Ch. 6: Thermochemistry

1. Ch. 6: Thermochemistry Dr. Namphol Sinkaset Chem 200: General Chemistry I

2. I. Chapter Outline • Introduction • Energy • Thermochemical Terminology • Changes in Internal Energy • Reactions at Constant V or P • Heats of Reaction • Thermochemical Equations • Formation Reactions

3. I. Chemistry and Energy • We have said before that chemistry is all about energy – energy lost, gained, and transferred. • Things always seek the lowest energy state which is equivalent to seeking highest stability. • Recall the reason for bonding; now, we look at the reasons for reacting.

4. II. What is Energy? • Energy is the ability to do work. • Recall KE and PE • thermal energy: energy associated w/ temp of an object (form of KE) • chemical energy: energy associated w/ relative positions of e-’s and nuclei in atoms and molecules (form of PE) • Energy can be transferred in two ways. • work: force acting through a distance • heat: flow of energy caused by temp. difference

5. II. Units of Energy • The SI unit of energy is the joule (J), 1 J = 1 kg·m2/s2. • An older unit is the calorie (cal), defined as the energy needed to raise the temp. of 1 g H2O by 1 °C. • Note that 1 cal = 4.184 J

6. III. System and Surroundings • We need to define points of view to observe where energy goes. • system: part of universe we will focus on • surroundings: everything else

7. III. The Internal Energy • Say we have a system comprised of reactants in solution. • Every particle moves constantly and feels different attractions/repulsions. • There is KE from movement and PE from position. • The sum of KE and PE for all particles in a system is called the internal energy (E).

8. III. Components of E • The total internal energy is comprised of many components. • Many different motions for KE and many different interactions for PE.

9. III. State Functions • Total internal energy (E) is a state function. • The value of a state function depends only on the state of the system.

10. III. State Functions • The E of a system does not depend on how it got there. Same goes for T, P, V. • “Independent of path.” • “It is what it is.” • State functions usually given upper-case variables. • Changes in state functions are easily calculated by subtracting initial value from final value (final – initial).

11. III. Changes in Internal Energy • When a rxn occurs, internal energy of products is usually different than internal energy of reactants. • The change can be easily calculated by taking the difference between the two.

12. III. Energy Flow • Since energy cannot be created or destroyed, a change in energy in the system is equal and opposite to the change in energy of the surroundings.

13. III. 1st Law of Thermodynamics • You already know this law by another name: Law of Conservation of Energy. • This can be restated as the total energy of the universe is constant, i.e. there’s never a change in Euniverse.

14. IV. Calculating Changes in E • The total change in internal energy is equal to the energy transferred via heat and/or work. • The sign convention for q and w is very important!!

15. IV. E Transfer of Heat Only • Negative values of q mean heat is lost by the system. • Positive values of q mean heat is absorbed by the system.

16. IV. E Transfer of Work Only • Negative values of w mean work is done by the system. • Positive values of w mean work is done on the system.

17. IV. Measuring Energy Changes • To find ΔE, we need to measure or calculate both q and w. • Heat is related to temperature change. • The 2 types of w that are most relevant are electrical work and pressure-volume work.

18. IV. Measuring Heat • Intuitively, we know temp. is related to heat, and different substances require different amounts of heat to change their temp. • q = C × ΔT • heat capacity (C): the amount of heat needed to change the temp of a substance by 1 K • What’s wrong with this?

19. IV. Specific Heat Capacity • specific heat capacity (Cs): the amount of heat required to change the temp. of 1 g of a substance by 1 K

20. IV. Sample Problem • Find the heat transferred (in kJ) when 5.50 L of ethylene glycol in a car radiator cools from 37.0 °C to 25.0 °C if the density of ethylene glycol is 1.11 g/mL, and its specific heat capacity is 2.42 J/g·K.

