1. Energy Relationships in Chemistry
Chemistry 100 Chapter 5
Energy Relationships in Chemistry

2. Thermochemistry
Thermodynamics – the study of energy and its transformations.
Thermochemical changes – energy changes associated with chemical reactions.
System ® that specific part of the universe of interest to us.
Surroundings ® the part of the universe not contained in the system.

3. 3 types of Systems open system ® exchanges mass and energy
closed system ® exchanges energy but no mass
isolated system ® no exchange of either mass or energy

4. Three Types of Systems insulation Isolated System Closed System cork
Open system

5. Different Types of Energy
Energy – the ability to do work.
Thermal energy – associated with the random motions of atoms and molecules
Heat energy – transfer of thermal energy between two objects at different temperature.

6. Energy (cont’d)
Chemical energy – energy stored within the structural units of chemical substance.
Potential energy – the ability of an object to do work because of its position in a field of force.

7. Kinetic Energy – the work that can be performed by a moving object.
The unit of energy
1 Joule (J)
=1 kg m2/s2
An older unit of energy
1 calorie (cal)
= J exactly

8. The Law of Conservation of Energy
Energy is neither created nor destroyed in ordinary chemical and physical processes
Converted from one type into another.

9. This is also stated in terms of the first law of thermodynamics.
E = internal energy change of the system
Ef and Ei  the energy of the final and initial states, respectively

10. First Law of Thermodynamics
Chemical reactions either absorb or release energy.
Two terms
Exothermic reaction ® heat is released to the surroundings.
Endothermic reaction ® heat is supplied to the system by the surroundings.

11. Exothermic

12. Endothermic

13. The First Law Restated
chemical systems – examine the conversion of heat energy into work.

14. Signs for Heat and Work Work done by system on surroundings
Work done by surroundings on system
w ‘+’
q < 0, heat flows to surroundings
Exothermic ‘-’
q > 0, heat flows to system
Endothermic ‘+’

15. Pressure-volume Work Pressure – volume work
w = -Pop V = -Pop (Vf -Vi)
This is the type of work done by the pistons in our automobile engines!
The greater the magnitude of Pop, the gas has to "work harder" to obtain the same volume change.

16. Pressure-Volume Work

17. State and Path Functions
E, H, V are examples of state functions.
State functions – numerical value doesn’t depend on how the process is carried out.
Work (w) and q (heat) are path functions
The amount of work done or heat released depends on how the system changes states.

18. E = E[CO2 (g)] – E[C(s)] – E[O2(g)]
Examine a chemical reaction.
C (s) + O2 (g)  CO2 (g)
E = E[CO2 (g)] – E[C(s)] – E[O2(g)]
This reaction has a negative enthalpy change (H = kJ).

19. From the first law
surrE + sysE = 0
surrE = -sysE
The energy "lost" from the system is "gained" in the surroundings.

20. Enthalpies of Formation – Standard Reaction Enthalpies
The enthalpy change for the reaction
rH = H(products) - H(reactants)
We cannot measure the absolute values of the enthalpies!!
How do we ‘measure’ enthalpies (or heat contents) of chemical species?

21. The Formation Reaction
A "chemical thermodynamic reference point."
For CO and CO2
C (s) + O2 (g)  CO2 (g)
C (s) + ½ O2 (g)  CO (g)
The "formation" of CO and CO2 from its constituent elements in their standard states under standard conditions.

22. The Formation Reaction
For the formation of 1.00 mole of Na2SO3(s)
2 Na(s) + S(s) + 3/2 O2 (g)  Na2SO3 (s)
The ‘formation enthalpy of Na2SO3(s)’,
symbolised fH[Na2SO3 (s)]

23. Standard Conditions for Thermodynamic Reactions
The degree sign, either  or , indicates standard conditions
P = 1.00 atm
[aqueous species] = 1.00 mol/L
T = temperature of interest (note 25C or 298 K is used in the tables in your text).

24. The Significance of the Formation Enthalpy
fH° is a measurable quantity!
Compare CO (g) with CO2 (g)
C (s) + 1/2 O2 (g)  CO (g)
fH° [CO(g)] = kJ/mole
C (s) + O2 (g)  CO2 (g)
fH° [CO2(g)] = kJ/mole
The formation enthalpy for CO2(g) is larger than the formation enthalpy of CO (g).

25. Reactions Enthalpies
Formation enthalpies – thermodynamic reference point,
Formation of the elements from themselves is a null reaction – fH (elements) = 0 kJ / mole.

26. The Combustion of Propane

27. The General Equation
Calculate enthalpy changes from the formation enthalpies as follows.
Reverse a reaction, the sign of the enthalpy change for the reaction is reversed.
Multiply a reaction by an integer, the enthalpy change is multiplied by the same integer.

