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30 Cards in this Set
- Front
- Back
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amphoteric
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acts as either acid or base
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define Ka and relate to pKa
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Ka = equilibrium constant for acidic dissociation
-intrinsic property of acid pKa = -log(Ka) Ka = 10^(-pKa) |
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strong/weak/very weak acid by pKa and Ka
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*substitute in base, Kb, pKb
strong acid: Ka >> 1 pKa < 0 weak acid: 1 > Ka> 10^(-14) 0 < pKa < 14 very weak acid: Ka < 10^(-14) pKa. > 14 *weaker than H2O |
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relate acid and conjugate base strength
(qualitative and quantitative) |
pKa (acid)+ pKb (conjugate base) = 14
strong acid/base had weak conjugate weak acid/base usually have weak conjugates |
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heterolytic vs. homolytic cleavage
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ionic vs. covenant bond breaking
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oxyacid (relative strength)
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*acidic H bonded to oxygen
*formed by hydration of non-metal oxide CO2 + H2O → H2CO3 1st : resonances electron withdrawal from central atom puts partial positive on acidic oxygen, increasing acidity "excess oxygen count" = #O - #H = # pi bond *each added O decreases pKa by 5 2nd: inductive effect: (if excess O is same) more electronegative central atom will be more acidic |
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normality (N)
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molar equivalents of H+: used to describe polyprotic acids
N for 1 M diprotic acid is 2 N |
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metal oxides and metal hydroxides
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metal oxide (lewis base)
CaO + H2O → Ca(OH)2 metal hydroxide (bronsted base) Ca(OH)2 |
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non-metal oxides and non-metal hydrides
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non-metal oxide (lewis acid)
SO3 + H2O → H2SO4 non-metal hydride (brownsted acid) H2SO4 |
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relative acidity strength (haloacids)
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compare in period:
more electronegative = more acidic compare in group: larger ionic radius = more acidic |
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neutralization of lewis acid + lewis base
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non-metal oxide + metal oxide→ neutral salt
SO3 + CaO → CaSO4 (s) |
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equations for pH of
1) either strong acid / base 2) either weak acid / base 3) both weak acid + weak base |
1) dissociates / hydrolizes water completely
[H+] = mol acid 2) pH = 1/2 pKa - 1/2 log (HA) pOH = 1/2 pKb -1/2 log (A-) 3) pH = pKa + log (base/acid) *Henderson-Hasselbalch *can use M or mols *addition of H2O doesn't affect ratio of acid to base so dilution doesn't affect pH |
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relate pH and pOH
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pH + pOH = 14
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approximate pka of and physiological preference
carboxylic acid + carboxylate amine + alkyl ammonium H2PO4(-) + HPO4 (2-) carbonate + bicarbonate (phenol) |
phenol pKa 9.5 -10.5
**pKa in lower range when R-group amino acid |
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diprotic and triprotic acid pKa / pKb relations
amino acids (pKa) |
pKa1 + pKb3 = 14
pKa2 + pKb2 = 14 pKa3 + pKb1 = 14 *for a given pH, pay attention to which values are applicable -amino acids with acid side-chains are triprotic -amino acids with neutral side-chains are diprotic (pKa for c-terminus = 2-3) (pKa for n-terminus (ammonium)= 9-10) |
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pH estimates of polyprotic acid
-half equivalence point -equivalence point |
1st half equivalence point:
pH = pKa1 1st equivalence point: pH = (pKa1 + pKa2) / 2 *this is not affected by initial concentrations |
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pH at half equivalence and equivalence point for weak acid/base (quantitative)
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half equivalence: pH = pKa
(*HA = A- in HH equation) equivalence point: pH = (pKa (acid) + pKa (titrant)) / 2 |
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pH of equiv point (qualitative)
-weak acid -weak base -strong acid/base |
pH at equiv point for weak acid will be >7
pH at equiv point for weak base will be <7 pH at equiv point for strong base/acid will be =7 *consider conjugates |
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explain how buffer work
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addition of OH- or H+ will convert weak acid or base to its conjugate which doesn't affect pH greatly
there is plenty of HA to A- and the ratio is close to 1 so little change in henderson-hasselbalch |
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2 ways to make a buffer
(pH considerations) |
1) combine equal molar proportions of weak acid and its weak conjugate base (salt)
2) titrate a weak acid/base with 1/2 equivalents of strong base/acid (half-equivalence point) pick acid whose pKa is (+/-) 1 from desired pH *carboxylic acid for pH 2-5 *amine for pH 8-11 *to work: acid:base ratio must not exceed 10:1 or 1 pH unit |
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respitatory acidosis
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retention of blood CO2
blood pH drops |
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respiratory alkalosis
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loss of blood CO2
increase in blood pH |
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metabolic acidosis
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loss of HCO3(-)
blood pH drops |
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metabolic alkalosis
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loss of H3O+
blood pH raises |
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shape of strong-strong titration curve and explanations
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starts close to pH =0 with no initial changes in pH
*concentration of H+ is high so it takes a lot of titrant to make 10-fold ratio change large pH vs. titrant slope at equivalence point *H+ concentration is low so 10-fold difference takes little titrant |
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shape of weak-strong titration curve and explanations
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-starting pH = (pKa - log(HA)) / 2
-"lip-o-weakness" initially there is a large slope because no buffering *larger lip for weaker acid -small pH change at and around half-titration point (pH = pKa) because of buffering (HA : A- ratio changes slowest) -sharp slope at equivalence point (lack of buffering) pH = (pKa acid + pH titrant) / 2 *higher pH for weaker acid -pH then dictated by titrant |
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Altering concentration of strong acid and weak acid (titrations)
*match with molarity of titrant |
strong: different starting pH
same equiv point at 7 weak: different pH at equiv. point, (same amount of titrant to reach) -different starting pH -same half-equiv point pKa = pH |
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-picking an indicator to determine equivalence point
-indicator range |
pH of equiv. point (+/-) 1 pKa of indicator
*pH equiv point = (pKa acid + pH titrant) / 2 *less particular for strong-strong because of large pH range of equivalence point indicators span 3 pH units with 3 colors base : acid 10 : 1 color A 1 : 1 mix A + B 1 : 10 color B |
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log (1/x) =
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-log(x)
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effect of diluting a buffer
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none
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