21. IV. Measuring Work

22. V. Reactions at Constant V

23. V. Reactions at Constant P • At constant P, we introduce a variable called enthalpy (H) that doesn’t necessitate that we calculate or measure q and w separately. • Enthalpy is defined as H = E + PV. • Typically we are interested in ΔH…

24. V. More Useful ΔH Equation • If only P-V work can be done under constant P…

25. V. More Useful ΔH Equation • We see that enthalpy is the heat gained/lost under constant pressure conditions. • Much easier to measure ΔH than ΔE, and typically they are close in value because few reactions do P-V work.

26. V. Measuring ΔH • We can measure temp. changes of a reaction in a controlled system and relate that to heat. This experiment is known as calorimetry. • We saw a constant V system before, now it’s a constant P system.

27. V. Constant P Calorimetry • In calorimetry, we run a controlled reaction and monitor the temp. change. • The important things to remember are: qsystem= -qsurroundings qrxn= -qsoln qrxn = ΔHrxn

28. VI. Heats of Reaction • H is a state function, so ΔH is simply Hfinal – Hinitial. • The enthalpy of a reaction is known as the heat of reaction (ΔHrxn). Reactions can be exothermic or endothermic.

29. VI. Sample Problem • 50.0 mL of 1.00 M NaOH is added to 25.0 mL of 1.25 M H2SO4 in a coffee cup calorimeter. If the initial temp. is 25.00 °C, and the final temp. is 33.83 °C, what is ΔHrxn in kJ/mole? Use d = 1.00 g/mL for all solutions and c = 4.184 J/g·K for the combined solution.

30. VII. Thermochemical Equations • A thermochemical equation is a balanced equation that includes the heat of reaction. 2H2O(l) 2H2(g) + O2(g)ΔHrxn = 572 kJ

31. VII. Rules for Manipulating Thermochemical Equations • If a reaction is reversed, the sign of ΔHrxn changes. • Changing the coefficients by a constant factor also changes ΔHrxn by the same factor. • The ΔHrxn can be used in a stoichiometric ratio.

32. VII. Hess’s Law • Some reactions are impossible or very difficult to carry out, yet we can still calculate what ΔH would be. • Hess’s Law: the enthalpy change of an overall process is the sum of the enthalpy changes of its individual steps • This means that you can break up a reaction into a series of smaller reactions that add up to the big reaction, and the ΔHrxn of the big reaction will be the sum of all the small reaction ΔHrxn’s.

33. VII. Sample Problem • Calculate ΔHrxn for the reaction 2NO2(g) + ½ O2(g) N2O5(g) given the information below. N2O5(g) 2NO(g) + 3/2 O2(g)ΔH = 223.7 kJ NO(g) + 1/2 O2(g) NO2(g)ΔH = -57.1 kJ

34. VII. Sample Problem • Calculate ΔHrxn for the reaction Ca(s) + ½ O2(g) + CO2(g) CaCO3(s) using the information below. Ca(s) + ½ O2(g) CaO(s)ΔH = -635.1 kJ CaCO3(s) CaO(s) + CO2(g)ΔH = 178.3 kJ

35. VII. Standard Heats of Reaction • Thermodynamic variables vary a little with conditions, so we need to define a set of standard conditions. • Standard conditions are defined at 1 atm and 25 °C. • Products and reactants at these conditions are said to be in their standard states. Note that solutions must be 1 M. • The “not” symbol is used to indicate standard conditions, e.g. ΔH°rxn.

36. VIII. Formation Reactions • A formation reaction is one that creates exactly 1 mole of a substance from its elements. • Of course, standard heats of formation occur at 1 atm and 25 °C. Na(s) + ½ Cl2(g) NaCl(s)ΔH°f = -411.1 kJ

37. VIII. Using ΔH°f’s • Standard heats of formation are very powerful; they allow calculation of ΔH°rxn for anything using a Hess’s Law type of calculation. ΔH°rxn = (sum ΔH°f products) – (sum ΔH°f reactants)

38. VIII. Sample Problem • Calculate ΔH°rxn for C2H2(g) + 5/2 O2(g) 2CO2(g) + H2O(g) given that the standard heats of formation for acetylene, carbon dioxide, and steam are 227.0 kJ/mole, -393.5 kJ/mole, and –241.8 kJ/mole, respectively.