Transcript Document

ACIDS
1) Sour taste:
Lemon Juice – Citric acid.
Vinegar – Acetic Acid.
Stomach ulcers are aggravated by
hydrochloric acid. HCl
2) Dissolve active metals, usually liberating H2.
3) Corrosive – dissolve compounds that are otherwise hard
to dissolve.
Examples:
Precious metals such as gold (Au) dissolve in HNO3 +
HCl (aqua regia).
Hard water deposits dissolve in vinegar.
4) Turn litmus paper RED (low pH).
BASES
1) Bitter taste.
2) Dissolve oil and grease.
Drano and lye soap contain NaOH.
Breaks ester and amide bonds
3) Slippery to the touch – dissolves hair and skin.
e.g., soap: Na+ -OOC(CH2)16CH3
4) React with many metal ions to form precipitates.
Mg2+ + 2OH-  Mg(OH)2
Example:
Hard water (=Ca2+, Mg2+) + soap  White precipitate.
(bathtub rings)
5) Turn litmus paper BLUE (high pH)
ARRHENIUS ACIDS AND BASES
Arrhenius ACID:
Any compound that releases H+ when dissolved in H2O.
Example:
HCl(g)  H+(aq) + Cl(aq)
Arrhenius BASE:
Any compound that releases OH- when dissolved in H2O.
Example:
KOH(s) + H2O (l)  K+(aq) + OH(aq)
BRØNSTED - LOWRY
ACIDS AND BASES
BrØnsted ACID:
Any compound capable of donating a H+ ion.
Example:
HCl(g)  H+(aq) + Cl(aq)
BrØnsted BASE:
Any compound capable of accepting a H+ ion.
Example:
NH3(g) + H2O(l)  NH4+(aq) + OH(aq)
WATER
Water electrolyzes slightly to produce H+ and OH- reversibly.
H2O
H+ + OH-
Autoionization of water
Kw = [H+][OH-] = 1.0 x 10-14 at 25oC
For pure water, [H+] = [OH-] = 10-7, so pH =7
Kw is constant even when [H+] and [OH-] are not equal
Calculate [H+] in a 0.05 M Ca(OH)2 solution
pH scale
pH = -log10[H+]
(low pH = acidic)
pH + pOH = -log10[H+] + -log10[OH-] = 14
Measuring pH
Most accurate method to measure pH is to use a pH meter.
However, certain dyes change color as pH changes. These
are indicators.
HIn = H+ + InIndicators are less precise than pH meters.
Many indicators do not have a sharp color change as a
function of pH.
Which bulbs light up?
Solution
Distilled water
Tap water
NaCl(aq)
1 M HCl (aq)
1 M CH3COOH (aq)
Sugar
CH3OH
Strong, weak, or non-electrolyte?
STRONG ACIDS
Strong Acids dissociate completely when dissolved in water to
form H+ and the corresponding BrØnsted base.
HA  H+(aq) + A-(aq)
Strong acids are strong electrolytes:
COMPLETE dissociation into ions
[H+]final = [HA]initial = CHA
(If the analytical concentration, CHA, is less than 10-6
M then the autoionization of water needs to be
taken into account.)
WEAK ACIDS
When dissolved in water weak acids only partially
dissociate to form H+ and the corresponding base.
HA (aq)
H+ (aq) + A- (aq)
Weak acids are weak electrolytes:
PARTIAL dissociation into ions
[H+]final < [HA]initial
Examples:
CH3CO2H
HF
H3PO4
Acid Dissociation Constant (Ka) <<1
[H  ][A ]
Ka 
[HA]
What is the [H+] of 0.10 M HI?
What is the [H+] of 0.10 M acetic acid?
1.
2.
3.
4.
1.8 x 10-5 M
4.2 x 10-3 M
1.8 x 10-6 M
1.3 x 10-5 M
What is the pH?
What is the % dissociation?
Ka = 1.8 x10-5
% Dissociation of CH3CO2H
CHA(M) [H+](M)
% Dissoc.
