Chem. 31 – 10/2 Lecture

Announcements I Exam 1 – on Oct. 4th Water Hardness Lab This Wednesday Will Cover Ch. 1, 3, 4, and 6-1 Review of topics after covering the last part of Ch. 4 and Ch. 6-1 Help Session Today in Sequoia 452; 5:30 to 6:30 PM Water Hardness Lab Due today Website has updated Appendix III (with report forms)

Announcements II Statistical Calculations Lab Today’s Lecture Data set is posted Due date is now 10/11 Today’s Lecture Chapter 4 Least Squares Regression (last bits) Chapter 6 – Material on Exam 1 Introduction to equilibria and equation manipulation Review of Exam 1 material Chapter 6 – Material not on Exam 1

Calibration Question A student is measuring the concentrations of caffeine in drinks using an instrument. She calibrates the instruments using standards ranging from 25 to 500 mg/L. The calibration line is: Response = 7.21*(Conc.) – 47 The response for caffeine in tea and in espresso are 1288 and 9841, respectively. What are the caffeine concentrations? Are these values reliable? If not reliable, how could the measurement be improved?

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)

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

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

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

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+

Review for Exam (just Ch. 1 to 4 topics – Ch. 6 covered later) Know the following (from Ch. 1) Common base units (m, kg, s, mol, K) + common multipliers (nano to kilo) How to convert between different units* Definitions of main concentration units (M, weight fractions including % and ppm, and mass/volume) How to convert between concentration units* Equipment and steps to make solutions of known concentration + calculations for preparation* How to do stoichiometry problems (involving solids or solutions)* Note: *means need quantitative knowledge 10

Review for Exam – cont. Know the following (from Ch. 3) Rules for significant figures (including for calculations with +, -, *, or / and when uncertainties are given)* Definitions for: systematic and random error, accuracy and precision, uncertainty, relative error and relative uncertainty How to do propagation of uncertainty problems (+, -, *, /, exponent, and mixed operations) and to convert between absolute and relative uncertainty* Know the following (from Ch. 4) What a Gaussian distribution represents How to calculate mean values and standard deviations (can use calculators)* 11

Review for Exam Know the following (from Ch. 4 – cont.) The differences between populations and samples How to calculate Z values* How to use Table 4-1 and Z values to calculate probabilities between limits* How to determine confidence intervals* (t- and Z-based) + factors which influence confidence intervals Difference between Z and t based confidence intervals (lecture only) How to calculate a standard deviation of a mean* What a confidence intervals tell you 12

Review for Exam – Ch. 4 (cont.) Know the following (from Ch. 4 – cont.) How to perform an F-Test* How to perform a case 1 or case 3 t-Test* How to recognize and select a proper test (3 t tests, F test and Grubbs test) How to deal with poor data points (including use of Grubbs test)* How method of least squares works (qualitatively) Steps to the calibration process Assumptions required for least squares analysis How to determine concentrations of unknowns* + limitations in this

Review for Exam Chapter 6 Be able to write equilibrium equations from given equilibrium reactions Manipulate equilibrium reactions/equations*

Equations given on Exam 1 Propagation of uncertainty Basic equations given for + and - for * and / and for exponents Standard Deviation Equation Equation for calculation of standard deviation F-test, case 3 t-test, and Grubb’s test equations

Thermodynamics ΔH, change in enthalpy, 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, change in entropy, is related to disorder of system - If the final system is “more random” than initial system, ΔS > 0

Thermodynamics Entropy A macroscopic analogy to entropy would be to have a box of 50 ping pong balls with half white and half black Even if placed on two separate halves of the box, if the box were shaken to mix the balls, roughly half of each color would be expected in each half leading to a positive DS v v initial state final state 17

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

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