Acid-Base Equilibria

Acid-Base EquilibriaAcid-Base Equilibria

REVIEWREVIEW

• Electrolyte: Substances that dissolves in water

to produce solutions that conduct electricity

• Nonelectrolytes: Substances whose aqueous

solutions do not conduct electricity

• Strong and weak relates to the degree of

dissociation or ionization

• In aqueous solutions protons occasionally exist

as hydrated molecules, H(H2O)n+ (n=1)

Definitions:

Arrhenius: An acid is a substance that increases the H+ (or H3O+)

concentration in an aqueous solution.

HCl + H 2O H3O+ + Cl-

A base is a substance that increases the OH- concentrationin an aqueous solution.

HCl H+ + Cl-

NaOH(s) Na+ + OH-

What about Na2CO3 ????

Bronsted-Lowry:

Acid = a proton donor in a Reaction

Base = a proton acceptor in a Reaction

HCl(aq) + NaOH(aq) → HOH + NaCl

Lewis:

An acid is an electron pair acceptor

H+

acid

A base is an electron pair donor

:O:H-....

H:O:H..

..

water

Acid/Base reactions:

Produce water and a salt (and sometimes carbon dioxide).

Hint: concentrate on the water first. Remember, water has the formula HOH.

Complete and balance the following:

HCl + KOH

HCl + Ca(OH)2

HOH + KCl

2

Require equal numbers

2HOH + CaCl2

1. Ba(OH)2 + H3PO4

2. HC2H3O2 + NaOH

3. H2SO4 + KOH

4. H2CO3 + NaOH

5. Na2CO3 + HCl

6. NH4OH + H2SO4

7. NH3 + HCl

Give a definition of an acid:

An acid is a proton donor (H+)

Give a definition of a base:

A base is a proton acceptor

Conjugate acids and Conjugate bases

HCl + KOH HOH + KClacid base conj. acid conj. base

Na2CO3 + 2HCl H2CO3 + 2NaCl base acid conj. acid conj. base

Na2CO3 + 2HCl H2O + CO2(g) + 2NaCl conj. baseconj. acidacidbase

NH3 + HCl NH4+ + Cl-

What is a strong Acid?

An Acid that is 100% ionized in water.

Strong Acids:100% ionized (completely dissociated) in water.

HCl + H2O H3O+ + Cl-

often written as:

HCl H+ + Cl-

Strong Acids:Perchloric HClO4

Chloric, HClO3

Hydrobromic, HBrHydrochloric, HClHydroiodic, HINitric, HNO3

Sulfuric, H2SO4



What is a strong Base?

A base that is completely dissociated in water (highly soluble).

NaOH(s) Na+ + OH-

Strong Bases:

Group 1A metal hydroxides(LiOH, NaOH, KOH,RbOH, CsOH)

Heavy Group 2A metal hydroxides[Ca(OH)2, Sr(OH)2, andBa(OH)2]

Weak Acids: “The Rest”



Note the “one way arrow”.

Weak Acids:Only a small % (dissociated) in water.

HC2H3O2 + H2O H3O+ + C2H3O2

-

Note the “2-way” arrow.

Why are they different?

Strong Acids:

HCl HCl HClHCl HCl

ADD WATER to MOLECULAR ACID

(H2O)

Strong Acids:

(H2O)H3O+

H3O+

H3O+

H3O+

H3O+

Cl-

Cl-

Cl-

Cl-

Cl-

Note: No HCl molecules remain in solution, all have been ionized in water.

HC2H3O2

HC2H3O2

HC2H3O2

HC

2H3O

2

HC2H3O2

(H2O)

Weak Acid Ionization:

Add water to MOLECULES of WEAK Acid

HC2H3O2

HC2H3O2

HC2H3O2

HC

2H3O

2

HC2H3O2

H30

+ C2H3O2-

(H2O)

Weak Acid Ionization:

Note: At any given time only a small portion of the acid molecules are ionized and since reactions are running in BOTH directions the mixture composition stays the same.

