1. Back Chapter 17: Energy and Kinetics Thermochemistry: Causes of change in systems Kinetics: Rate of reaction progress (speed) Heat, Energy, and Temperature changes Standard unit of heat is the Joule, J Standard unit of temperature is Kelvin, K Pages 510-547

2. Back Bires, 2010-b Slide 2 Heat vs Temperature Heat Heat –measure of energy change in a system. Temperature Temperature –measure of the average kinetic energy (movement) of the particles in a system. Exothermic Exothermic –System loses energy to surroundings Endothermic Endothermic –System gains energy from surroundings K = C + 273.15

3. Back Bires, 2010-b Slide 3 Specific Heat Specific Heat Specific Heat –measure of how a substance reacts to heat energy changes. –Think thermal inertia –is the heat energy required to raise one gram of a pure substance one degree Celsius. –is a property of matter; different species have different Specific Heats. The symbol we use is c p. The symbol we use is c p. –The “p” stands for constant pressure while heat is added or lost.

4. Back Bires, 2010-b Slide 4 Specific Heat Capacity Metals have very low c p, Metals have very low c p, – which is why metals often feel cold to the touch. Water has a very high c p, Water has a very high c p, –4.184 J/g· 0 C Substances with lower c p will rise in temperature faster and require less energy to do so than do substances with high c p. Substances with lower c p will rise in temperature faster and require less energy to do so than do substances with high c p. Substance J/(g x o C) or J/(g x K) cal/(g x o C) or cal/(g x K) Water (0 o C to 100 o C) 4.1841.000 Zinc.3870.093 Ice (-10 o C to 0 o C) 2.0930.500 Steam (100 o C) 2.0090.480 Brass.3800.092 Wood (typical) 1.6740.400 Soil (typical) 1.0460.250 Air (50 o C) 1.0460.250 Aluminum.9000.215 Tin.2270.205 Glass (typical).8370.200 Iron/Steel.4520.108 Copper.3870.0924 Silver.2360.0564 Mercury.1380.0330 Gold.1300.0310 Lead.1280.0305

5. Back Bires, 2010-b Slide 5 Specific Heat Capacity C p units are J/g 0 C) C p units are J/g 0 C) Change in heat (joules, J) Change in heat (joules, J) = Change in temperature (degree, 0 C) Change in temperature (degree, 0 C) x Mass (mass, g) Mass (mass, g) x Specific Heat Capacity (4.184 for water) Specific Heat Capacity (4.184 for water) 1 calorie = 4.184 Joules Cp (H 2 O) = 4.184

6. Back Bires, 2010-b Slide 6 Specific Heat Example Exercise Determine the specific heat of 34 grams of an unknown material if 485 J of heat are absorbed to change the temperature by 20.0 o C. Determine the specific heat of 34 grams of an unknown material if 485 J of heat are absorbed to change the temperature by 20.0 o C. If 950 J of heat are added to 5.4 mL of water at 280 K, what will be the resulting temperature of the water? (hint: mL  g) If 950 J of heat are added to 5.4 mL of water at 280 K, what will be the resulting temperature of the water? (hint: mL  g)

7. Back Bires, 2010-b Slide 7 The Calorimeter The Calorimeter (shown) Heat energy is transferred from a reaction inside the calorimeter to the water in the calorimeter. The temperature change of the water is observed. Text page 519 When two objects are in contact, they eventually obtain Thermal Equilibrium; their temperatures become equal.

8. Back Bires, 2010-b Slide 8 Enthalpy, ΔHEnthalpy, ΔHEnthalpy, ΔHEnthalpy, ΔH Enthalpy Enthalpy –heat energy transferred for a specific change to take place. We specify enthalpy with ΔH. “ Δ ” means “change in”. We specify enthalpy with ΔH. “ Δ ” means “change in”. Exothermic reaction Exothermic reaction –negative enthalpy ( -ΔH ) Endothermic reaction Endothermic reaction –Positive enthalpy (+ ΔH). Elements in their standard (elemental) state have a ΔH of zero. Elements in their standard (elemental) state have a ΔH of zero. –O 2, Fe, Cu, N 2, He, etc are have H f = 0 kj/mol The universe favors LOW energy states - if the products have lower energy reaction is favored.

