter
Made out of two styrofoam coffee cups, filled partially with water.
Works well for reactions in aqueous solution.
Not suitable for reactions involving gases.
A simple calorimeter can be constructed from two polystyrene cups. A thermometer and stirrer extend through the cover into the reaction mixture.

16. Calculation of Heat Gained or Lost
Heat gained or lost: q = m.s.Dt
where m = mass; s = specific heat capacity of the substance, and Dt = change in temperature
Enthalpy change (DH):
Heat gained or lost during a chemical reaction at constant pressure.

17. Specific Heat Capacity Determination
To determine the specific heat of a metal using coffee cup calorimeter:
55.0 g of hot metal at 99.5oC is added to 40.0 g of cold water at 22.5 oC in a Styrofoam cup calorimeter. When metal and water reach thermal equilibrium, the final temperature is 32.4oC. (a) Calculate the amount of heat absorbed by water (and lost by metal); (b) Calculate the specific heat of metal.
(Specific heat of water = J/(g.oC); ignore heat absorbed by calorimeter.)
(Answer: (a) 1.6 x 103 J; (b) 0.45 J/g.oC)

18. Heat Capacity of Calorimeter
Heat capacity is the amount of heat absorbed by an object when its temperature increases by 1.00oC.
A 50.0-g sample of warm water at 40.0oC is added to a 50.0 g of cold water in a Styrofoam cup calorimeter; the initial temperature of cold water is 20.0oC. If the final temperature of mixture is 29.6oC, calculate the heat absorbed by calorimeter and the heat capacity, Ccal, of the calorimeter.
(Specific heat of water is J/g.oC)
(Answer: qcal = 170 J; Ccal = 17 J/oC)

19. Measuring enthalpy of dissolution of an ionic compound using calorimeter.
When 7.50 g of ammonium nitrate, NH4NO3, is dissolved in g of water at 25.0 oC in a calorimeter, the temperature of the solution decreases to 19.9 oC. The specific heat of water is J/g.oC, and the heat capacity of calorimeter is 15 J/oC.
Calculate the value of q for the dissolution of NH4NO3 sample. What is the algebraic sign of q for the process? Is the dissolution of ammonium nitrate endothermic or exothermic?
Calculate the molar enthalpy of ammonium nitrate dissolution.
(Answer: (a) q = +2.4 kJ; positive ”+”; endothermic; (b) DH = 25 kJ/mol)

20. Determination of Reaction Enthalpy
The enthalpy of a reaction in aqueous solution can be determined using coffee cup calorimeter.
When g sample of magnesium is added to 100. mL of 1.0 M HCl solution in a Styrofoam coffee cup calorimeter, the temperature of the solution increases from 22.5oC to 34.2 oC.
(a) Calculate the total amount of heat absorbed by solution and calorimeter.
(b) What is q for the following reaction for sample used?
Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
(c) Calculate the molar enthalpy for the above reaction (DH, in kJ/mol).
(Assume the specific heat of solution = 4.18 J/g.oC; heat capacity of calorimeter = 10. J/oC, and the density of HCl solution = 1.0 g/mL)

21. Heat of Neutralization
Determination of molar enthalpy of acid-base reaction using coffee cup calorimeter.
When 50.0 mL of 1.0 M HCl is reacted with 50.0 mL of 1.0 M NaOH in a coffee cup calorimeter, the temperature of the solution increases from 22.0oC to 28.8oC. (a) Calculate the total amount of heat absorbed by the solution and calorimeter; (b) What is q for this reaction? (c) Calculate the molar enthalpy for the following reaction (DH, in kJ/mol).
HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
Assume density of solution = 1.0 g/mL; specific heat of solution = 4.18 J/g.oC, and the heat capacity of calorimeter = 10.J/oC.
(Answer: (a) 2.9 kJ; (b) q = -2.9 kJ; (c) DH = -58 kJ/mol)

