ACIDS AND BASES
acid when dissolved in H2O, it increases the concentration of H+(aq)
strong acid completely dissociates into ions in aqueous solution
examples: Hydrochloric acid HCl, Hydrobromic acid HBr, Hydroiodic acid HI, Perchloric acid HClO 4,
Nitric acid HNO3, Sulfuric acid H2SO4)
weak acid when dissolved in water, only a small percentage of the molecules dissociate into ions; most are
organic acids
examples: HF, H2CO3, H3PO4, H3AsO4, HClO3, HClO2, HClO
base when dissolved in H2O, it increases the concentration of OHGAS-FORMING REACTIONS
metal carbonate + acid  metal salt + H2CO3
 metal salt + CO2 (g) + H2O (l)
example: Na2CO3 + 2HCl  2NaCl + CO2 (g) + H2O (l)
metal sulfite + acid  metal salt + H2SO3
 metal salt + SO2 (g) + H2O (l)
example: Na2SO3 + 2HCl  2NaCl + CO2 (g) + H2O (l)
metal sulfide + acid  metal salt + H2S(g)
example: Na2S + 2HCl  2NaCl + CO2 (g) + H2O (l)
ammonium salt + strong base  metal salt + NH3 (g) + H2O (l)
example: NH4Cl + NaOH  NaCl + NH3 (g) + H2O (l)
ACID-BASE THEORIES
Arrhenius acids increase H3O+ concentration, bases increase OH- concentration
example: HCl + H2O  H3O+ + ClKOH  K+ + OHBrønsted-Lowry acids donate protons, bases accept protons
example: CH3COOH + Na(HCO)3  H2CO3 + CH3OONa
 H2O + CO2 (g) + CH3OONa
Lewis acids accept electron pairs, bases donate electron pairs
example: Ammonia (base) Boron trichloride (acid)
CONJUGATE ACID-BASE PAIRS
H2O (ACID) + H2O (BASE)  H3
O+
(CONJUGATE ACID)
+ OH-
(CONJUGATE BASE)
H
|
H–N:
|
H
+
..
: Cl :
| ..
B – Cl :
. .
|
: Cl :
..
H : Cl :
|
| ..
 H – N – B – Cl :
. .
|
|
H : Cl :
. .
. .
K  [ H 3O  ][OH  ]
*example of autoionization (ionization with another identical molecule) and the amphoteric/amphiprotic property (acts as both acid and base)
More examples of the amphiprotic property:
H3PO4  H2PO4-  HPO42-  PO43H2CO3  HCO3-  CO32Use ICE charts to determine position of equilibrium from Ka or Kb:
CH3COOH(aq)  CH3COO-(aq) + H+(aq)
(given) Ka = 1.8x10-5
[H+] if 0.5M acetic acid is used:
CH3COOH(aq)
0.5M
Initial
-x
Change
Equilibrium
0.5 – x  0.5 (CHECK!)
Ka = 1.8x10-5 =
[ H  ][CH 3COO  ] x 2

[CH 3COOH ]
0 .5
 x = 3x10-3M = [H+] Check: 0.5 – 0.003 = 0.497  under 10%
pH = -log[H+] = -log[3x10-3M] = 2.523 (acidic)
Percent dissociation =
[ H  ]eq
[ H  ]0
 100
=
CH3COO-(aq)
0
+x
x
3x10 3 M
 100  0.60%
0.50M
H+(aq)
0
+x
x
WEAK ACIDS AND BASES
HA(aq) + H2O(l)  H3O+(aq) + A-(aq)
B(aq) + H2O(l)  BH+(aq) + OH-(aq)
Ka =
[ H 3O  ][ A ]
for weak acid: Ka < 1
[ HA]
[ BH  ][OH  ]
Kb =
for weak base: Kb < 1
[ B]
-equilibrium lies toward the side of the reaction having the weaker acid and base
Acid Used
Strong
Strong
Weak
Weak
Kw = KaKb = 1x10-14
neutral: [H3O+] = [OH-] = 1x10-7M
acidic: [H3O+] > [OH-]
basic: [H3O+] < [OH-]
Base Used
Strong
Weak
Strong
Weak
Salt Type
Neutral
Acidic
Basic
Mixed
Salt Solution Will Be
pH  7
pH < 7
pH > 7
Depends on Ka and Kb
pH = -log[H3O+]
 [H3O+] = 10-pH
pOH = -log[OH-]  [OH-] = 10-pOH
pH + pOH = 14
Monoprotic acids contain one acidic hydrogen (HCl, HNO3, HCN, HBrO3)
Polyprotic acids contain more than one acidic hydrogen. A diprotic acid will have two hydrogens, a triprotic acid will have three
hydrogens. (H2SO4, H3PO4, H2S, H2CO3) A titration curve of a polyprotic acid will show a steep rise (endpoint) for each of the
hydrogens in the acid.
TITRATION CURVES
Titration of a strong monoprotic acid (HCl) with a strong base (NaOH):
Before addition
- [H3O+] = [HCl]
As strong base is added
- [OH-] = [NaOH]
- moles H3O+ = moles of original H3O+ - moles of added OH- volume VT = VA + VB
- [H3O+] = moles H3O+/Vtotal
At half-point
- Ka = [H3O+]
- pH = pKa
- [HA] = [A-]
At equivalence point
- moles acid = moles base
- CAVA = CBVB
- pH = 7
After equivalence point
- [OH-] = (moles of added OH- - moles of original H3O+)/VT
- [H3O+] = 1 x 10-14/[OH-]
Strong monoprotic acid titrated with strong base
http://www.bcpl.net/~kdrews/titration/titrationcurve.html
Weak diprotic acid titrated with strong base
- Strong acid titrated with strong base: starts out with very slow or
moderate change in pH as base is added to acid
- Weak acid titrated with strong base: starts out with quicker change in pH
than above situation
- At endpoint, pH changes more dramatically (vertical line). After endpoint,
pH rate of change diminishes.
•Standardization is used to determine the concentration of an analytical reagent.
A sample of a pure solid acid or base is weighed, then titrated with a solution of
base or acid to be standardized. • NAVACA= NBVBCB
•Common Ion Effect The limiting of acid (or base) ionization caused by addition of tits conjugate base (or conjugate acid)
HA + H2O  H3O+ + A- -Weak acid HA ionizes to produce less H3O+ in the presence of A-, therefore increasing pH
• Buffers a solution that resists a change in pH when hydroxide or hydronium ions are added; to make a buffer solution, large quantities of both a
weak acid and conjugate base (or weak base and conjugate acid)
• Henderson-Hasselbach equation: pH = pKa + log ([A-]/[HA])