ACIDS1) 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).
BASES1) 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 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)
ARRHENIUS 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)
BRØNSTED - LOWRYACIDS AND BASES
WATERWater 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
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+ + In-
Indicators are less precise than pH meters.Many indicators do not have a sharp color change as a function of pH.
Measuring pH
pH = -log10[H+] (low pH = acidic)
pH + pOH = -log10[H+] + -log10[OH-] = 14
pH scale
Which bulbs light up?
Solution Strong, weak, or non-electrolyte?
Distilled water
Tap water
NaCl(aq)
1 M HCl (aq)
1 M CH3COOH (aq)
Sugar
CH3OH
STRONG ACIDSStrong 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.)
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
WEAK ACIDS
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Ka = [H +][A−][HA]
What is the [H+] of 0.10 M HI?
What is the [H+] of 0.10 M acetic acid? Ka = 1.8 x10-5
• 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?
% 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
OXYACIDSMany 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 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) H+ (aq) + HCO3
-(aq) Ka = 4.3 x 10-7
HCO3-(aq) H+ (aq) + CO3
2- (aq) Ka = 5.6 x 10-11
H3PO4 (aq) H+(aq) + H2 PO4- (aq) Ka = 7.5 x 10-3
H2PO4-(aq) H+ (aq) + HPO4
2- (aq) Ka = 6.2 x 10-8
HPO42-(aq) H+(aq) + PO4
3-(aq) Ka = 4.2 x 10-13
POLYPROTIC ACIDS
What are the concentrations of H+, HCO3-, CO3
2- in 1 x 10-3 M H2CO3?
Which one of the following are not strong acids?
1. HNO3 5. HOBr2. HF 6. HBr3. HClO3 7. HPO4
2-
4. HClO4 8. H2SO3
Strong Acids
STRONG BASESGroup 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 BASESWhen 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 neutralExample: 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
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Kb = [HB+][OH−][B]
What is the pH of 0.1 M NH3?
Kb = 1.8 x 10-5
1. 2.872. 4.743. 7.004. 9.255. 11.1
CONJUGATE ACID BASE PAIRSCONJUGATE ACID BASE PAIRSDiffer only by the presence or absence of a proton (H+).
Conjugate Acid = Conjugate Base + H+
Examples:H3O+ / H2O 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 1014
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 stronger conjugate base
H-F + OH- F- + H2O
Stronger acid6.9 x 10-4
Weaker acidKa = 10-14
Stronger base Weaker baseKb = 1.4 x 10-11
ACETIC ACID
Acid: CH3COOH H+ + CH3COO
Base: CH3COO + H2O CH3COOH + OH-
----------------------------------------------- H2O H+ + OH-
Ka = 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) • 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+
• 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. 3.212. 5.493. 7.004. 8.515. 10.8
LEWIS ACIDSAny 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 BASESAny 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 CATIONSTo 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+ 1.0Li+ 1.5Ca2+ 2.1Mg2+ 3.1Zn2+ 2.7Cu2+ 2.8Al3+ 6.7Cr3+ 4.8Fe3+ 4.7
HYDRATIONMetal ions attract the lone pairs on the oxygen in water molecules. This is a Lewis acid – Lewis base reaction.
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+ M(H2O)n-1(OH)(z-1)+ + H+
Fe(H2O)63+ Fe(H2O)5(OH)2+ + H+ (Ka=6.7 x 10-3)
Mz+ OH
H
:
:-
+
ACIDS AND BASES SO FARACIDS AND BASES SO FAR1) Arrhenius, Brønsted, and Lewis definitions2) pH, pOH
3) Acid and Base Dissociation Constants – Ka and Kb
4) [H+] [OH-] = 1 x 10-14 = Ka x Kb
1) pH and % ionization calcn for strong and weak acids/bases
2) Conjugate Acid-Base Pairs:ArrheniusBronsted-LowryLewis
7) Salts – Hydrolysis8) Structure Related to Acid-Base Properties (Oxyacids)
YOU SHOULD KNOW
GIVEN FIND
pH [H+], [OH-], pOH[H+] or [OH-] pHList of acids Weaker /StrongerList of pKa’s or Ka’s Weaker /StrongerKa or pKa and [HX] pH, [H+], [OH-]pH and [HX] Ka
Recall that a small Ka means high pKa, and both mean weak acid and not much dissociation.
Acid/Base SALTS Review1) Which one of the following salts would have a
basic aqueous solution?
1. KF 3. NaI 2. Al(NO3)3 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. Cl- 4. NH3 2. OH- 5. H+
3. CN-
