1. OWL HOMEWORK Announcements

2. CH 3 CO 2 H(aq) + NaOH(aq)  1. CH 3 CO 2 H 2 + (aq) + NaO(aq) 2. CH 3 CO 2 - (aq) + H 2 O(l) + Na + (aq) 3. CH 4 (g) + CO 2 (g) + H 2 O(l)

3. HCN(aq) + NH 3 (aq)  1. NH 4 + (aq) + CN - (aq) 2. H 2 CN + (aq) + NH 2 - (aq) 3. C 2 N 2 (s) + 3 H 2 (g)

4. The pH Scale Quantitative measure of solution acidity Remember solution concentration:   [NaCl]=0.25M means 0.25 moles of NaCl are in 1L of solution

5. How much of a 1.oM HCl solution should I add to neutralize 1.0L of 0.1M NaOH? 1.0L 0.1M NaOH 1. 1.0 mL 2. 10.0 mL 3. 100.0 mL 4. 1.0 L 5. 10.0 L

6. The pH Scale In pure water, some molecules ionize to form H 3 O + and OH - H 2 O + H 2 O  OH – + H 3 O + In acidic and basic solutions, these concentrations are not equal acidic: [H 3 O + ] > [OH – ] basic: [OH – ] > [H 3 O + ] neutral: [H 3 O + ] = [OH – ]

7. The pH Scale pH scale= measure of [H 3 O + ] pH < 7.0 = acidic pH > 7.0 = basic pH = 7.0 = neutral Measure of H 3 O + concentration (moles per liter) in a solution As acidity increases, pH decreases

8. The pH Scale The pH scale is logarithmic 10010 2 log(10 2 ) = 2 1010 1 log(10 1 ) = 1 110 0 log(10 0 ) = 0 0.110 –1 log(10 –1 ) = –1 0.0110 –2 log(10 –2 ) = –2 pH = –log [H 3 O + ] pH = –log [H 3 O + ]

9. Logs- a quick review log N = a10 a = N Example log 4 = 0.60210 0.602 = 4

10. The pH Scale pH = –log [H 3 O + ] pH = –log [H 3 O + ] pH if [H 3 O + ] = 10 –5 ? 10 –9 ? pH if [H 3 O + ] = 10 –5 ? 10 –9 ? Acidic or basic? pH if [H 3 O + ] = 0.000057 M? pH if [H 3 O + ] = 0.000057 M?

11. Finding [H 3 O + ] from pH [H 3 O + ] = 10 -pH or [H 3 O + ] = log -1 (-pH) [H 3 O + ] = 10 -pH or [H 3 O + ] = log -1 (-pH) Finding the inverse log (or log -1 )of a number on your calculator: Finding the inverse log (or log -1 )of a number on your calculator: Enter the number, press the inverse (inv) or shift button, the press the log button (it might be labeled 10 x ) What is [H 3 O + ] if pH = 8.6? What is [H 3 O + ] if pH = 8.6?

12. pH: Quantitative Measure of Acidity Acidity is related to concentration of H + (or H 3 O + )  pH = -log[H 3 O + ]  [H 3 O + ]= 10 -pH = log -1 (-pH) pOH measures basicity  pOH=-log[OH - ]  pH + pOH = 14

13. Precipitation Reactions Solubility of Ionic Compounds dissolving

14. How to determine if an ionic compound is solubl e Identify the two ions Check the “solubility rules”  Soluble ions with no “exceptions” never form precipitates  If one of the ions is insoluble, the compound is insoluble  Make sure to check for “exceptions”  Example: K 2 CO 3  If a reaction product is insoluble, it will form a precipitate

16. Examples Soluble or Insoluble? 1. NaNO 3 2. FeCl 3 3. Fe(OH) 3 4. BaSO 4 5. AgNO 3 6. AgCl

17. More Examples (on your handout- try at home) Soluble or Insoluble? 7. K 3 PO 4 8. Fe 3 (PO 4 ) 2 9. PbCl 2 10. FeSO 4 11. (NH 4 ) 2 S 12. PbS

18. When potassium chromate and barium nitrate react, which product precipitates? 1. BaCrO 4 2. KNO 3 3. Both products are insoluble

19. Precipitation Reactions Net Ionic Equations Pb(NO 3 ) 2 + K 2 CrO 4  ? Pb(NO 3 ) 2 + KI  ? BaCl 2 + KNO 3 