Chapter 5 Thermochemistry

1. Chemistry, The Central Science, 10th edition Theodore L. Brown; H. Eugene LeMay, Jr.; and Bruce E. Bursten Chapter 5Thermochemistry

2. Energy • The ability to do work or transfer heat • Work: the amount of energy transferred by a force acting through a distance in the direction of the force • Heat: Energy used to cause the temperature of an object to rise Energy released from chemical reactions can be used to “do work or produce heat” – examples? How would knowing about energy release help us predict the direction of a chemical reaction?

3. Potential Energy Energy an object possesses by virtue of its position or chemical composition

4. 1 KE =  mv2 2 Kinetic Energy Energy an object possesses by virtue of its motion.

5. kg m2 1 J = 1  s2 Units of Energy • The SI unit of energy is the joule (J) • An older, non-SI unit is still in widespread use: The calorie (cal) 1 cal = 4.184 J • 1 Calorie = 1000 calories. Used in food labeling

6. System and Surroundings • The system includes the molecules we want to study (here, the hydrogen and oxygen molecules) • The surroundings are everything else (here, the cylinder and piston)

7. First Law of Thermodynamics • Energy is neither created nor destroyed • In other words, the total energy of the universe is a constant; if the system loses energy, it must be gained by the surroundings, and vice versa

8. Exchange of Heat between System and Surroundings • When heat is absorbed by the system from the surroundings, the process is endothermic.

9. Exchange of Heat between System and Surroundings • When heat is absorbed by the system from the surroundings, the process is endothermic. • When heat is released by the system to the surroundings, the process is exothermic.

10. State Functions Usually we have no way of knowing the internal energy of a system; finding that value is simply too complex a problem

11. State Functions • However, we do know that the internal energy of a system is independent of the path by which the system achieved that state. • In the system below, the water could have reached room temperature from either direction.

12. State Functions • Therefore, internal energy is a state function. • It depends only on the present state of the system, not on the path by which the system arrived at that state. • And so, E depends only on Einitial and Efinal.

13. Enthalpy, H • Enthalpy is the internal energy of a body plus the product of pressure and volume: H = E + PV • A change in enthalpy is a measure of the heat gained or lost (at constant pressure) from a system H = q

14. Endothermicity and Exothermicity • A process is endothermic, then, when H is positive.

15. Endothermicity and Exothermicity • A process is endothermic when H is positive • A process is exothermic when H is negative.

16. Enthalpies of Reaction The change in enthalpy, H, is the enthalpy of the products minus the enthalpy of the reactants: H = Hproducts−Hreactants ∆H/mol = ∆Hrxn = change in enthalpy “per mole” of reaction as written

17. Enthalpies of Reaction This quantity, H, is called the enthalpy of reaction, or the heat of reaction. Units are kJ and refer to the stated chemical equation

18. Facts about Enthalpy • Enthalpy is an extensive property: changes in enthalpy (H) depend on reaction quantities • H for a reaction in the forward direction is equal in size, but opposite in sign, to H for the reverse reaction • H for a reaction depends on the state of the products and the state of the reactants

19. Calorimetry Since we cannot know the exact enthalpy of the reactants and products, we measure H through calorimetry, the measurement of heat flow. Usually the heat flow is determined by measuring the rise in temperature of water surrounding the system

20. Constant Pressure Calorimetry By carrying out a reaction in aqueous solution in a simple calorimeter such as this one, one can indirectly measure the heat change for the system by measuring the heat change for the water in the calorimeter

21. Heat Capacity and Specific Heat • The amount of energy required to raise the temperature of a substance by 1 K (1C) is its heat capacity • We define specific heat capacity (or simply specific heat) as the amount of energy required to raise the temperature of 1 g of a substance by 1 K

22. heat transferred Specific heat = mass  temperature change q s = m T Heat Capacity and Specific Heat Specific heat, then, is Usual units: J/g.K or, q = s x m xT

23. Constant Pressure Calorimetry Because the specific heat for water is well known (4.184 J/mol-K), we can measure H for the reaction with this equation: q = m  s  T Assumptions: use water specific heat, mass is mass of ?

24. Bomb Calorimetry Reactions can be carried out in a sealed “bomb,” such as this one, and measure the heat absorbed by the water

25. Bomb Calorimetry • Because the volume in the bomb calorimeter is constant, what is measured is really the change in internal energy, E, not H • For most reactions, the difference is very small Bomb Calorimetry

26. Hess’s Law • H is well known for many reactions, and it is inconvenient to measure H for every reaction in which we are interested • However, we can estimate H using H values that are published for known reactions and the properties of enthalpy

27. Hess’s Law Hess’s law states that “If a reaction is carried out in a series of steps, H for the overall reaction will be equal to the sum of the enthalpy changes for the individual steps.”

28. Hess’s Law Because H is a state function, the total enthalpy change depends only on the initial state of the reactants and the final state of the products

29. Hess’s Law • The change in enthalpy for a reaction is the same whether the reaction takes place in one step or a series of steps. Germain Hess 1802 - 1850 Example: N2(g) + 2O2(g) 2NO2(g)H1 = 68 kJ N2(g) + O2(g) 2NO(g) H2 = 180 kJ 2NO(g) + O2(g)  2NO2(g)H3 = -112 kJ ---------------------------------------- ------------------

30. Using Hess’s Law ΔH C2H2(g) + 5/2O2(g) → 2CO2(g) + H2O(l) -1300kJ C(s) + O2(g) → CO2 - 394kJ H2(g) + 1/2O2(g) → H2O(l) - 286kJ 2C(s) + H2(g) C2H2(g) ??

31. Enthalpies of Formation An enthalpy of formation, Hf, is defined as the enthalpy change for the reaction in which a compound is made from its constituent elements in their elemental forms e.g. C(s) + 2H2 (g) → CH4 (g) ∆Hf CH4 = -74.8kJ/molrxn

32. Enthalpy of Formation Hf is the enthalpy of formation of onemole of a substance from its constituent elements • 2Na + Cl2 2NaCl H =  822kJ/mol • Na + 1/2Cl2 NaCl Hf =  411kJ/mol NaCl Formation

33. Standard Enthalpies of Formation  Standard enthalpies of formation, Hf, are measured under standard conditions (25oC and 1.00 atm pressure)

34. Calculation of H We can use Hess’s law in this way: H = nHf(products) - mHf(reactants) where n and m are the stoichiometric coefficients  

35. Using ΔHf C + O2 CO +1/2O2 → CO2ΔH? ΔHf CO = -110.5kJ/mol ΔHf CO2 = -393.5kJ/mol