Chapter 17buffers- resist changes in pH by neutralizing added acid or base

-acid will neutralize added OH-(base) and base will neutralize added H+(acid)

-a WA by itself does not contain enough CB to be a buffer

-a WB also does not contain enough CA to be a buffer

-a buffer contains significant amounts of both a WA and its CB or WB and its CA

Ex: a simple buffer can be made by dissolving CH3COOH and NaCH3COO

-both have a common acetate ion

#1 NaCH3COO(aq) Na+(aq) + CH3COO-(aq)

#2 CH3COOH(aq) ⇌ H+(aq) + CH3COO-(aq)

-when mixed, acetate ion from #1 causes a shift to left in #2 b/c adding to product side, dec [H+]

#1 NaCH3COO(aq) Na+(aq) + CH3COO-(aq)

#2 CH3COOH(aq) ⇌ H+(aq) + CH3COO-(aq)

-causes acetic acid to ionize less than it normally would produces higher pH (less acidic)

common-ion effect- weak electro. and strong electro. with common ion in a solution causes weak electro. to ionize less than it would if it were alone

Calculating pH of a Buffer

1. identify equilibrium that is source of [H+] determines pH (the acid)

2. ICE table- be sure to include initial [ ] for acid and its CB (common ion)

3. use Ka to find [H+] and pH

Ka = ([H+][A-])/[HA]

*if initial [ ] of acid and its CB are 102 or 103 times > Ka you can neglect the x value (change)

Can also use:

Henderson-Hasselbalch equation:

pH = pKa + log([base]/[acid])

-base and acid are [ ] of conj acid-base pair

-when [base] = [acid] pH = pKa

*can only be used for buffers and when Ka is small compared to [ ]

buffer capacity- the amount of acid or base the buffer can neutralize before the pH begins to change

-inc with inc [ ] of acid and base used to prepare buffer

-the pH range of any buffer is the pH range which the buffer acts effectively

-buffers are most effective when [ ] of WA and CB are about the same

*remember when [base] = [acid], pH = pKa

-this gives optimal pH of any buffer

-if [ ] of one component is more than 10X(1 pH unit) the other, the buffering action is poor

-effective range for a buffering system is:

pH = pKa ± 1

Ex:

-a buffering system with pKa= 5.0 can be used to prepare a buffer in the range of 4.0-6.0

-amounts of acid and CB can be adjusted to achieve any pH within this range

*most effective at pH=5.0

*b/c pH= pKa

Addition of Strong Acids or Bases to Buffers

-if SB is added:

OH-(aq) + HX(aq) H2O(l) + X-(aq)

*OH- reacts with HX (WA) to produce X-

*[HX] will dec and [X-] will inc

*inc pH slightly

-if strong acid is added:

H+(aq) + X-(aq) HX(aq)

*H+ reacts with X-(CB) to produce HX

*[X-] will dec and [HX] will inc

*pH dec slightly

Acid-Base Titrations

-a base in a buret is added to an acid in small increments (or acid added to base)

pH titration curve- graphs pH vs volume of titrant added

*page 714

equivalence point- moles base = moles acid

Strong Acid-Strong Base Titrations

Titration Curve (finding pH values)

1. initial pH: pH before any base is added

*initial conc of SA ([H+]) = initial pH

2. between initial pH and equivalence pt: as

base is added pH inc slowly and then rapidly near equiv pt.

*pH = conc of excess acid (not yet neutralized)

3. equiv. pt.: mol base = mol acid, leaving only

solution of the salt

*pH = 7.00

4. after equiv pt: pH determined by conc of excess base

Weak Acid-Strong Base Titrations

How does this differ from titration curve of strong-strong?

1. weak acid has higher initial pH than strong

2. pH change in rapid-rise portion is smaller for weak acid than strong

3. pH at equiv. is > 7.00

-page 717 fig 17.9

Titration Curve (finding pH values)

1. initial pH: use Ka to calculate

2. between initial pH and equiv. pt: acid is being neutralized and CB is being formed

*done using ICE table and [ ]

3. equiv. pt.: mol base = mol acid

*pH > 7 b/c anion of salt is a weak base

4. after equiv pt: pH determined by conc of excess base

Titrations of Polyprotic Acids

-occurs in a series of steps

-has multiple equiv. pt. (one for each H+)

-page 720 fig 17.12

Acid-Base Indicators

endpoint- point where indicator changes color (closely approximates equiv pt)

-must make sure the equiv. pt. falls within the color-change interval

-page 721 and 722

Solubility Equilibria

-looking at dissolution of ionic compounds

-heterogeneous reactions

ex: BaSO4(s)

BaSO4(s) ⇌ Ba2+(aq) + SO42-(aq)

-to determine solubility (saturated solution):

solubility product constant Ksp = [ions]coeff

Ex: Ksp = [Ba2+][SO42-]

*the smaller the Ksp, the lower the solubility

Ex: Write the Ksp expression for:

1) calcium fluoride

CaF2(s) ⇌ Ca2+(aq) + 2F-(aq)

Ksp = [Ca2+][F-]2

2) silver sulfate

Ag2SO4(s) ⇌ 2Ag+(aq) + SO42-(aq)

Ksp = [Ag+]2[SO42-]

Factors Affecting Solubility

1. Common-Ion Effect

*solubility of an ionic compound is lower in a solution containing a common ion than in pure water

2. Solubility and pH

-pH of a solution affects the solubility of any substance whose anion is basic

*solubility of a compound with a basic anion (anion of WA) inc. as the solution becomes more acidic

3. Formation of Complex Ions

*involves transition metal ion in solution and a Lewis base

complex ion- contains a central metal ion bound to one or more ligands

ligand- a neutral molecule or ion that acts as a Lewis base with the central metal ion

-forms a coordinate covalent bond (one atom contributes both electrons for a bond)

page 731 and 732

4. Amphoterism

-behave as an acid or a base

amphoteric oxides/hydroxides- soluble in strong acids and bases b/c they can behave as an acid or a base

ex: Aℓ3+, Cr3+, Zn2+, Sn2+

Precipitation and Separation of Ions

-if two ionic compounds are mixed, a precipitate will form if product of initial ion [ ] > Ksp

-use Q (reaction quotient)

*if Q > Ksp, precip occurs, dec ion [ ] until Q = Ksp

*if Q = Ksp, equilibrium exists (saturated solution)

*if Q < Ksp, solid dissolves, inc ion [ ] until Q = Ksp