28. The Measurement of Energy Changes – Calorimetry
Calorimetry – the measurement of heat and energy changes in chemical and physical processes.
Heat capacity (C) – the amount of heat (energy) needed to raise the temperature of a given mass of substance by 1°C.
Specific heat capacity (s) – the amount of heat energy (in Joules, J) required to raise 1 g of a substance by 1°C (units = J/g °C).

29. General expression for heat capacity C = m s
m is the mass of the substance (in grams).
Molar heat capacity
Cm = M s
M – molar mass of the substance
s – its specific heat capacity.

30. The Calorimeter
A calorimeter – a device which contains water and/or another substance with a known capacity for absorbing energy (heat).
Calorimeters are adiabatic systems.
All energy changes take place within the calorimeter.

31. Adiabatic System
Adiabatic system – thermally insulated from the rest of the universe
No heat exchange between system and surroundings!
For an adiabatic system,
qtot = qrxn + qH2O + qcal = 0
-qrxn = qH2O + qcal

32. The Constant Volume (Bomb) Calorimeter
E = qv

33. The Constant Pressure Calorimeter
H = qp

34. Relating the Enthalpy to the Internal Energy
The enthalpy and the internal energy both represent quantities of heat.
E = qv.
H = qp.
E and H are related as follows
H = E +Pop V
V = the volume change for the reaction.

35. For reactions involving gases V = ng /(RT Pop)
ng =  np (g) -  nr (g)
For most reactions, ng is small.
The difference between the internal energy change and the enthalpy change is small.

36. Other important Enthalpy changes
Many other important processes have associated enthalpy changes.
The measurement of the heat changes for these process can give us some insight into the changes in intermolecular forces that occur during the transformation.

37. Heat of dilution and solution.
solH = the heat absorbed or given off when a quantity of solute is dissolved in a solvent.
solH = H(sol’n) - H(component)
H(component) = H (solid) + H(solvent)

38. dilH = H(sol’n 2) – H(sol’n ,1)
For the process,
HCl (aq, 6 M)  HCl (aq, 1 M).
A significant amount of heat is released when the acid solution is diluted.
This is the enthalpy of dilution of the acid.
dilH = H(sol’n 2) – H(sol’n ,1)

39. Lattice Enthalpies Look at the following process.
NaCl (s)  Na+ (g) + Cl- (g)
H = latH = 788 kJ/mole  the lattice enthalpy
A very endothermic reaction!
Due to the strength of the ionic bond!

40. Latent Heats
Latent heats are the enthalpy changes associated with phase transitions.
H2O (l)  H2O (g)
rH = vapH  the enthalpy of vapourization.
H2O (s)  H2O (l)
rH = fusH  the enthalpy of fusion.
H2O (s)  H2O (g)
rH = subH  the enthalpy of sublimation.

41. Foods and Fuels
Most of the chemical reactions that produce heat are combustion reactions.
Note – all combustion reactions are exothermic.
Fuel values are generally reported as positive quantities.
Obtaining fuel values – calorimetry.

42. Calories, Food Calories, and Kilojoules
When we read our cereal boxes, we may see the following
1 bowl cereal = 30 g cereal = 132 Cal (490 kJ).
Isn’t 1 calorie = J (not kJ)?
The fuel values of foods are reported as food calories (Cal).
1.00 food calorie (Cal) = 1000 thermal calories (cal) = 4184 J = kJ.

43. Combustion of Carbohydrates and Fats
Most of the energy our body needs comes form the combustion of sugar and fats.
For the glucose (blood sugar) combustion
C6H12O6 (s) + 6 O2 (g)  6 CO2 (g) + 6 H2O (l)
rH = kJ
This energy is supplied quickly to the body!
Average fuel value of carbohydrates = 17 kJ/g.

44. C57H110O6 (s) + 163/2 O2 (g)  57 CO2 (g) + 55 H2O (l).
Fats
The combustion (metabolism) of fats also produces CO2 and H2O.
The combustion of tristearin
C57H110O6 (s) /2 O2 (g)  57 CO2 (g) + 55 H2O (l).
rH = x 104 kJ

45. Fuel Value of Fats Fats are the body’s ‘energy stockpiles!’
Insoluble in water.
Average fuel value = 38 kJ/g – about twice that of the carbohydrates.

46. Caloric Contents
For proteins – average fuel value = 17 kJ/g, about the same value as for the carbohydrates.
The relative amounts of proteins, fats, and carbohydrates in foods determines the caloric content.

47. Fossil Fuels
Coal, petroleum, and natural gas are known as fossil fuels. They are collectively the major source of energy for commercial and personal consumption.
Fossil fuels are mixtures of many different kinds of organic compounds.
The fuel values of fossil fuels is directly related to the amount of carbon and hydrogen in the fuel.

48. Hydrogen As a Fuel Hydrogen has a huge fuel value (142 kJ/g).
The combustion product is innocuous – water.
Obviously, there are problems!
Two major difficulties with H2 as a fuel source.
Where do we get the hydrogen?
How do we store the hydrogen?