Chem. 31 – 9/10 Lecture

1. Chem. 31 – 9/10 Lecture

2. Announcements • Due today: • corrected diagnostic quiz • HW1.1 (only additional problem 1.1) • Quiz today • Lab Procedures Quiz in Lab Thursday and Monday • Today’s Lecture • Stoichiometry problems • Titrations • Uncertainty and Error (definitions and signficant figures)

3. Stoichiometry • Stoichiometry refers to ratios between moles of reactants and products in chemical reactions • The ratio of moles of reactants and products is equal to the ratio of their stoichiometric coefficients Example: aA + bB ↔ cC + dD Moles A/moles B = a/b

4. Stoichiometry • Example problem: How many moles of H2O2 are needed to completely react with 25 mL of 0.80 M MnO4-? Reaction: 5H2O2(aq) + 2MnO4- +6H+↔ 2Mn2+ + 5O2(g) + 8H2O(l)

5. Stoichiometry • Remember: there are two (common) ways to deliver a known amount (moles) of a reagent: • Mass (using formula weight) • Volume (if molarity is known) • Titrations = A practical way of using stoichiometry with precise measure of added volume

6. TitrationsDefinitions • Titrant: • Reagent solution added out of buret (concentration usually known) • Analyte solution: • Solution containing analyte • Equivalence Point: • point where ratio of moles of titrant to moles of analyte is equal to the stoichiometric ratio titrant analyte solution for: Al3+ + 3C2O42-→ Al(C2O4)33- n(Al3+)/n(C2O42-) = 3/1 at equivalence pt.

7. TitrationsPractical Requirements • The equilibrium constant must be large • Precision of titration will depend on size of K and concentration of analyte • Typically K ~ 106 is marginal, K > 1010 is better • The reaction must be fast • It must be possible to “observe” the equivalence point • observed equivalence point = end point

8. Standardization vs. Analyte Titrations To accurately determine an analyte’s concentration, the titrant concentration must be well known This can be done by preparing a primary standard (high purity standard) Alternatively, the titrant concentration can be determined in a standardization titration (e.g. vs. a known standard) Rationale: many solutions can not be prepared accurately from available standards Example: determination of [H2O2] by titration with MnO4- neither compound is very stable so no primary standard instead, [MnO4-] determined by titration with H2C2O4 in standardization titration then, H2O2 titrated using standardized MnO4- TitrationsOther Definitions

9. TitrationsOther Definitions • Direct vs. Back Titration • In a direct titration, the titrant added slowly to the analyte until reaching an end point • In a back titration, a reagent is added to the analyte in excess, and then that reagent is titrated to an end point • Often done to get sharper endpoint

10. TitrationsBack Titration Example Titration to determine moles of Na2CO3 in a sample: First, direct titration: Na2CO3 + 2HCl → H2CO3 + NaCl (we will do following AA lab) HCl not that sharp Na2CO3

11. TitrationsBack Titration Example NaOH HCl Titration to determine moles of Na2CO3 in a sample: Now, via back titration: excess HCl added to sample Na2CO3 + HCl → NaCl + H2CO3 +heat → NaCl + H2O + CO2(g) After heating only NaCl and excess HCl left Excess HCl titrated with NaOH to NaCl + H2O Na2CO3 Very Sharp Excess HCl

12. TitrationsWhat Makes a Titration Sharp? SHARP TITRATION • A sharp titration has a large slope (absolute value) • Slope at endpoint seen in plot of -log[analyte] vs. V(titrant) • With a sharp titration, errors or uncertainties in V(equivalence point) are small uncertainties in log[analyte] [reactant] at eq. point -Log[analyte] V(eq. pt.) V(titrant) small uncertainty in V results NON-SHARP TITRATION -Log[analyte] V(titrant) larger unc. in V

13. TitrationsSome Questions • List two requirements for a titration to be functional. • In a back titration, what is actually being titrated? (a) analyte b) reagent added c) excess reagent d) secondary reagent) • Why might one want to standardize a prepared solution of 0.1 M NaOH rather than prepare it to exactly 0.100 M? NaOH is a hygroscopic solid that also absorbs CO2.

14. TitrationsBack Titration Example Sulfur dioxide (SO2) in air can be analyzed by trapping in excess aqueous NaOH (see 1). With addition of excess H2O2, it is converted to H2SO4 (see 2), using up additional OH- (see 3). SO2 (g) + OH- (aq) → HSO3- HSO3- + H2O2(aq) → HSO4- + H2O HSO4- + OH- (aq) → SO42- + H2O 208 L of air is trapped in 5.00 mL of1.00 M NaOH. After excess H2O2 is added to complete steps 1 to 3 (above), the remaining NaOH requires 21.0 mL of 0.0710 M HCl. What is the SO2 concentration in mmol/L?

15. Chapter 3 – Error and Uncertainty Error is the difference between measured value and true value or error = measured value – true value Uncertainty Less precise definition The range of possible values that, within some probability, includes the true value

16. Measures of Uncertainty Explicit Uncertainty: Measurement of CO2 in the air: 389 + 3 ppmv The + 3 ppm comes from statistics associated with making multiple measurements (Covered in Chapter 4) Implicit Uncertainty: Use of significant figures (389 has a different meaning than 400 and 389.32)

17. Significant Figures(review of general chem.) Two important quantities to know: Number of significant figures Place of last significant figure Example: 13.06 4 significant figures and last place is hundredths Learn significant figures rules regarding zeros

18. Significant Figures - Review Some Examples (give # of digits and place of last significant digit) 21.0 0.030 320 10.010