1. Acid-Base Equilibria Chapter 16

2. The __________________is the shift in equilibrium caused by the addition of a compound having an __________in common with the______________ substance. The presence of a common ion _____________the ionization of a ________acid or a __________base. Consider mixture of CH 3 COONa (_______electrolyte) and CH 3 COOH (_______ acid). CH 3 COONa (s) Na + (aq) + ) CH 3 COOH (aq) H + (aq) + ) common ion 16.2

3. A ________________is a solution of: 1.A weak ______or a weak _______and 2.The _________of the weak acid or weak base Both must be present! A ____________has the ability to _________changes in ___upon the addition of small amounts of either _______or __________. 16.3 CH 3 COOH (aq) H + (aq) + CH 3 COO - (aq) Consider an equal molar mixture of CH 3 COOH and CH 3 COONa Adding more _______creates a shift left IF enough _________ions are present

4. Which of the following are buffer systems? (a) KF/HF (b) KCl/HCl, (c) Na 2 CO 3 /NaHCO 3 (a) KF is a weak acid and F - is its conjugate base _______________ (b) HCl is a strong acid _______________ (c) CO 3 2- is a weak base and HCO 3 - is it conjugate acid _______________ 16.3

5. What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOOK? HCOOH (aq) H + (aq) + HCOO - (aq) Initial (M) Change (M) Equilibrium (M) 16.2 Mixture of weak acid and conjugate base! K a for HCOOH = 1.8 x 10 -4

6. What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOOK? HCOOH (aq) H + (aq) + HCOO - (aq) Initial (M) Change (M) Equilibrium (M) Common ion effect 16.2 Mixture of weak acid and conjugate base!

7. HCl H + + Cl - HCl + CH 3 COO - CH 3 COOH + Cl - 16.3

8. Calculate the pH of the 0.30 M NH 3 /0.36 M NH 4 Cl buffer system. What is the pH after the addition of 20.0 mL of 0.050 M NaOH to 80.0 mL of the buffer solution? Initial End 16.3 Change 

9. pH = 9.20 Calculate the pH of the 0.30 M NH 3 /0.36 M NH 4 Cl buffer system. What is the pH after the addition of 20.0 mL of 0.050 M NaOH to 80.0 mL of the buffer solution? start (M) end (M) final volume = 80.0 mL + 20.0 mL = 100 mL 16.3

10. = Calculate the pH of the 0.30 M NH 3 /0.36 M NH 4 Cl buffer system. What is the pH after the addition of 20.0 mL of 0.050 M NaOH to 80.0 mL of the buffer solution? [ [[ start (M) end (M) 16.3

11. Chemistry In Action: Maintaining the pH of Blood 16.3

12. Titrations In a ____________a solution of accurately known _____________is added gradually added to another solution of __________concentration until the chemical reaction between the two solutions is complete. ___________________– the point at which the reaction is complete ___________– substance that changes color at the ________ (hopefully close to the equivalence point) Slowly add base to unknown acid UNTIL The indicator changes color (________) 4.7

13. Strong Acid-Strong Base Titrations NaOH (aq) + HCl (aq) H 2 O (l) + NaCl (aq) OH - (aq) + H + (aq) H 2 O (l) 16.4 100% ionization! ___________ _

14. _________Acid-________ Base Titrations CH 3 COOH (aq) + NaOH (aq) CH 3 COONa (aq) + H 2 O (l) CH 3 COOH (aq) + OH - (aq) CH 3 COO - (aq) + H 2 O (l) CH 3 COO - (aq) + H 2 O (l) OH - (aq) + CH 3 COOH (aq) At equivalence point (pH ______ 7): 16.4

15. _________Acid-__________ Base Titrations HCl (aq) + NH 3 (aq) NH 4 Cl (aq) NH 4 + (aq) + H 2 O (l) NH 3 (aq) + H + (aq) At equivalence point (pH _____7): 16.4 H + (aq) + NH 3 (aq) NH 4 Cl (aq)

