1. Chemical Thermodynamics
Unit 3
Chemical Thermodynamics

2. 1st Law of Thermodynamics
Thermodynamics is the study of the changes in energy that accompany chemical and physical processes.
These energy changes usually involve heat (“thermo”)
Energy is the capacity to do work or to transfer heat
Two general types of energy:
Potential energy
Kinetic energy

3. 1st Law of Thermodynamics
Potential energy
Is stored energy or the energy a system possess by virtue of its position or composition.
E.g. the chemical energy in fuel or food comes potential energy stored in the atoms of the molecules
Kinetic energy
Is the energy of motion
Ekinetic = ½ mv2 where m = mass of the object and v is its velocity

4. The First Law of Thermodynamics
Reactions that release energy in the form of heat are called exothermic rxns.
For example, the combustion of propane is exothermic.
The combustion of butane is also exothermic.

5. 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.

6. The First Law of Thermodynamics
A process that absorbs energy from its surroundings is called endothermic
(a) When solid hydrated Ba(OH)2 and excess solid NH4NO3 are mixed, an endothermic rxn occurs (b) If the flask is placed in a wet wooden block, the water freezes and attached the block to the flask

7. 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.
Some examples of this include:
H2O flows downhill.
Objects fall when dropped.

8. The First Law of Thermodynamics
Chemical systems tend toward a state of maximum disorder.
Common examples of this are:
A mirror shatters when dropped and does not reform.
It is easy to scramble an egg and difficult to unscramble it.
Food dye when dropped into water disperses.

9. The First Law of Thermodynamics
This law can be stated as, “The total amount of energy in the universe is constant.”
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.

10. Some Thermodynamic Terms
The substances involved in the chemical and physical changes under investigation are called the system.
In chemistry lab, the system is the chemicals inside the beaker.
The environment around the system is called the surroundings.
The surroundings are outside the beaker.
The system plus the surroundings is called the universe.

11. Some Thermodynamic Terms
The set of conditions that specify all of the properties of the system is called the thermodynamic state of a system.
For example the thermodynamic state could include:
The number of moles and identity of each substance.
The physical states of each substance.(s, l, g, aq)
The temperature of the system.
The pressure of the system.

12. Some Thermodynamic Terms
A variable that defines the state of the system and independent of the pathway by which a process occurs is called a state function.
State functions are always written using capital letters
E.g. P, T, V.
E.g. consider 1 mol of pure liquid water at 300C and 1 atm. Some time later the temp is 220C and same pressure, then it is in a different thermodynamic state
There is net change in temp. however it does not matter whether
cooling took place directly (slowly or rapidly)
It was heated first and then cooled
Or any other conceivable path was followed form initial to final state.

13. Some Thermodynamic Terms
In thermodynamics we are often interested in changes in functions.
We will define the change of any function X as:
X = Xfinal – Xinitial
If X increases X > 0
If X decreases X < 0.

14. Enthalpy Change
Chemistry is commonly done in open beakers on a desk top at atmospheric pressure.
Because atmospheric pressure only changes by small increments, this is almost at constant pressure.
The enthalpy change, H, is the change in heat content at constant pressure.
H = qp
Enthalpy change also called heat change or heat of reaction

15. Enthalpy Change
Hrxn is the heat of reaction or enthalpy change of a reaction.
This quantity will tell us if the reaction produces or consumes heat.
If Hrxn < 0 the reaction is exothermic.
If Hrxn > 0 the reaction is endothermic.
Hrxn = Hproducts - Hreactants
Hrxn = Hsubstances produced - Hsubstances consumed
Notice that this is Hrxn = Hfinal – Hinitial

16. Calorimetry
This is one method to measure heat produced or absorbed in a reaction at constant P (qP) for reactions in solution.
A coffee-cup calorimeter is the device used.
Features include:
Stirring rod for thorough mixing & to ensure uniform heating
Polystyrene walls and cover for insulation and prevention of heat loss
Thermometer

17. Calorimetry
If an exothermic reaction is performed in a calorimeter, the heat evolved by the reaction is determined from the temperature rise of the solution.
This requires a two part calculation.
Amount of heat gained by calorimeter is called the heat capacity of the calorimeter or calorimeter constant.
The value of the calorimeter constant is determined by adding a specific amount of heat to calorimeter and measuring the temperature rise.

18. Calorimetry
Refer to examples in textbook, pg 554.

19. Calorimetry
Example 15-1: When kJ of heat is added to a calorimeter containing g of water the temperature rises from 24.00oC to 36.54oC. Calculate the heat capacity of the calorimeter in J/oC. The specific heat of water is J/g oC.