9. Back Bires, 2010-b Slide 9 Enthalpy, ΔHEnthalpy, ΔHEnthalpy, ΔHEnthalpy, ΔH Some common changes involving ΔH: Some common changes involving ΔH: ΔH fus = heat of fusion ΔH fus = heat of fusion ΔH vap = heat of vaporization ΔH vap = heat of vaporization ΔH cond = heat of condensation ΔH cond = heat of condensation ΔH sub = heat of sublimation ΔH sub = heat of sublimation ΔH rxn = heat of reaction ΔH rxn = heat of reaction ΔH f = heat of formation ΔH f = heat of formation ΔH sol = heat of solution ΔH sol = heat of solution ΔH comb = heat of combustion ΔH comb = heat of combustion State change to/from? Sign of ΔH?

10. Back Bires, 2010-b Slide 10 Phase Changes 1. Solid + heat = temp ▲ 2. Solid + heat = phase change 3. Liquid + heat = ? 4. And then…? 5. Gas + heat = ? temperature Heat added temperature Heat added

11. Back Bires, 2010-b Slide 11 Heat of… Heat of fusion : solid to liquid Heat of fusion : solid to liquid Heat of vaporization : liquid to a gas Heat of vaporization : liquid to a gas Heat of condensation : gas to liquid Heat of condensation : gas to liquid Heat of sublimation : solid to gas Heat of sublimation : solid to gas Heat of reaction : during a chemical reaction Heat of reaction : during a chemical reaction Heat of formation : energy (used) to form 1 mole of a compound from its elements Heat of formation : energy (used) to form 1 mole of a compound from its elements Heat of solution : dissolving in a solvent Heat of solution : dissolving in a solvent Heat of combustion : energy released when a substance reacts with O 2 to form CO 2 and H 2 O. Heat of combustion : energy released when a substance reacts with O 2 to form CO 2 and H 2 O. Remember: if ΔH is positive- energy is going in, if negative, energy is coming out.

12. Back Bires, 2010-b Slide 12 Reaction Enthalpy If ΔH is negative, the reaction is exothermic. If ΔH is negative, the reaction is exothermic. C 6 H 12 O 6 + 6O 2  6CO 2 + 6H 2 O + 2870kJ C 6 H 12 O 6 + 6O 2  6CO 2 + 6H 2 O + 2870kJ ΔH rxn = -2870 kJ/mol ΔH rxn = -2870 kJ/mol If ΔH is positive, the reaction is endothermic. If ΔH is positive, the reaction is endothermic. 2H 2 O + 571.6kJ  2H 2 + O 2 2H 2 O + 571.6kJ  2H 2 + O 2 ΔH rxn = +571.6 kJ/mol ΔH rxn = +571.6 kJ/mol energy

13. Back Bires, 2010-b Slide 13Spontaneity Spontaneous Spontaneous –A reaction that will proceed on its own once started. Sometimes, all the reaction needs to get going is the kinetic energy of nearby colliding atoms. Sometimes, all the reaction needs to get going is the kinetic energy of nearby colliding atoms. Kinetic Molecular Theory: Kinetic Molecular Theory: –All Matter is made of particles in constant motion –Some collisions are more energetic than others. Why? Spontaneous combustion Spontaneous combustion –occurs when the kinetic energy of colliding oxygen molecules striking a fuel have enough energy on their own to start the combustion reaction.

