chapter 17 buffers- resist changes in ph by neutralizing added acid or base -acid will neutralize...
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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
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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)
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-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
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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)
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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 [ ]
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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
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-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
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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
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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
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-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
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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
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Strong Acid-Strong Base Titrations
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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)
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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
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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
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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
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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
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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
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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-]
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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
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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
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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
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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+
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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