10
0.013
0.13
1
0.004
0.4
0.1
0.0013
1.3
0.01
0.0004
4.0
0.001
0.00013
13.4
OXYACIDS
Many Brønsted acids consist of a central atom with
several attached oxygen atoms. These are called
oxyacids.
Acid strength increases with increasing oxidation
number of the central atom:
HOClO3 > HOClO2 > HOClO > HOCl
General rule for uncharged oxyacids HxEOy:
If y-x > 2 then strong (H2SO4, HNO3,…)
If
< 2 then weak (H2CO3, HBrO, HNO2,…)
Increasing electronegativity of the central atom
increases acid strength
HOCl > HOBr > HOI
POLYPROTIC ACIDS
Polyprotic acids are capable of donating more than one proton.
Contain more than one ionizable proton.
The Ka always gets smaller with each ionization
Examples:
H2CO3(aq)
HCO3-(aq)
H+ (aq) + HCO3-(aq)
H+ (aq) + CO32- (aq)
Ka = 4.3 x 10-7
Ka = 5.6 x 10-11
H3PO4 (aq)
H2PO4-(aq)
HPO42-(aq)
H+(aq) + H2 PO4- (aq)
H+ (aq) + HPO42- (aq)
H+(aq) + PO43-(aq)
Ka = 7.5 x 10-3
Ka = 6.2 x 10-8
Ka = 4.2 x 10-13
What are the concentrations of H+, HCO3-, CO32- in
1 x 10-3 M H2CO3?
Strong Acids
Which one of the following are not strong
acids?
1. HNO3
2. HF
3. HClO3
4. HClO4
5. HOBr
6. HBr
7. HPO428. H2SO3
STRONG BASES
Group I and II hydroxides (except Mg and Be).
Arrhenius bases donate OH-.
Brønsted bases accept H+
Examples:
NaOH, KOH, Ca(OH)2
KOH + H2O  K+ (aq) + OH- (aq)
Strong bases are strong electrolytes.
[OH-] = Cbase
WEAK BASES
When dissolved in water weak bases only partially react to form
OH and the corresponding BrØnsted acid.
B + H2O
HB+(aq) + OH(aq)
Weak bases are weak electrolytes: [OH-] < Cbase
Weak bases can be neutral
Example: NH3, amines
NH3 + H2O = NH4+(aq) + OH(aq)
Or Anions (any ion derived from a weak base) Example: F, NO2,
CH3COO
F(aq) + H2O = HF(aq) + OH(aq)
Base Dissociation Constant Kb << 1
[HB ][OH ]
Kb 
[B]
What is the pH of 0.1 M NH3?
Kb = 1.8 x 10-5
1.
2.
3.
4.
5.
2.87
4.74
7.00
9.25
11.1
CONJUGATE ACID BASE PAIRS
Differ only by the presence or absence of a proton (H+).
Conjugate Acid = Conjugate Base + H+
Examples:
H 3 O+ / H 2 O
H2O / OH
NH4+ / NH3
HCl / Cl
Ka x Kb = constant = 1 x 10-14
• The conjugate of a weak acid is a weak base (and vice versa)
• The conjugate of a strong acid is a spectator ion (example: Cl is the
conjugate base of HCl).
• The conjugate acid of OH (strong base) is water.
When we add two reactions together, we multiply
their equilibrium constants.
For conjugate acid-base pairs:
Ka x Kb = Kw = 1 x 1014
Larger Ka means smaller Kb
The stronger the acid, weaker its conjugate base
pKa = log Ka
pKb = log Kb
log ( Ka x Kb ) = log Kw = 14
log Ka  log Kb = 14
pKa + pKb = 14
Weaker acid
H-F +
OH-
stronger conjugate base
F-
+
Weaker acid
Ka = 10-14
Stronger acid
6.9 x 10-4
Stronger base
H2O
Weaker base
Kb = 1.4 x 10-11
ACETIC ACID
Acid:
CH3COOH
Base: CH3COO + H2O
H+ + CH3COO
CH3COOH + OH-
----------------------------------------------H2O
H+ + OHKa =
Kb =
Kw = [H+][OH-] = Ka x Kb = 1 x 10-14
pKa + pKb = 14
Hydrolysis: when a cation or anion reacts with
H2O to form H+(aq) or OH(aq)
Will a salt be acidic or basic?
