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 hardto 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. Theseare indicators.

HIn = H+ + In-

Indicators are less precise than pH meters.Many indicators do not have a sharp color change as afunction 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 betaken into account.)

When dissolved in water weak acids only partiallydissociate 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: CH3CO2HHFH3PO4

Acid Dissociation Constant (Ka) <<1

WEAK ACIDS

!

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. 1.8 x 10-5 M2. 4.2 x 10-3 M3. 1.8 x 10-6 M4. 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 withseveral attached oxygen atoms. These are calledoxyacids.

Acid strength increases with increasing oxidationnumber 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 atomincreases 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- in1 x 10-3 M H2CO3?

Which one of the following are not strongacids?

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

!

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 multiplytheir 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 = 14pKa + 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 withH2O 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 weakbase

Acidic solution (pH <7)

Example: NH4Cl (NH3 and HCl)

NH4+ + H2O NH3 + H3O+

4. Salt derived form a weak acid and a weakbase

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 watermolecules. This is a Lewis acid – Lewis base reaction.

Hydrated metal ions are acidic. Acidity increases withincreasing 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

:

:δ-

δ+

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

5) pH and % ionization calcn for strong and weak acids/bases

6) 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 weakacid 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 increasingbase strength:

N3- NO3

- HPO42- CN-

3) Which of the following cannot act as a Lewisbase?

1. Cl- 4. NH32. OH- 5. H+

3. CN-