Chapter 17

1. Chapter 17 Reaction Energy and Reaction Kinetics

2. Sect. 17-1: Thermochemistry • Thermochemistry – the study of the transfers of energy as heat that accompany chemical reactions and physical changes • Calorimeter – device used to measure the energy absorbed or released as heat in a chemical or physical change

3. Temperature – a measure of the average KE of the particles in a sample of matter • Joule (J) – the SI unit of heat and energy; kJ is also commonly used • Heat – energy transferred between samples of matter because of a difference in their temperatures

4. Specific heat – the amount of energy required to raise the temperature of one gram of a substance by one degree Celsius. • Units are typically J/(g x°C) or cal/(g x°C) • q = Cp x m x ΔT, where q is heat, Cp is specific heat, m is mass, & ΔT is change in temperature

5. Example: a 4.0 g sample of glass was heated from 274 K to 314 K, a temperature increase of 40 K and was found to have absorbed 32 J of energy as heat. What is the specific heat of this type of glass?

6. Heat of reaction – quantity of energy released or absorbed as heat during a chemical reaction (difference between stored energy of reactants and products) • Thermochemical equation – an equation that includes the quantity of energy released or absorbed as heat Ex: 2 H2 (g) + O2 (g) 2 H2O (g)+ 483.6kJ

7. Enthalpy change (Δ H) – the amount of energy absorbed or lost by a system as heat during a process at constant pressure • ΔH = Hproducts – Hreactants • Negative for exothermic reactions and positive for endothermic reactions

8. Important things to remember • Coefficients represent # moles & can be written as fractions if need be • Physical state of reactants/products matters • Change in energy is directly proportional to number of moles reacting • ΔH is usually not significantly influence by changing temperature

9. Molar heat of formation – the energy released or absorbed as heat when one mole of a compound is formed by combination of its elements • When given for the standard state of that substance it is written as ΔH0f; the 0 is for standard state and the f for heat of formation

10. Substances that have a large negative ΔH0f are very stable • Small negatives or small positive ΔH0f are relatively unstable and will decompose easily • Large positive ΔH0f are very unstable and will decompose or react violently

11. Heat of combustion ΔH0c – the energy released as heat by the complete combustion of one mole of a substance

12. Calculating Heats of Reaction • Hess’s Law – the overall enthalpy change in a reaction is equal to the sum of the enthalpy changes for the individual steps in the process • If a reaction is reversed, the sign of ΔH is also reversed • Multiply the coefficients as needed

13. Example • Calculate the heat of reaction for the combustion of nitrogen monoxide gas to form nitrogen dioxide gas. • NO + ½ O2 NO2 • From table A-14 on pg. 902: ½ N2 + ½ O2 NO ΔH0f= +90.29kJ/mol ½ N2 + O2 NO2ΔH0f = +33.2 kJ/mol

14. Heats of formation can be determined by combining the heat of formation and heat of combustion for various substances

15. Sect 17-2: Driving Force of Reactions • Most spontaneous reactions tend toward products that have a lower energy state than the reactants (exothermic reactions) • However, some endothermic reactions do occur spontaneously

16. Entropy (S) – a measure of the degree of randomness of the particles in a system • Reactants tend towards a less ordered state of matter (example: ice melting – liquid is less organized than solid) • In general, gases have the highest entropy, then liquids, and then solids • ΔS is postive for an increase in entropy and negative for a decrease in entropy

17. Free energy • Nature drives processes toward lowest enthalpy and highest entropy, when these are opposite directions, the dominant factor determines the direction • Free energy (G) – the combined enthalpy-entropy function • ΔG0 = ΔH0 - TΔS0

18. Example • For the reaction NH4Cl  NH3 + HCl at 298.15 K, ΔH0= 176 kJ/mol and ΔS0= 0.285 kJ/(molK). Calculate ΔG0, and tell whether this reaction can proceed in the forward direction at 298.15K. • + 91 kJ/mol, so it does not occur naturally at this temperature

19. Sect. 17-3: The Reaction Process • Reaction mechanism – the step-by-step sequence of reactions by which the overall chemical change occurs • Intermediates – species that appear in some steps but not in the net equation • Homogeneous reaction – a reactions whose reactants and products exist in a single phase

20. Collision theory – set of assumptions regarding collisions and reactions • if collision is too gentle, the species rebound unchanged • If colliding species are poorly oriented, they will not react

21. Activation energy (Ea) – the minimum energy required to transform the reactants into an activated complex • Activated complex – a transitional structure that results from an effective collision and that persists while old bonds are breaking and new bonds are forming (not the same as intermediate)

22. activation energy for forward reaction • Activation energy for reverse reaction • Energy change in reaction http://www.bbc.co.uk/scotland/education/bitesize/higher/img/chemistry/calculations_1/pe_diags/fig10.gif

23. Sect. 17-4: Reaction Rate • Reaction rate – the change in concentration of reactants per unit time as a reaction proceeds • Chemical kinetics – the area of chemistry that is concerned with reaction rates and reaction mechanisms

24. Rate-influencing factors • Nature of reactants • Surface area • Temperature • Concentration • Presence of catalysts

25. Catalyst – increases rate of reaction with out being used up • Catalysis – action of a catalyst • Homogeneous catalyst – same phase as reactants/products • Heterogeneous catalyst – different phase as reactants/products

26. Rate law – an equation that relates reaction rate and concentrations of reactants for a reaction • If multiple steps in reaction mechanism, the slow step always controls the rate