1. Buffers Made Easy

2. Common ion effect Common Ion: the presence of an ion which appears in both the acid (or base) and a salt in the solution Common ion effect: The shift that occurs because of the addition of an ion already involved in the equilibrium reaction.

4. Buffered Solutions A solution that resists a change in it’s pH when either hydroxide ions or protons(hydrogen) ions are added. Consists of a WEAK acid and it’s SALT Weak acids or bases and a common ion.

5. How does it work If OH - are added to the system it moves right. Not allowing OH - to accumulate not changing the pH If H + are added it will proceed in the reverse to produce HA not allowing the pH to change

6. Look at the Eq expression The Eq concentration of H + is determined by the ratio of [HA]/[A - ].

7. Bases are the opposite You will use Kb

8. Hendersen- Hasselbach

9. Simpler method One equation: Work buffers as an equilbrium problem Go back to the reaction HA  H + + A - If acid is added to the buffer, simply add acid to the numerator AND subtract the same quantity from the base since it was self-sacrificing and neutralized the acid. If base is added, simply add the base to the denominator and subtract from the numerator. Add or subtract in moles NOT molarity. Moles = molarity x volume

10. When equal concentrations (or moles) of Acid and Base are present [which occurs at the. equivalence point of a titration] the ratio of acid to base equals ONE and therefore, the pH = pKa. IF you are asked to construct a buffer of a specific pH and given a table of Ka’s, choose a Ka with an exponent close to the desired pH and use equal concentrations of the acid and base

13. RECAP Go to MOLES. Moles = M x V Add to what ever is being added and subtract from the other.

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

15. 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! No equilibrium

16. Weak Acid-Strong 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

17. Strong Acid-Weak 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)

18. Acid-Base Indicators 16.5

19. pH 16.5

20. 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.

21. 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)

22. 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.

23. 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.

24. 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

25. 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.