16. Acid-Base Indicators 16.5

17. The titration curve of a strong acid with a strong base. 16.5

18. Which indicator(s) would you use for a titration of HNO 2 with KOH ? _______acid titrated with _________base. At equivalence point, will have __________________acid. At equivalence point, pH _____ 7 Use __________or__________________ 16.5

19. Finding the Equivalence Point (calculation method) ________Acid vs. ________Base –________% ionized! pH ____No__________! ________Acid vs. _______ Base –______is____________; Need ____for conjugate ________equilibrium _________Acid vs. ________Base –Base is________________; Need ____for conjugate ________equilibrium _________ Acid vs. ______ Base –Depends on the ______of both; could be conjugate______, conjugate ______, or pH ___

20. Exactly 100 mL of 0.10 M HNO 2 are titrated with 100 mL of a 0.10 M NaOH solution. What is the pH at the equivalence point ? HNO 2 (aq) + OH - (aq) NO 2 - (aq) + H 2 O (l) start (moles) end (moles) NO 2 - (aq) + H 2 O (l) OH - (aq) + HNO 2 (aq) Initial (M) Change (M) Equilibrium (M) Final volume = 200 mL K b == =

21. _____________Ion Equilibria and Solubility A __________is an ion containing a central metal cation bonded to one or more molecules or ions. Co 2+ (aq) + 4Cl - (aq) CoCl 4 (aq) 2- Co(H 2 O) 6 2+ CoCl 4 2- 16.10

23. Complex Ion Formation These are usually formed from a transition metal surrounded by ligands (polar molecules or negative ions). As a "rule of thumb" you place twice the number of ligands around an ion as the charge on the ion... example: the dark blue Cu(NH 3 ) 4 2+ (ammonia is used as a test for Cu 2+ ions), and Ag(NH 3 ) 2 +. Memorize the common ligands.

24. Common Ligands LigandsNames used in the ion H2OH2Oaqua NH 3 ammine OH-hydroxy Cl-chloro Br-bromo CN-cyano SCN-thiocyanato (bonded through sulphur) isothiocyanato (bonded through nitrogen)

25. Names Names: ligand first, then cation Examples: –tetraamminecopper(II) ion: Cu(NH 3 ) 4 2+ –diamminesilver(I) ion: Ag(NH 3 ) 2 +. –tetrahydroxyzinc(II) ion: Zn(OH) 4 2- The charge is the sum of the parts (2+) + 4(-1)= -2.

26. When Complexes Form Aluminum also forms complex ions as do some post transitions metals. Ex: Al(H 2 O) 6 3+ Transitional metals, such as Iron, Zinc and Chromium, can form complex ions. The odd complex ion, FeSCN 2+, shows up once in a while Acid-base reactions may change NH 3 into NH 4 + (or vice versa) which will alter its ability to act as a ligand. Visually, a precipitate may go back into solution as a complex ion is formed. For example, Cu 2+ + a little NH 4 OH will form the light blue precipitate, Cu(OH) 2. With excess ammonia, the complex, Cu(NH 3 ) 4 2+, forms. Keywords such as "excess" and "concentrated" of any solution may indicate complex ions. AgNO 3 + HCl forms the white precipitate, AgCl. With excess, concentrated HCl, the complex ion, AgCl 2 -, forms and the solution clears.

27. Coordination Number Total number of bonds from the ligands to the metal atom. Coordination numbers generally range between 2 and 12, with 4 (tetracoordinate) and 6 (hexacoordinate) being the most common.

28. Some Coordination Complexes molecular formula Lewis base/ligand Lewis aciddonor atom coordination number Ag(NH 3 ) 2 + NH 3 Ag + N2 [Zn(CN) 4 ] 2- CN-Zn 2+ C4 [Ni(CN) 4 ] 2- CN-Ni 2+ C4 [PtCl 6 ] 2- Cl-Pt 4+ Cl6 [Ni(NH 3 ) 6 ] 2+ NH 3 Ni 2+ N6