Chem. 31 – 10/01 Lecture

1. Chem. 31 – 10/01 Lecture

2. Announcements I • Pipet/Buret Calibration Lab Report • Resubmissions must be turned in within 1 week of receiving report • Attach new report form to graded report • Exam 1 • Next Wednesday • Will review topics on next Monday • Will cover Ch. 1-4 + sections of Ch. 6 covered today (+ in lecture notes today but not gotten to) • Help Session ?? Times (Monday 12 to 1 – in place of office hour or 6 to 7)

3. Announcements II • Today’s Lecture • Chemical Equilibrium • relating reactions to equilibrium equations • manipulating reactions • thermodynamics • le Chatelier’s principle

4. Ch. 3 and 4– What you need to know Equations you need to know: Average calculation t and Z based confidence intervals line equation Equations I will provide: Propagation of uncertainty for +/-, *//, and exponent Standard deviation Case 2 and 3 t-test, F-test and Grubbs test (if needed)

5. Equilibrium Equations Equilibrium Equations from Chemical Equations (Reactions) Generic Example: aA + bB ↔ cC + dD (Reaction) Equilibrium Equation Compounds are in equation if in solution (not present as solid, or solvent). Concentrations are in M but K is unitless Similar equation for gases (except with PAa replacing [A]a)

6. Equilibrium Equations Example problem: Write equation for reaction: AgCl(s) + 2NH3(aq) ↔ Ag(NH3)2+(aq) + Cl-(aq) AgCl not included because it is a solid

7. Equilibrium Equations- manipulating reactions Flipping Directions - If for A ↔ B, K = K1, then for B ↔ A, K = 1/K1 b) Adding Reactions NH4+↔ NH3(aq) + H + H+ + OH-↔ H2O(l) NH4+ + OH-↔ NH3(aq) + H2O(l) Reaction 3) = rxn1) + rxn2) So K3 = K1K2

8. Equilibrium Equations- manipulating reactions c) Multiplication 2x[½ N2 (g) + ½ O2 (g) ↔ NO (g)] K = K1 N2 (g) + O2 (g) ↔ 2NO (g) K = K12

9. Equilibrium Equation Example Problem: If the following reactions have the given equilibrium constants: Ag+ + 2NH3(aq) ↔ Ag(NH3)2+ K = 1.70 x 107 NH3(aq) + H2O(l) ↔ NH4+ + OH-K = 1.76 x 10-5 H2O(l) ↔ H+ + OH-K = 1.0 x 10-14 Determine the equilibrium constant for the following reaction: Ag(NH3)2++ 2H+→ Ag+ + 2NH4+

10. Thermodynamics ΔH is related to heat of reaction - if a reaction produces heat, ΔH < 0 and reaction is “exothermic” - a reaction that requires heat has ΔH > 0 and is endothermic ΔS is related to disorder of system - If the final system is “more random” than initial system, ΔS > 0

11. Thermodynamics Entropy Examples: (Is ΔS > or < 0?) H2O(l) ↔ H2O(g) H2O(s) ↔ H2O(l) NaCl(s) ↔ Na+ + Cl- 2H2(g) + O2(g) ↔ 2H2O(g) N2(g) + O2(g) ↔ 2NO(g) ΔS > 0 ΔS > 0 ΔS > 0 ΔS < 0 ΔS > 0

12. Thermodynamics ΔG = Change in Gibbs free energy This tells us if a process is spontaneous (expected to happen) or non-spontaneous ΔG < 0 process is spontaneous (favored) ΔG = ΔH - TΔS (T is absolute temperature) processes that are exothermic (Δ H < 0) and increase disorder (Δ S > 0) are favored at all T processes that have Δ H > 0 and Δ S > 0 are favored at high T

13. Example question The reaction N2(g) + O2 (g) ↔ 2NO(g) has a positive DH. Under what conditions is this process spontaneous? - all temperatures - low temperatures - high temperatures - never

14. Thermodynamics ΔG and Equilibrium ΔG = ΔG° + RTlnQ Q = Reaction Quotient (for A ↔ B, Q = [B]/[A]) At equilibrium, ΔG = 0 and Q = K ΔG° = -RTlnK

15. Thermodynamics Example Question: The ΔG° for the reaction Ca2+ + 2OH- => Ca(OH)2(s) is -52 kJ/mol Determine K at T = 20°C for Ca(OH)2(s) => Ca2+ + 2OH-

16. Section 6 – 2: Le Châtelier’s Principle The position of a chemical equilibrium always shifts in a direction that tends to relieve the effect of an applied stress Types of stresses: Addition/removal of reactant/product Change in volume (gases) or dilution (aqueous solutions) Changes in temperatures Can take intuitive or mathematical approaches to solving problems

17. Le Châtelier’s Principle Intuitive Method Addition to one side results in switch to other side Example: Mathematical Method reaction shifts to reactants (more AgCl(s)) When Q>K, ΔG>0 (toward reactants) When Q<K, ΔG<0 (toward products) Example: Q = [Ag+][Cl-] As Ag+ increases, Q>K AgCl(s) ↔ Ag+ + Cl- Addition of Ag+

18. Le Châtelier’s Principle Stress Number 1 Reactant/Products: Addition of reactant: shifts toward product Removal of reactant: shifts toward reactant Addition of product: shifts toward reactant Removal of product: shifts toward product

19. Le Châtelier’s Principle Stress Number 1 Example: CaCO3(s) + 2HC2H3O2(aq) ↔ Ca(C2H3O2)2(aq) + H2O(l) + CO2(g) 1. Add HC2H3O2(aq) 2. Remove CO2(g) 3. Add Ca(C2H3O2)2(aq) 4. Add CaCO3(s) No effect because (s)

20. Le Châtelier’s Principle Stess Number Two: Dilution Side with more moles is favored at lower concentrations Example: HNO2(aq) ↔ H+ + NO2- If solution is diluted, reaction goes to products If diluted to 2X the volume: So Q<K, products favored

21. Le Châtelier’s Principle Stess Number Two: Dilution – Molecular Scale View Concentrated Solution Diluted Solution – dissociation allows ions to fill more space H+ NO2- H+ NO2- H+ NO2- H+ NO2- H+ H+ H+ NO2- H+ NO2- H+ NO2- H+ NO2- NO2- NO2-

22. Le Châtelier’s Principle Stress Number 3: Temperature If ΔH>0, as T increases, products favored If ΔH<0, as T increases, reactants favored Easiest to remember by considering heat a reactant or product Example: OH- + H+↔ H2O(l) + heat Increase in T

23. Some Le Chatelier’s Principle Examples Looking at the reaction below, that is initially at equilibrium, AgCl(s) ↔ Ag+(aq) + Cl-(aq) (ΔH°>0) determine the direction (toward products or reactants) each of the following changes will result in increasing the temperature addition of water addition of AgCl(s) addition of NaCl