22. Bomb Calorimeter for determination of Heat of Combustion

23. Determination of Heat of Combustion
To determine the enthalpy of combustion of glucose.
2.250-g of glucose (C6H12O6) is completely combusted in a “bomb” calorimeter, the temperature of calorimeter increases from oC to oC. The calorimeter contains 775 g of water (specific heat = J/g.oC), and the “bomb” has a heat capacity of 893 J/oC.
(a) How much heat is produced by the combustion of the glucose sample? (b) What is the heat of combustion for 1.00 g of glucose? (c) Calculate the molar enthalpy of combustion of glucose (DH, in kJ/mol).
(Answer: (a) kJ; (b) kJ/g; (c) kJ/mol)

24. Energy Change during Reactions: Exothermic & Endothermic Reactions

25. Enthalpy of Formation

26. Enthalpy Diagram of a Reaction
Reaction: B2H6(g) + 3O2(g)  B2O3(s) + 3H2O(g)
B2H6(g) ___________ DHf = 36 kJ
2B(s) + 3H2(g) + 3O2(g) __________ DHf = 0.0 kJ
3H2O(g) ____________ DHf = -726 kJ
B2O3(s) ____________ DHf = kJ

27. Enthalpy Diagram Reaction: C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)
3C(s) + 4H2(g) + 5O2(g) ____________ DHf = 0.0 kJ
C3H8(g) ____________ DHf = -104 kJ
4H2O(g) ____________ DHf = -968 kJ
3CO2(g) ____________ DHf = kJ

28. Hess’s Law of Enthalpy of Reactions
According to Hess’s law:
The net enthalpy change of a given process is independent of the number of steps taken to complete the process.
If a reaction can be broken down into several steps, then the overall enthalpy for the reaction is equal to the sum of enthalpies of individual steps.

29. Hess’s Law of Enthalpy of Summation
According to Hess’s law:
For a given reaction, the overall enthalpy change is equal to the difference between the algebraic sum of enthalpy of formation of products and the algebraic sum of enthalpy of formation of reactants.
DHrxn = S(npDHf[products]) – S(nrDHf[reactants])

30. Enthalpy Diagram for Formation of CO & CO2
Examples:
C(s) + O2(g)  CO2(g); DH = -394 kJ
C(s) + ½O2(g)  CO(g); DH = -283 kJ
C(s) + ½O2(g) ____________ DHf = 0.0 kJ
CO(g) + ½O2(g) ____________ DHf = -111 kJ
CO2(g) ____________ DHf = -394 kJ

31. Enthalpy Diagram for the Combustion of Methane

32. Enthalpy Diagram for Combustion Propane

33. Using Hess’s law to calculate Enthalpy of Reactions
Applying Hess’s law to calculate enthalpy of reactions.
Given:
C(s) + O2(g)  CO2(g); DHo = -394 kJ (1)
CO(g) + ½O2(g)  CO2(g); DHo = -283 kJ (2)
Calculate DH for the following reaction:
C(s) + ½O2(g)  CO(g)

34. Using Enthalpy of Formation to calculate Heat of Combustion
Calculating enthalpy of combustion of diborane, B2H6, using the molar enthalpy of formation data:
Given:
(1) 4B(s) + 3O2(g)  2B2O3(s); DHf(1) = kJ
(2) 2B(s) + 3H2(g)  B2H6(g); DHf(2) = 36.4 kJ
(3) 2H2(g) + O2(g)  2H2O(g); DHf(3) = kJ
Calculate DH for the following combustion reaction:
B2H6(g) + 3O2(g)  B2O3(s) + 3H2O(g);

35. Using Enthalpy of Formation to calculate Heat of Combustion
To calculate heat of reaction using molar enthalpy of formation:
Given: DHf[C3H8(g)] = -104 kJ/mol;
DHf[H2O(g)] = -242 kJ/mol, and
DHf[CO2(g)] = -394 kJ/mol;
Calculate DH for the following combustion of C3H8(g):
C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(g)