20. CH3COOH(aq) + NaOH(aq)  NaCH3COO(aq) + H2O()
Calorimetry
Example 15-2: A coffee-cup calorimeter is used to determine the heat of reaction for the acid-base neutralization
CH3COOH(aq) + NaOH(aq)  NaCH3COO(aq) + H2O()
When we add mL of M NaOH at oC to mL of M CH3COOH already in the calorimeter at the same temperature, the resulting temperature is observed to be oC.

21. Calorimetry
Example 15-2: The heat capacity of the calorimeter has previously been determined to be 27.8 J/0C. Assume that the specific heat of the mixture is the same as that of water, 4.18 J/g0C and that the density of the mixture is 1.02 g/mL.

22. Thermochemical Equations
Thermochemical equations are a balanced chemical reaction plus the H value for the reaction.
For example, this is a thermochemical equation.
The stoichiometric coefficients in thermochemical equations must be interpreted as numbers of moles.
1 mol of C5H12 reacts with 8 mol of O2 to produce 5 mol of CO2, 6 mol of H2O, and releasing 3523 kJ is referred to as one mole of reactions.

23. Thermochemical Equations
This is an equivalent method of writing thermochemical equations.
H < 0 designates an exothermic reaction.
H > 0 designates an endothermic reaction

24. Thermochemical Equations
∆H is directly proportional to the quantity of a substance that reacts or is produced by a reaction.
Amount of energy produced is directly proportional to mass. Therefore, if you double the coefficients in an equation, then the value of ∆H is multiplied by two. E.g.
C2H4 (g) + 3O2 (g)  2CO2 + 2H2O (l) ∆H = kJmol-1
2C2H4 (g) + 6O2 (g) 4CO2 + 4H2O (l) ∆H = kJmol-1

25. Thermochemical Equations
∆H for the reverse reaction is equal in magnitude but opposite in sign to ∆H for the forward rxn
For example:
2H2 (g) + O2 (g)  2H2O (g) ΔH= kJ/mol
2H2O (g) 2H2 (g) + O2 (g) ΔH= kJ/mol

26. Thermochemical Equations
The physical state of matter must be shown, since changes of state require energy changes

27. Thermochemical Equations
Example: Sodium nitrate dissolves readily in water according to the equation:
NaNO3(s)→ NaNO3(aq) ∆H = kJ mol–1
Determine the energy change when 1.0 g of solid NaNO3 is dissolved in water.
Would you expect the temperature to rise or fall when sodium nitrate dissolves in water? Explain your answer.

28. Thermochemical Equations
? mols (NaNO3) = 1 mol x 1.0g = mol
85g
1 mol NaNO3 absorbs 21 kJ energy
mol absorbs × 21 kJ = 0.25 kJ
The temperature will fall. Since the reaction is endothermic the enthalpy of reactants is higher than the enthalpy of products, i.e. the reaction absorbs energy

29. Standard States and Standard Enthalpy Changes
Thermochemical standard state conditions
T = 250C or 298 K.
P = atm.
Be careful not to confuse these values with STP(0 oC and 1 atm)
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.

30. Standard Molar Enthalpies of Formation, Hfo
The standard molar enthalpy of formation is defined as the enthalpy for the reaction in which one mole of a substance is formed from its constituent elements.
The symbol for standard molar enthalpy of formation is Hfo.
The standard molar enthalpy of formation for MgCl2 is:

31. Standard Molar Enthalpies of Formation, Hfo
Standard molar enthalpies of formation have been determined for many substances and are tabulated in Table 15-1 and Appendix K in the text.
Standard molar enthalpies of elements in their most stable forms at 298. K and atm are zero.
Example 15-4: The standard molar enthalpy of formation for phosphoric acid is kJ/mol. Write the equation for the reaction for whichHof = kJ.
P in standard state is P4
Phosphoric acid in standard state is H3PO4(s)
You do it!

32. Standard Molar Enthalpies of Formation, Hfo
Remember: We ALWAYS interpret coefficients in thermochemical equations as numbers of mols AND NOT molecules
Therefore 3/2 H2 (g) above refer to 3/2 mol H2 of molecules

33. Hess’s Law
Hess’s Law of Heat Summation 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

34. Hess’s Law The H for Fe2O3 is given in Equ [1] below:-
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.

35. Hess’s Law
Arrange reactions [2] and [3] so that they also have FeO and O2 as reactants and Fe2O3 as a product.
Any reaction can be doubled, tripled, or multiplied by half, etc.
However, the Ho values are also doubled, tripled, etc. accordingly
If a reaction is reversed the sign of the Ho is changed.
Reverse & double Equ 2

36. Hess’s Law Example 15-7: Given the following equations and Hovalues
calculate Ho for the reaction below.

37. Hess’s Law
The + sign of the Hovalue tells us that the reaction is endothermic.
The reverse reaction is exothermic, i.e.,