1. Thermochemistry
Unit 7

2. Energy Changes in Chemical Reactions
Heat is the transfer of thermal energy between two bodies that are at different temperatures.
Temperature is a measure of the thermal energy.
Temperature = Thermal Energy
400C
900C
greater thermal energy
A Video – Eureka!

3. Thermochemistry is the study of heat change in chemical reactions and changes of state (physical changes)
System - is the specific part of the universe that is of interest in the study.

4. 2H2 (g) + O2 (g) 2H2O (l) + energy
Exothermic process is any process that gives off heat – transfers thermal energy from the system to the surroundings.
Exothermic Phase changes – gas to liquid to solid take heat out
2H2 (g) + O2 (g) H2O (l) + energy
Endothermic process is any process in which heat has to be supplied to the system from the surroundings.
Endothermic Phase Changes – solid to liquid to gas – put heat in
energy + H2O (s) H2O (l)

5. Exothermic and Endothermic Videos Mythbusters Exothermic Reaction
Death of a Gummy Bear
The Most Common Endothermic Reaction

6. Thermochemical Equations
Enthalpy (H) is used to quantify the heat flow into or out of a system in a process that occurs at constant pressure.

7. Drawing Enthalpy Diagrams

8. Draw the energy diagram & thermochemical equation for the following reactions:
4 H O2  H2O kJ
2H2O(g)  kcal → 2H2(g) + O2(g)

9. The MATH of Thermochemical Equations
How much heat is evolved when 266 g of white phosphorus (P4) burn in air?
P4 (s) + 5O2 (g) P4O10 (s) DH = kJ
6470kJ

10. The MATH of Thermochemical Equations
How much heat is absorbed when 32.5g of O2(g) is reacted with N2 (g)?
N2(g) + 2 O2(g) ----> 2 NO2(g) ΔH = kJ

11. Introduction to Specific Heat and Heat Capacity

12. The specific heat (s) of a substance is the amount of heat (q) required to raise the temperature of one gram of the substance by one degree Celsius.
Broader Definition
The heat capacity (C) of a substance is the amount of heat (q) required to raise the temperature of a given quantity (m) of the substance by one degree Celsius. (general amount)

13. How much heat is given off when an 869 g iron bar cools from 940C to 50C?

14. Multiple Methods to Calculate ΔH
Heats of Formation
Calorimetry
Hess’s Law

15. Coffee Cup Calorimetry
Method #1 - Calorimetry
Experimental way of measuring heat generation/consumption by essentially catching all the heat energy in a water bath or water bath + metal apparatus.
Coffee Cup Calorimetry
Styrofoam cup insulates the contents so any heat generated or consumed in the water can be measured by the temperature change
q = -(Hrxn) = msH2O∆T
Reactions must take place in water, then you measure the change in temperature

16. Calorimetry
When 100ml of 0.1M HCl is mixed with 100ml of 0.1M NaOH in a coffee cup calorimeter, the temperature increases from 25oC to 29.8oC. Assume that the coffee cup is a perfect insulator and specific heat capacity for the mixture is that of water, 4.18 J/g°C. What is the enthalpy of reaction?

17. Method #2 – Heats of Formation
Standard enthalpy of formation (DH°) is the heat change that results when one mole of a compound is formed from its elements at a pressure of 1 atm. f
The standard enthalpy of formation of any element in its most stable form is zero.
DH° (O2) = 0 f

18. 6.5

19. What is the ΔH of the following reaction using Heats of Formation?
The standard enthalpy of reaction (DH° ) is the enthalpy of a reaction carried out at 1 atm.
rxn
aA + bB cC + dD
DH0
rxn
dDH° (D) f
cDH° (C) = [ + ] -
bDH° (B)
aDH° (A)
DH0
rxn
nDH° (products) f = S
mDH° (reactants) -
Example
What is the ΔH of the following reaction using Heats of Formation?
Zn AgNO3 → Zn(NO3) Ag

20. Benzene (C6H6) burns in air to produce carbon dioxide and liquid water
Benzene (C6H6) burns in air to produce carbon dioxide and liquid water. How much heat is released per mole of benzene combusted? The standard enthalpy of formation of benzene is kJ/mol.
2C6H6 (l) + 15O2 (g) CO2 (g) + 6H2O (l)

21. Method #3 - Hess’s Law
When reactants are converted to products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps.
(Enthalpy is a state function. It doesn’t matter how you get there, only where you start and end.)
6.5

22. From the following data,
CH4(g) + 2O2(g) → CO2(g) + 2H2O(l)         ΔHo = -890 kJ/mol                 H2O(l) → H2O(g)         Δ Ho = 44 kJ/mol
Calculate the enthalpy of the reaction
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g)         Δ Ho = ?

23. Calculate the standard enthalpy of formation of CS2 (l) given that:
C(graphite) + O2 (g) CO2 (g) DH0 = kJ
rxn
S(rhombic) + O2 (g) SO2 (g) DH0 = kJ
rxn
CS2(l) + 3O2 (g) CO2 (g) + 2SO2 (g) DH0 = kJ
rxn
C(graphite) + 2S(rhombic) CS2 (l)

24. Spontaneous Processes
Spontaneous: process that does occur under a specific set of conditions
Nonspontaneous: process that does not occur under a specific set of conditions
Copyright McGraw-Hill 2009

25. Spontaneous vs Nonspontaneous

26. Why would an endothermic reaction happen spontaneously?
Entropy
Entropy (S): Can be thought of as a measure of the disorder of a system
In general, greater disorder means greater entropy

27. Entropy in Action

28. Entropy Values

29. Trends in Entropy
Entropy for gas phase is greater than that of liquid or solid of same substance
I2 (g) has greater entropy than I2 (s)

30. More complex structures have greater entropy
C2H6 (g) has greater entropy than CH4 (g)

31. Spure < Saqueous

32. Slower temp < Shigher temp

33. Sfewer moles < Smore moles

34. Entropy Changes in a System Qualitative
Determine the sign of S for the following
1. Liquid nitrogen evaporates
2. Salt is dissolved in water.
3. Liquid water is heated from 22.5 C to 55.8 C
4. Water condenses on outside of bottle

35. Entropy Changes in the System
Entropy can be calculated from the table of standard values just as enthalpy change was calculated.
Srxn = nS products  mS reactants

36. Standard Entropy
2NH3(g)  N2(g) + 3H2(g)
Srxn = J/K · mol

37. Your Turn!
Calculate the standard entropy change for the following using the table of standard values. (first, predict the sign for S qualitatively)
2H2(g) + O2(g)  2H2O (g)

38. Gibbs Free Energy G = H – T S
The Gibbs free energy, expressed in terms of enthalpy and entropy can be used to predict spontaneity.

39. Gibbs Free Energy
If G < 0,negative, the forward reaction is spontaneous.
If G = 0, the reaction is at equilibrium.
If G > 0, positive, the forward reaction is nonspontaneous

40. Predicting Sign of G

41. The hydrogenation of ethene gas under standard conditions (T = 298
The hydrogenation of ethene gas under standard conditions (T = K) shows a decrease in disorder (ΔS˚ = kJ/(mol•K)) during an exothermic reaction (ΔH˚ = kJ/mol). Determine whether the reaction is spontaneous or nonspontaneous by calculating ΔG˚
C2H4 (g) + H2 (g) → C2H6 (g)

42. Calculate the standard free energy of the following reaction from standard free energies of formation.
2 CO (g) + O2 (g) → 2 CO2 (g)