36. Calculating Enthalpy of Reaction in Aqueous Solution using the enthalpy of formation of Ions
Enthalpy change for reactions in aqueous solution:
Standard condition for gas: P = 1 atm at 25oC
Standard condition for solution: 1 M at 25oC
Under standard condition: DHf[H3O+(aq)] = 0.0 kJ
Given the following thermodynamic values:
DHf[Mg2+(aq)] = -467 kJ/mol;
DHf[OH-(aq)] = -230 kJ/mol, and DHf[Mg(OH)2(s)] = -925 kJ/mol,
Calculate DH for the reaction:
MgCl2(aq) + 2NaOH(aq)  Mg(OH)2(s) + 2NaCl(aq)

37. Enthalpy of Ionic Reactions
Given the following enthalpy of formation (in kJ/mol):
DHf[Ba2+(aq)] = ; DHf[CO32-(aq)] = ;
DHf[BaCO3(s)] = -1219; DHf[BaSO4(s)] = -1465;
DHf[OH-(aq)] = -230; DHf[H2SO4(aq)] =
DHf[H2O(l)] = -286 kJ/mol
Calculate enthalpy changes for the following reactions in aqueous solution:
(1) BaCl2(aq) + Na2CO3(aq)  BaCO3(s) + 2NaCl(aq)
(2) HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
(3) Ba(OH)2(aq) + H2SO4(aq)  BaSO4(s) + 2H2O(l)

38. Global Energy Resources
Biomass (mainly wood) – major sources of energy in many under-developed countries;
Coal was once the major source of energy in U.S.A. and industrialized European countries;
Petroleum replaces coal in the middle of 20th Century as the major source of energy for power plants and transportation;
Hydroelectric power and nuclear energy are used in certain developed countries. Geothermal energy is used as a secondary source of energy
Solar energy is a secondary source of energy, but mainly for household heating.

39. Comparison of Enthalpy of Combustion
————————————————————————
Substances Energy (kJ/g fuel)
Hydrogen gas (H2) 120
Natural gas (CH4) 50
Gasoline
Crude petroleum 43
Animal fat
Coal
Charcoal
Ethanol
Methanol
Paper
Dry biomass (wood)

40. Petroleum and Natural Gas
Origin of petroleum and natural gas:
most likely from fossilized remains of marine organisms that lived approximately 500 millions years ago
Petroleum –
thick, dark liquid composed of mixture of hydrocarbons
Composition varies from one location to another, but mostly hydrocarbon compounds containing C5 to > C25
Natural Gas
Consists mostly methane (CH4, >90%) and some ethane (C2H6), propane (C3H8) and butane (C4H10)

41. Petroleum Refining Petroleum refining
Fractional distillation of crude petroleum yields the following fractions:
Gasoline (C5 – C10);
Kerosene & jet fuel (C10 – C18);
Diesel fuel, heating, and lubricating oil (C15 – C25),
Asphalt (>C25)
More gasoline is produced by pyrolytic (high temperature) cracking of larger HC compounds (> C25)

42. Coal
Formed from fossilized plant remains that have been subjected to high temperature and pressure for many millions years
Coal matures through 4 stages:
Lignite, subbituminous, bituminous, and anthracite;
Composition by mass%:
Lignite: 71% C, 4% H, 23% O, 1% N, and 1% S;
Subbituminous: 77% C, 5% H, 16% O, 1% N, and 1% S;
Bituminous: 80% C, 6% H, 8% O, 1% N, and 5% S;
Anthracite: 92% C, 3% H, 3% O, 1% N, and 1% S.
The relative carbon content increases and those of hydrogen and oxygen decrease as coal matures.

43. Coal as Energy Source
Coal furnishes about 23% of energy needs in U.S.A.
Underground coal mining is dangerous
Strip mining destroys lands and the environments
Coal contains sulfur and burning coal causes severe air pollution due to:
Air particulate matters, CO2, CO, and SO2 (from sulfur in coal)
In atmosphere, SO2 is oxidized to SO3, which yields acid rain (SO3 forms H2SO4 when mixed with rain water)

44. Processing Coal C(s) + ½ O2(g)  CO(g); DH = -111 kJ
Coal gasification - converting coal into gaseous fuel
Treating coal with air and steam at high temperature produces mixture of CO, H2, and CH4. Some CO2 and SO2 are also formed
Mixture of CO and H2 is also called syngas
Reactions in coal gasification:
C(s) + O2(g)  CO2(g); DH = -394 kJ
C(s) + ½ O2(g)  CO(g); DH = -111 kJ
C(s) + H2O(g)  CO(g) + H2(g); DH = 131 kJ
C(s) + 2H2(g)  CH4(g); DH = -75 kJ
Note: both exothermic and endothermic reactions occur. An energy balance can be maintained by controlling the temperature, the rate of coal feed, and the flow of air and steam.

