Thermodynamics

1. Part I (Yep, there’ll be a Part II) Thermodynamics

2. Energy • The capacity to do work or transfer heat • Measured in Joules • Two Types • Kinetic (motion) • Potential (based on position)

3. 1st Law of Thermodynamics

4. 1st Law of Thermodynamics • aka Law of Conservation of Energy • Energy can be converted from one form to another but cannot be created or destroyed

5. Heat vs Temperature Heat Temperature • A measure of energy content • What is transferred during a temperature change • Units Joules • Reflects random motion of particles in a substance • Indicates the direction in which heat energy will flow • Units °C or K

6. Energy, Heat, and Work • ∆E = q + w • Change in energy equals heat plus work • Energy and heat are state functions. Work is not • A state function is independent of the pathway taken to get to that state. (only the beginning and end matter)

7. Sign Conventions for q, w, and ∆E

8. Work, Pressure, and Volume • w = -P∆V

9. Chemical Energy Endothermic Exothermic • Rxns in which energy is absorbed from the surroundings • Energy flows into the system to increase the potential energy of the system • Rxns that give off energy as they progress • Some of the potential energy stored in the chemical bonds is converted to thermal energy (random KE) through heat • Products are more stable (stronger bonds) than reactants

10. Enthalpy • H = E + PV • H – enthalpy • E – internal energy • P- pressure • V – volume • In systems at constant pressure, where the only work is PV, the change in enthalpy is due only to energy flow as heat • ∆H = heat of rxn

14. Sample Problem • Indicate the sign of the enthalpy change, ∆H, in each of the following processes carried out under atmospheric pressure, and indicate whether the process is endothermic or exothermic • An ice cube melts • 1g of butane is combusted

15. Enthalpies of Reaction • ∆H = Hproducts- Hreactants • Enthalpy is an extensive property •  The enthalpy change for a reaction is equal in magnitude, but opposite in sign, to ∆H for the reverse reaction. •  The enthalpy change for a reaction depends on the state of the reactants and products

17. Sample Problem • How much heat is released when 4.50g of methane gas is burned in a constant pressure system? (∆H = -890 kJ)

18. Calorimetry • Science of measuring heat • Heat Capacity (C) • C= heat absorbed/Temp increase • Specific Heat Capacity • Energy required to raise the temp of 1 gram of a substance 1 °C • Molar Heat Capacity • Energy required to raise the temp of 1 mole of a substance 1 °C

19. Calorimetry • ∆H = mc∆T or q = mc∆T • m – mass (g) • c – specific heat (J/g·K) • ∆T – change in temperature (K)

20. Sample Problem • How much heat is needed to warm 250 g of water from 22°C to near its boiling point, 98°C? (The specific heat of water is 4.18 J/g·K) • What is the molar heat capacity?

21. Constant Pressure Calorimetry Coffee Cup Calorimeter

22. Sample Problem • When a student mixes 50 mL of 1.0 M HCl and 50 mL of 1.0 M NaOH in a coffee-cup calorimeter, the temperature of the resultant solution increases from 21.0°C to 27.5°C. Calculate the enthalpy change for the reaction in kJ/mol HCl, assuming that the calorimeter loses only a negligible quantity of heat, that the total volume of the solution is 100 mL, that its density is 1.0 g/mL, and that its specific heat is 4.18 J/g·K.

23. Constant Volume Calorimetry Bomb Calorimeter

24. Hess’s Law • In going from a particular set of reactants to a particular set of products, the change in enthalpy (∆H) is the same whether the reaction takes place in one step or a series of steps

25. Hess’s Law

26. Using Hess’s Law • Work backward from the final reaction • Reverse reactions as needed, being sure to also reverse the sign of ∆H • Remember that identical substances found on both sides of the summed equation cancel each other out.

27. Sample Problem • The enthalpy of reaction for the combustion of C to CO2 is -393.5 kJ/mol·C, and the enthalpy for the combustion of CO to CO2 is -283.0 kJ/mol·C. Using this data, calculate the enthalpy for the combustion of C to CO.

28. Standard State • For a compound • Gaseous state • Pressure of 1 atm • Pure liquid or solid • Standard state IS the pure liquid or solid • Substance in a soln • Concentration of 1 M • For an element • The form in which the element exists at 1 atm and 25°C

29. Standard Enthalpies of Formation (∆Hf°) • The change in enthalpy that accompanies the formation of one mole of a compound from its elements with all elements in their standard state

30. Calculating Enthalpy Change • When a rxn is reversed, the magnitude of ∆H remain the same, but its sign changes. • When the balanced eqn for a rxn is multiplied by an integer, the value of ∆H must be multiplied by the same integer • The change in enthalpy for a rxn can be calculated from the enthalpies of formation of the reactants and products ∆H°rxn=Σ∆Hf°products - Σ∆Hf°reactants • Elements in their standard states are not included • For elements in their standard state, ∆Hf° = 0

31. Sample Problem • For which of the following reactions at 25°C would the enthalpy change represent a standard enthalpy of formation? For those where it does not, what changes would need to be made in the reaction conditions? • 2Na(s) + ½ O2(g)  Na2O(s) • 2K(l) + Cl2(g) 2KCl(s) • C6H12O6(s)  6C (diamond) + 6H2(g) + 3O2(g)

32. Sample Problem • Calculate the standard enthalpy change for the combustion of 1 mol of benzene, C6H6.

33. Sample Problem • Compare the quantity of heat produced by combustion of 1.00g propane, C3H8, to that produced by 1.00 g benzene.