Definitions and Basic Concepts

Transcription

1 Chemical Thermodynamics: Energy Changes in Chemical Systems Conversion of energy from one form to another Transfer of energy from one place to another Why do we care about Thermodynamics? Practical applications: energy production; fuels & foods energy conversion, storage & transfer Fundamental to understanding most areas of chemistry (equilibrium, kinetics, electrochemistry ) Allows us to predict reaction spontaneity

2 Definitions and Basic Concepts 1. Energy: the capacity to do work Units: Joule calorie BTU 1 cal = J a. Kinetic Energy: energy of motion KE = ½ m v 2 b. Potential Energy: stored energy -energy available due to an object s position -chemical energy: energy stored in chemical bonds

3 c. thermal energy : energy associated with the random motion of particles (atoms/ions/molecules) heat = transfer of thermal energy from warmer to cooler object symbol = q 2.a. System: specific part of universe under study 2.b. Surroundings: everything else in the universe Types of Systems: -open: - closed: - isolated:

4 3. State Functions: properties that are determined only by the current physical state of the system. Example: T = Fictitious Temperature - Time Plot Examples of State Functions include: Temperature, F :00 am 6:00 am 7:00 am 8:00 am 9:00 am 10:00 am 11:00 am noon 1:00 pm 2:00 pm

5 4. Specific Heat: ρ water = 5. Heat Capacity:

6 Calculating thermal energy transfer using C or ρ Let q = Conventions for q: q ( ) q (-) Note: 1 st Law of Thermodynamics

7 Energy Changes Associated with Phase Transitions A. Melting/freezing: B. Vaporization/condensation:

8 Enthalpies of vaporization and fusion for some selected substances Silberberg Fig. 12.1

9 Heating/cooling Curve for Water at 1 atm Temperature, C Thermal Energy Added ->

10 How are Thermal Energy Changes Measured? 1. Constant Pressure Calorimetry Silberberg Fig 6.10

11 2. Thermal Energy Change at Constant Volume: Silberberg Fig 6.11

12 In addition to gaining or losing thermal energy (q) in a process, Sign conventions q (+) q (-) w (+) w (-)

13 E = internal energy: many factors contribute to the value of E for a system. Components of Internal Energy Contributions to the kinetic energy: The molecule moving through space, E k(translation) The molecule rotating, E k(rotation) The bound atoms vibrating, E k(vibration) The electrons moving within each atom, E k(electron) Contributions to the potential energy: Forces between the bound atoms vibrating, E p(vibration) Forces between nucleus and electrons and between electrons in each atom, E p(atom) Forces between the protons and neutrons in each nucleus, E p(nuclei) Forces between nuclei and shared electron pair in each bond, E p(bond)

14 Some components of internal energy

15 Note: it is not possible to determine absolute values for the internal energy of a system (E), but it is possible to determine changes in internal energy, E. C 8 H 18 (l) O 2 8CO 2 (g) + 9H 2 O(l) C 8 H 18 (l) O 2 8CO 2 (g) + 9H 2 O(l)

16 How does a system (like a chemical reaction) do work? Work done by a chemical reaction: Zn(s) + 2HCl(aq) ZnCl 2 (aq) + H 2 (g) V = V f -V i Note: opposing pressure P is constant Silberberg Fig 6.7 Silberberg Fig 6.4

17 A chemical reaction can do work if Zn(s) + 2HCl(aq) ZnCl 2 (aq) + H 2 (g) n = 3C(s) + 4H 2 (g) C 3 H 8 (g) n =

18 Determining Work Done by (or on) a Reaction System at Constant Pressure

19 Enthalpy of Reaction = H or H rctn Some Important Types of Enthalpy Change heat of combustion ( H comb ) C 4 H 10 (l) + 13/2O 2 (g) 4CO 2 (g) + 5H 2 O(g) heat of formation ( H f ) K(s) + 1/2Br 2 (l) KBr(s) heat of fusion ( H fus ) NaCl(s) NaCl(l) heat of vaporization ( H vap ) C 6 H 6 (l) C 6 H 6 (g)

20 Thermochemical Equations Information about both reaction stoichiometry and enthalpy change for the reaction as written H 2 O(l) H 2 O(g) H = o C (= H vap ) CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O(l) H = o C

21 Rules for Working with Thermochemical Equations

22 Standard Enthalpy of Reaction H o When a reaction is carried out under thermodynamic standard conditions, the enthalpy change is H o CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O(l) H o = o C CH 4 (g) + 2O 2 (g) CO 2 (g) + 2H 2 O(g) H o = o C

23 H o f = Standard Enthalpy of Formation H o f C(s) + O 2 (g) CO 2 (g) H o = H o f of CO 2 (g)

24 Table 6.5 Selected Standard Heats of Formation at 25 0 C(298K) Formula calcium Ca(s) CaO(s) CaCO 3 (s) carbon C(graphite) C(diamond) CO(g) CO 2 (g) CH 4 (g) CH 3 OH(l) HCN(g) CS s (l) chlorine Cl(g) H 0 f (kj/mol) Formula Cl 2 (g) HCl(g) hydrogen H(g) H 2 (g) nitrogen N 2 (g) NH 3 (g) NO(g) oxygen O 2 (g) O 3 (g) H 2 O(g) H 2 O(l) H 0 f (kj/mol) Formula silver Ag(s) AgCl(s) sodium Na(s) Na(g) NaCl(s) H 0 f (kj/mol) sulfur S 8 (rhombic) 0 S 8 (monoclinic) 2 SO 2 (g) SO 3 (g)

25 Using H o f to Calculate Ho rctn

26 Entropy = Disorder = S Third Law of Thermodynamics:

27 Second Law of Thermodynamics:

28 For a system at constant pressure

29 Under standard state conditions, If G < 0 (-) If G > 0 (+) If G = 0 H S G

30 When H is (-) and S is (-): 200 G vs Temp: G = H - T S H = -150 kj; S = -250 J/K Temp, C H T S G Π G, kj Temp, C

31 When H is (+) and S is (+): 200 Free Energy Change vs Temp H= +95kJ; S = +225 J/K delta G, kj Temp, C

32 Two Ways to Calculate Go