Introduction to Inorganic Chemistry: Thermodynamics and Kinetics

1. CHEM 120: Introduction to Inorganic Chemistry Instructor: Upali Siriwardane (Ph.D., Ohio State University) CTH 311, Tele: 257-4941, e-mail: upali@chem.latech.edu Office hours: 10:00 to 12:00 Tu & Th ; 8:00-9:00 and 11:00-12:00 M,W,& F

2. Chapters Covered and Test dates • Tests will be given in regular class periods  from  9:30-10:45 a.m. on the following days: September 22,     2004 (Test 1): Chapters 1 & 2 • October 8,         2004(Test 2):  Chapters  3, & 4 • October 22,         2004 (Test 3): Chapter  5 & 6 • November 12,     2004 (Test 4): Chapter  7 & 8 • November 15 Brief survey of chpater 9-10 • November 17,      2004 MAKE-UP: Comprehensive test (Covers all chapters 1-8) • Grading: • [( Test 1 + Test 2 + Test3 + Test4 + Test5)] x.70 + [ Homework + quiz average] x 0.30 = Final Average •                               5

3. Chapter 8. Chemical andPhysical Change: Energy, Rate, and Equilibrium Thermodynamics 1. Endothermic and exothermic based on heat flow between a system and its surroundings. 2. Enthalpy (DH), entropy (DS) , and free energy (DG) 3. Experiments to get thermochemical information and fuel values. Reaction Rates 4. Reaction rate and the role of kinetics in chemical and physical change. 5. Activation energy and the activated complex and effects on reaction rate. 6. Predict the affect of, concentration, temperature, and catalysis on the rate of a chemical reaction. 7. Write rate equations for elementary processes. Chemical Equilibrium 8. Equilibrium chemical reactions. 9. Equilibrium-constant expressions and equilibrium constants. 10.LeChatelier's principle for predicting equilibrium position.

4. Thermodynamics • Thermodynamics is study of energy, heat and work. • As chemists we are interested in heat changes in chemical and physical changes. • Aim: to predict whether a change (both physical and chemical ) will occur spontaneously when left to itself--do we have to do anything other than mix the reactants together to make it occur?

5. Chemical reaction and energy • In a chemical reaction heat is released or absorbed by our system to or from the surroundings. In a chemical reaction convert energy in bonds into heat energy (and vice versa). • We can measure the energy changes in these processes.

6. universe system surroundings Universe • System: • Surroundings: • Universe:

7. First law of thermodynamics • Law of Conservation of Energy:

8. Thermochemistry Heat changes during chemical reactions Thermochemical equation. eg. H2 (g) + O2 (g) ---> 2H2O(l) DH =- 256 kJ; DH is called the enthalpy of reaction. if DH is + reaction is called endothermic if DH is - reaction is called exothermic

9. Exothermic processes • In an exothermic process

10. Exothermic rxn • Reactants • DE • products

11. Endothermic Reaction • In an endothermic reaction

12. Endothermic process • products • DE • reactants

13. Enthalpy, H • Define enthalpy (H): • Normally talk about a change in enthalpy (DH, DHo):

14. Enthalpy, H • If DH is negative

15. Exothermic Reaction • Reactants • DH • products

16. Enthalpy, H • If DH is positive

17. Endothermic process • products • DH • reactants

18. Are these exo- or endothermic? • When solid NaOH is dissolved in water the soln gets hotter. • S(s) + O2(g) g SO2(g) DHo = -71kcal • N2(g) + 2O2(g +16.2kcal g 2NO2(g)

19. Spontaneous Processes • A spontaneous reactions is one that occurs without us having to do anything to it (once it has started). (No external energy input) • Name some spontaneous processes. • Note that one direction is spontaneous, the reverse is not.

20. We want to come up with a way of predicting whether something will be spontaneous. • It has been observed that many exothermic processes are spontaneous. • Question: Does exothermicity guarantee that something is spontaneous?

21. Look at H2O(s) gH2O(l) DHorxn=1.44kcal

22. Entropy, S • Entropy (S, So) is a measure of the disorder, or randomness of a system. • The greater

23. Entropy info

24. For a rxn: DSrxn= sum of entropy of all the productsminus the sum of the entropy of all the reactants.

25. Get entropy increases • When go from a system of

26. Is there an entropy increase when • A log burns in a fireplace • Water vapor condenses on a cold surface • A solid metal melts • Water boils

27. Does a positive entropy change insure a spontaneous process • Look at H2O(l) gH2O(s) • DS<0 but

28. Second law of thermodynamics • The entropy of the universe increases in a spontaneous process and is equal to zero in a system at equilibrium. • Not always easy to calculate the entropy change of the universe. • Define a new quantity, G, Gibbs free energy. G refers to the system we are studying.