1. Salt derived from a strong acid and a strong base
Neutral solution
(pH = 7)
Example: NaCl (from NaOH and HCl)
2. Salt derived form a weak acid and a strong base
Basic solution (pH > 7)
Examples: NaClO (NaOH and HClO)
ClO (aq) + H2O
HClO (aq) + OH(aq)
(CH3COO)2Ba (Ba(OH)2 and CH3COOH)
CH3COO(aq) + H2O
CH3COOH(aq) +OH(aq)
3. Salt derived from a strong acid and a weak
base
Acidic solution (pH <7)
Example: NH4Cl (NH3 and HCl)
NH4+ + H2O
NH3 + H3O+
4. Salt derived form a weak acid and a weak
base
pH depends on acid/base involved
Example: NH4CN (NH4+ and CN)
What is the pH of 0.02 M KN3
Ka (HN3) = 1.9 x 10-5
1.
2.
3.
4.
5.
3.21
5.49
7.00
8.51
10.8
LEWIS ACIDS
Any substance that can accept a pair of electrons.
• Small cations
• Molecules with unfilled octets
e.g. H+, BF3
Examples of Lewis Acids:
Highly charged transition metal cations, e.g. Fe3+, Fe2+, Co3+
Group III cations (Al3+, Ga3+) and compounds (AlCl3)
Smaller group II cations: Be2+ and Mg2+
LEWIS BASES
Any substance that can donate a pair of electrons.
• Has lone pair electrons
• May be neutral or anionic.
Examples: NH3, OH-, Brønsted bases, H2O, Cl-
LEWIS CATIONS
To compare acidity of Lewis acids, first compare charge. If
charge is the same then compare size.
Charge/Size Ratios
Metal Ion
Charge/Ionic radius (Å)
Na+
1.0
Li+
1.5
Ca2+
2.1
Mg2+
3.1
Zn2+
2.7
Cu2+
2.8
Al3+
6.7
Cr3+
4.8
Fe3+
4.7
HYDRATION
Metal ions attract the lone pairs on the oxygen in water
molecules. This is a Lewis acid – Lewis base reaction.
Mz+
H
-
:O
+
H
Hydrated metal ions are acidic. Acidity increases with
increasing charge/size ratio of the metal ions.
Hydrolysis is a reaction that dissociates water:
M(H2O)nz+
Fe(H2O)63+
M(H2O)n-1(OH)(z-1)+ + H+
Fe(H2O)5(OH)2+ + H+ (Ka=6.7 x 10-3)
ACIDS AND BASES SO FAR
1) Arrhenius, Brønsted, and Lewis definitions
2) pH, pOH
3) Acid and Base Dissociation Constants – Ka and Kb
4) [H+] [OH-] = 1 x 10-14 = Ka x Kb
5) pH and % ionization calcn for strong and weak acids/bases
6) Conjugate Acid-Base Pairs:
Arrhenius
Bronsted-Lowry
Lewis
7) Salts – Hydrolysis
8) Structure Related to Acid-Base Properties (Oxyacids)
YOU SHOULD KNOW
GIVEN
pH
[H+] or [OH-]
List of acids
List of pKa’s or Ka’s
Ka or pKa and [HX]
pH and [HX]
FIND
[H+], [OH-], pOH
pH
Weaker /Stronger
Weaker /Stronger
pH, [H+], [OH-]
Ka
Recall that a small Ka means high pKa, and both mean weak
acid and not much dissociation.
Acid/Base SALTS Review
1) Which one of the following salts would have a
basic aqueous solution?
1. KF
2. Al(NO3)3
3. NaI
4. NH4Br
2) Arrange the following in the order of increasing
base strength:
N3-
NO3-
HPO42-
CN-
3) Which of the following cannot act as a Lewis
base?
1. Cl4. NH3
2. OH5. H+
3. CN-