14. Back Bires, 2010-b Slide 14 Exothermic Reaction: products have lower energy than do the reactants. What if endothermic?  H rxn

15. Back Bires, 2010-b Slide 15 In the diagram, the hump is called a activation energy barrier - the amount of energy required for the reaction to begin. All reactions require some sort of activation energy, E a. We can reduce the activation energy with a catalyst. Activated complex

16. Back Bires, 2010-b Slide 16 Hess’s Law: Hess’s Law: –If two reactions begin with the same reactants in the same condition and end with the same products in the same condition, they must have the same enthalpy change. –It doesn’t matter if you perform a reaction in several steps or produce your final product in one step, the enthalpy change will be the same. Consider the reaction A + B  D : Consider the reaction A + B  D : A + B + 100 kJ  C then C + 50 kJ  D A + B + 100 kJ  C then C + 50 kJ  D Must be the same as A + B + 150 kJ  D Must be the same as A + B + 150 kJ  D

17. Back Bires, 2010-b Slide 17 Hess’s Law Enthalpy of Reaction ΔH rxn = ∑H products – ∑H reactants = Σ H f, all the products – Σ H f, all the reactants “ Sum of ” Enthalpy of Formation

18. Back Bires, 2010-b Slide 18 Hess’s Law Example Exercise ΔH rxn = ∑H products – ∑H reactants Calculate the heat of reaction when 350 grams of methane, CH 4 are burned in excess oxygen. H f book values for each species are: Calculate the heat of reaction when 350 grams of methane, CH 4 are burned in excess oxygen. H f book values for each species are: CH 4(g) = -74.8 kJ/mol CH 4(g) = -74.8 kJ/mol O 2(g) = ? O 2(g) = ? H 2 O (g) = -285.83 kJ/mol H 2 O (g) = -285.83 kJ/mol CO 2(g) = -393.5 kJ/mol CO 2(g) = -393.5 kJ/mol

19. Back Bires, 2010-b Slide 19 Entropy, ΔS Entropy Entropy –is a measure of relative disorder. Thermodynamics tells us that the universe tends towards disorder or entropy. –Temperature affects entropy (why?) Entropy calculations are very similar to enthalpy calculations: Entropy calculations are very similar to enthalpy calculations: ΔS rxn = ΣS products – ΣS reactants Entropy has the unit J/K*mol

20. Back Bires, 2010-b Slide 20 Entropy, ΔS The universe tends towards entropy The universe tends towards entropy –entropy plays a part in predicting whether or not a reaction will be spontaneous. Solids have very low entropy Solids have very low entropy Gases have very high entropy Gases have very high entropy Solutions also have high entropy Solutions also have high entropy

21. Back Bires, 2010-b Slide 21 Qualitative Entropy Values We can make generalizations about a reaction’s entropy; We can make generalizations about a reaction’s entropy; 2KClO 3(s)  2KCl (s) + 3O 2(g) 2KClO 3(s)  2KCl (s) + 3O 2(g) 2 solids  2 solids + 3 gases 2 solids  2 solids + 3 gases Entropy appears to increase in this reaction. Entropy appears to increase in this reaction.

22. Back Bires, 2010-b Slide 22 Quantitative Entropy Values 2KClO 3(s)  2KCl (s) + 3O 2(g) 2KClO 3(s)  2KCl (s) + 3O 2(g) S of KClO 3(s) = 143.7 J/mol*K S of KClO 3(s) = 143.7 J/mol*K S of KCl (s) = 82.6 J/mol*K S of KCl (s) = 82.6 J/mol*K S of O 2(g) = 205.1 J/mol*K S of O 2(g) = 205.1 J/mol*K Using Δ S rxn = S products – S reactants, the reaction has a total entropy change of + 493.1 J/mol*K Using Δ S rxn = S products – S reactants, the reaction has a total entropy change of + 493.1 J/mol*K

23. Back Bires, 2010-b Slide 23 Entropy Values A positive ΔS = increase in entropy A positive ΔS = increase in entropy A negative ΔS = decrease in entropy A negative ΔS = decrease in entropy Do not confuse entropy and enthalpy! Tending toward spontaneity: Tending toward spontaneity: Negative Enthalpy (-ΔH) Negative Enthalpy (-ΔH) Positive Entropy (+ΔS) Positive Entropy (+ΔS)

24. Back Bires, 2010-b Slide 24 Free Energy, ΔG Free energy, ΔG: Free energy, ΔG: –allows us to assign a value to an entire reaction to predict whether a reaction is spontaneous, product favored. –or nonspontaneous, reactant-favored. –Named for American Chemist, J. Willard Gibbs ΔG = ΔH -TΔS Free Energy kJ/mol Enthalpy kJ/mol temperature in Kelvin Entropy J/mol·K

25. Back Bires, 2010-b Slide 25 Gibbs Free Energy, ΔG rxn Negative Gibbs Energy (-ΔG rxn ) Spontaneous, Product favored Positive Gibbs Energy (+ΔG rxn ) Nonspontaneous, Reactant favored A ΔG of zero means that neither the products nor reactants are favored-the reaction is in equilibrium.