29. DG = DH - T DS • This equation combines the exothermicity and positive entropy criteria. • Criterion for spontaneity • If DG < 0 • If DG > 0 • If DG = 0

30. How does DG = DH - T DSwork? • Look at H2O(l) gH2O(g) water boiling • DH= 10.6kcal and DS =0.0284kcal • DG = DH -T DS • at 50oC DG = • at 100oC DG = • at 120oC DG =

31. DG = DH - T DS • An exothermic rxn that has a positive entropy change is • An endothermic rxn that has a negative entropy change is

32. DG = DH - T DS • An exothermic rxn that has a negative entropy change is spontaneous • An endothermic rxn that has a positive entropy change is spontaneous

33. Calorimetry: how to measure heat changes in reactions • Measure heat change (temp inc or dec) in a quantity of water or solution that is in contact with the reaction of interest and is isolated from the surroundings.

34. Constant Volume (Bomb) Calorimeter • Used for combustion reactions, etc. • CxHy + (x+y/2)O2g xCO2 +yH2O • Measure t of H2O bath, calculate Hsurr, determine Hrxn

35. To calculate the amt of heat (Q) absorbed or released • Q= m x SH x T • Q = heat change; T = temperature change = Tfinal-Tinitial • Add heat, temp inc; remove heat temp dec • Specific heat (SH) = the amount of heat (cal) required to raise the temperature of one g of a substance by 1°C

36. Units of SH : cal/(g oC); [SH of H2O= 1.00 cal/(g oC); of Al = 0.21 cal/(g oC)] • Calc the amt of heat liberated (in cal and kcal) from 366 g of aluminum when it cools from 77.0oC to 12.0oC. • 10kJ of heat is supplied to 1000g of H2O and to 1000g of Al. Calc the inc in temp for both.

37. To raise the temp of a mass of water from 25oC to 50oC requires 7.5 kcal. What is the mass of the water? • A sample of Al weighs 67 g. If 854 cal of heat are required to raise the temp of this sample from 25oC to 85oC, calculate the specific heat of aluminum.

38. Fuel value • Fuel value is amt of energy per gram of food. • One nutritional Calorie (C) = 1 kcal = 1000 cal; 1 cal = 4.184 J; 1 kcal=1 Cal =4.184 kJ • 8.7: A 1.00 g sample of a candy bar (which contains a lot of sugar was burned in a bomb calorimeter. A 3.0oC temp increase was observed for 1.00 x 103 g of water. The entire candy bar weighed 2.5 ounces. Calc the fuel value (in nutritional Calories) of the sample and the total caloric content of the candy bar.

39. 8.8: If the fuel value of 1.00g of a certain carbohydrate is 3.00 nutritional Calories, how many grams of water must be present in the calorimeter to record a 5.00oC change in temp?

40. Kinetics (Reaction Rates) • Thermodynamics tells us whether a reaction should occur spontaneously, but does not tell us how fast the reaction will occur. Kinetics tells us • For example, thermodynamics says that diamond will spontaneously change into graphite. • Kinetics tells us

41. Look at H2(g) + I2(g) g 2HI(g) • Let’s envision how the rreaction might occur on a molecular basis. • To react • 1. H2 must collide with I2 • 2. There must be enough energy to break a H-H and a I-I bond to initiate reaction • 3. The molecules must collide in the correct geometry.

42. A molecule that is moving has kinetic energy; faster the motion, the greater the KE. When molecules collide some of the KE is changed into vibrational energy of the bonds . Sometimes there is enough energy gained through collision to break a bond and initiate reaction. If the energy is not enough to break the bond, the molecules bounce of one another with no reaction occurring. • So there is some minimum collision energy below which no reaction occurs.

43. Activation energy, Ea

45. Factors that affect reaction rate • I. Structure of reacting species: A. oppositely charged species often react faster than neutral species. B. bond strength can influence rxn rate (Ea) C. size and shape of molecule can be important

46. II. Concentration of reactants: if increase the conc of the reactants:

47. III. Increase the temperature:

48. IV. Physical state of reactants: reactions in solution (liquids) are often very fast. In the solid state molecules have limited motion, in the gas phase have large distances between molecules and not many collisions so these reactions may be slower. In the liquid phase the molecules (ions) are able to move and are close to each other.

49. V. Add a catalyst. A catalyst speeds up the rate of reaction by • . The catalyst generally works by giving a different pathway (of lower energy) for the reaction to occur. A catalyst is not used up (consumed) in a reaction. The catalyst appears to be unchanged at the end. • Biological catalysts: enzymes