45. Coal Gasification Process
Coal + Steam + Air
Heat
CH4 + CO, CO2, H2, H2O
(+ sulfur containing impurities)
Separate
Treatment
CO + H2O  CO2 + H2
CO + 3H2  CH4 + H2O
After removal of CO2 and H2O, the remaining mixture contains CH4 and syngas (CO + H2)

46. Coal to Coal Slurry
Coal is pulverized and mixed with water to form a thick slurry
Coal slurry burns like residual oil, with less CO and SO2 produced

47. Hydrogen Fuel Combustion of H2: H (g) + ½O2(g)  H2O(l); DH = -286 kJ
H2(g) + ½O2(g)  H2O(g); DH = -242 kJ
Combustion of H2 produces 2.5 times more energy per gram than natural gas
Combustion of H2 only produces water.
However, production, storage and transportation of the gas pose major problems

48. Problems in Hydrogen Production
H2 does not exist in the free form (like N2 or O2)
All existing methods to produce H2 gas involve endothermic reactions or high temperature:
Steam reformation:
CH4(g) + H2O(g)  CO(g) + 3H2(g); DH = 206 kJ
Electrolysis: H2O(l)  H2(g) + ½O2(g); DH = 286 kJ
Thermochemical decomposition of H2O:
2HI(g)  H2(g) + I2(s); (425oC)
2H2O(l) + SO2(g) + I2(s)  H2SO4(aq) + 2HI(aq); (90oC)
H2SO4(aq)  SO2(g) + H2O(l) + ½O2(g); (825oC)
Net Reaction: H2O(l)  H2(g) + ½O2(g);

49. Problems in Storage & Transportation of H2
H2 decomposes to H-atoms on metal surface
H-atoms are small enough to be absorbed into the metal lattices and makes the metal to become brittle
H2 requires a much larger container for storage and transportation
Storage or transportation of H2 under high pressure (say as liquefied gas) poses high explosion hazard
Energy produced per unit volume by H2 is only one-third that produced by natural gas under similar conditions
An alternative method suggested for H2 storage/transportation is to convert it into solid metal hydrides, MH2.

50. Alternative Fuels Hydrogen gas is most efficient and clean fuel
It produces the most energy per gram,
Transportation and storage are difficult
It requires a much large fuel tank than gasoline;
May be compressed into liquid form, but poses explosion hazard
An alternative way is to convert it into a solid metal hydride (MH2) – one that is able to release H2 when heated.
Most likely the hydrogen will be used to power fuel cells that will be installed in automobiles

51. Alternative Fuels Oil shale
Shale rocks contain complex HC called kerogen
Huge deposits in western states, especially Colorado;
Kerogen is not fluid like petroleum - cannot be pumped;
Rocks containing fuel must be heated to >250oC to decompose the kerogen into smaller HC molecules;
Process produces large quantities of waste rocks – a negative environmental impact.

52. Alternative Fuels Ethanol:
Produced by fermentation of sugar (sugar cane or corn)
Pure ethanol produces about 27 kJ/g of energy
Use as fuel supplement
Added to gasoline (gasohol contains ca. 10% ethanol)
Pure ethanol not suitable for motor fuel in USA, especially in winter, because it does not vaporize easily in cold climate.
Pure ethanol is widely used as motor fuel in Brazil where the climate is warmer.

53. Alternative Fuels Methanol:
Methanol produces about 20 kJ/g of energy when completely combusted
It has been used in race cars as a mixture of 85% methanol and 15% gasoline.
California is evaluating methanol as motor fuel
Arizona and Colorado are considering similar step.