26. Back Bires, 2010-b Slide 26  G =  H - T  S

27. Back Bires, 2010-b Slide 27 Reaction Rates Reaction Rates Reaction rates Reaction rates –how fast a reaction proceeds. Some factors will affect reaction rate: Some factors will affect reaction rate: Temperature of reactants: higher = faster Temperature of reactants: higher = faster Concentration of reactants: greater = faster Concentration of reactants: greater = faster Surface area of reactants: greater = faster Surface area of reactants: greater = faster –(powders react faster than chunks) Pressure of gaseous reactants: greater = faster Pressure of gaseous reactants: greater = faster Catalyst presence: catalysts make rxns faster Catalyst presence: catalysts make rxns faster –reduce activation energy! –are not used up (not reactants)

28. Back Bires, 2010-b Slide 28 Rate Laws For any reaction: For any reaction: The rate is based on the [reactants]: The rate is based on the [reactants]: [X] : “1 st order” : 2x [A], 2x rate [X] : “1 st order” : 2x [A], 2x rate [X] 2 : “2 nd order” : 2x [A], 4x rate [X] 2 : “2 nd order” : 2x [A], 4x rate [X] 3 : “3 rd order” : 2x [A], 8x rate [X] 3 : “3 rd order” : 2x [A], 8x rate End of C17, conclusion follows

29. Back Bires, 2010-b Slide 29 In conclusion… Specific Heat Capacity, c p (J/gK) Specific Heat Capacity, c p (J/gK) –the amount of heat energy required to raise 1 gram, 1 degree Enthalpy, ΔH (kJ/mol) Enthalpy, ΔH (kJ/mol) –the heat energy transferred in a reaction Entropy, ΔS (J/mol-K) Entropy, ΔS (J/mol-K) –the change in disorder of the species in a reaction Gibbs Free Energy, ΔG (kJ/mol) Gibbs Free Energy, ΔG (kJ/mol) –measure of spontaneity; how product favored or reactant favored a reaction is Recall that K = C + 273.15

30. Back Bires, 2010-b Slide 30 CCSD Syllabus Objectives 16.1: Thermodynamics, definition 16.1: Thermodynamics, definition 16.2: Exothermic/Endothermic 16.2: Exothermic/Endothermic 16.3: Changes in Enthalpy 16.3: Changes in Enthalpy 16.4: Thermochemical Calculations 16.4: Thermochemical Calculations 16.5: Energy Diagrams 16.5: Energy Diagrams 16.6: Enthalpy-Entropy-Free Energy 16.6: Enthalpy-Entropy-Free Energy 17.1: Kinetics Definition 17.1: Kinetics Definition 17.2: Factors that Affect Reaction Rate 17.2: Factors that Affect Reaction Rate

31. Back Bires, 2010-b Slide 31 Aligned Labs and Demos Lab: Flaming Cheeto Calorimetry Lab Lab: Flaming Cheeto Calorimetry Lab Lab: Metals Calorimetry Lab Lab: Metals Calorimetry Lab Lab: NaOH-HCl Enthalpy of Reaction Lab Lab: NaOH-HCl Enthalpy of Reaction Lab Lab: Ba(OH) 2 -8H 2 O w/ NH 4 NO 3 and H 2 O 2 with a catalyst Free Energy Lab Lab: Ba(OH) 2 -8H 2 O w/ NH 4 NO 3 and H 2 O 2 with a catalyst Free Energy Lab Lab: KI-H 2 O 2 Kinetics Lab Lab: KI-H 2 O 2 Kinetics Lab Demo: Carbon Snake with powdered vs granular sugar (Kinetics) Demo: Carbon Snake with powdered vs granular sugar (Kinetics)