# Chapter 17 buffers- resist changes in pH by neutralizing added acid or base -acid will neutralize added OH - (base) and base will neutralize added H +

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Chapter 17 buffers- 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 CH 3 COOH and NaCH 3 COO -both have a common acetate ion #1 NaCH 3 COO(aq)  Na + (aq) + CH 3 COO - (aq) #2 CH 3 COOH(aq) ⇌ H + (aq) + CH 3 COO - (aq)

-when mixed, acetate ion from #1 causes a shift to left in #2 b/c adding to product side, dec [H + ] #1 NaCH 3 COO(aq)  Na + (aq) + CH 3 COO - (aq) #2 CH 3 COOH(aq) ⇌ H + (aq) + CH 3 COO - (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 K a to find [H + ] and pH K a = ([H + ][A - ])/[HA] *if initial [ ] of acid and its CB are 10 2 or 10 3 times > K a you can neglect the x value (change)

Can also use: Henderson-Hasselbalch equation: pH = pK a + log([base]/[acid]) -base and acid are [ ] of conj acid-base pair -when [base] = [acid] pH = pK a *can only be used for buffers and when K a 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 = pK a -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 = pK a ± 1

Ex: -a buffering system with pK a = 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= pK a

Addition of Strong Acids or Bases to Buffers -if SB is added: OH - (aq) + HX(aq)  H 2 O(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 K a 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: BaSO 4 (s) BaSO 4 (s) ⇌ Ba 2+ (aq) + SO 4 2- (aq) -to determine solubility (saturated solution): solubility product constant  K sp = [ions] coeff Ex: K sp = [Ba 2+ ][SO 4 2- ] *the smaller the K sp, the lower the solubility

Ex: Write the K sp expression for: 1)calcium fluoride CaF 2 (s) ⇌ Ca 2+ (aq) + 2F - (aq) K sp = [Ca 2+ ][F - ] 2 2)silver sulfate Ag 2 SO 4 (s) ⇌ 2Ag + (aq) + SO 4 2- (aq) K sp = [Ag + ] 2 [SO 4 2- ]

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: Aℓ 3+, Cr 3+, Zn 2+, Sn 2+

Precipitation and Separation of Ions -if two ionic compounds are mixed, a precipitate will form if product of initial ion [ ] > K sp -use Q (reaction quotient) *if Q > K sp, precip occurs, dec ion [ ] until Q = K sp *if Q = K sp, equilibrium exists (saturated solution) *if Q < K sp, solid dissolves, inc ion [ ] until Q = K sp

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