This gives rise to an Equilbrium expression, Ka

H30+ C2H3O2

-HC2H3O2

ACID-BASE CONCEPTSACID-BASE CONCEPTS

• Acid-base reactions may be one of the most

important class of reactions

• The most basic of the acid-base concepts is

the Arrhenius theory

• Acids are substances that dissociate in water

to produce hydronium ions, H3O+ and bases are

substances that dissociate in water to produce

hydroxide ions, OH-

ACID-BASE STRENGTHACID-BASE STRENGTH

• A strong acid is almost completely dissociated

or ionized

• A weak acid is only partially dissociated or

ionized

HCl(aq) H+(aq) + Cl- (aq)

HF(g) + H2O(l) H3O+(aq) + F- (aq)

HCl(aq) H+(aq) + Cl-(aq)

• So each reaction is an equilibrium with a K

STRONG ACIDS AND BASESSTRONG ACIDS AND BASES

• Common strong acids are either monoprotic or

diprotic

HA(aq) + H2O(l) H3O+(aq) + A-(aq) 100%

• Means [HA] = [H3O+]

• Similar situation for strong bases

• Typical strong bases: Group 1A metal and Ca,

Sr, and Ba hydroxides

• Strong acids: HX, HNO3, HClO3, HClO4, H2SO4

WEAK ACIDSWEAK ACIDS

• Weak acids and bases are weak electrolytes

• For a weak acid, HA

• Ka: the dissociation or ionization constant

• For acetic acid

[H3O+][A- ]

Ka = [HA]

HA(aq) + H2O(l) H3O+(aq) + A- (aq)

CH3COOH(aq) + H2O(l) CH3COO- (aq) + H3O

+(aq)

Ka = [H3O+][CH3COO-]/[CH3COOH]

• Acids with larger ionization constants ionize or

dissociate to a greater extent than acids with

smaller ionization constants

• The larger the value of Ka, the higher [H3O+]

and the stronger is the acid

HIO3; Ka = 1.6 x 10-1

CH3COOH; Ka = 1.8 x 10-5

HCN; Ka = 6.2 x 10-10

AUTOIONIZATION OF WATERAUTOIONIZATION OF WATER

• Reaction called autoionization of water

K = [H3O+][OH-]/[H2O]2

K[H2O]2 = [H3O+][OH-]

Kw = [H3O+][OH-]

• Kw is called ion product of water

• At 25 °C, Kw = 1.0 x 10-14

• Valid also for dilute aqueous solutions

H2O(l) + H2O(l) H3O+(aq) + OH- (aq)

THE pH SCALETHE pH SCALE

• The hydronium ion concentration is a measure

of a solution’s acidity

• Usually small numbers

• The pH scale is used express acidity and

basicity

pH = -log[H3O+] so [H3O+] = 10-pH

• Note that as pH increases [H3O+] decreases

• Value of Kw varies with temperature

• “Acidity” depends on the concentration of

H3O+ ions

Acidic: [H3O+] > [OH-]

Basic: [H3O+] < [OH-]

Neutral: [H3O+] = [OH-]

• Notice that neutral does NOT necessarily

mean pH 7

• pH is usually quoted with the same number of

significant digits as the concentration

• pH can be measured using a pH meter or by an

acid-base indicator

• An indicator is usually an organic acid that have

different colors in solutions of different pH

• Indicators exist that cover the entire pH scale

e.g. bromothymol blue: 6.0 – 7.6

phenolphthalein: 8.0 – 10

HI n(aq) + H2O(l) H3O+(aq) + I n- (aq)

• The p-scale can also be applied to

ionization

constants

pKa = - log Ka

• The larger the value of Ka the smaller

the value of pKa and the stronger the acid

• From the acid-dissociation constant we

can

calculate equilibrium concentrations as

well as

pH

BRØNSTED-LOWRY THEORY

• An acid is a proton donor

• A base is a proton acceptor

• An acid-base reaction is the transfer of a

proton from an acid to a base

• Reactions can be described in terms of what

are called conjugate acid-base pairs

- species which differ by a proton

- a charge

- B is a proton acceptor; it is a base

- BH+ is a proton donor; it is an acid

So A- is the conjugate base of HA and

BH+ is the conjugate acid of B

• HF/F- and H3O+/H2O are conjugate pairs

HA + B BH+ + A-

Conjugate acid- basepairs

HF(g) + H2O(l) H3O+(aq) + F- (aq)

Relationship of Ka and concentration

• In dilute solutions of weak acids, the assumption

- is that all the H3O+ is coming from the acid

- the concentration change of a species is small

compared to the initial concentration of that

species

- at equilibrium, [HA] [HA]init

• Always be cautious when using the assumptions

given above

%-dissociation

• Amount of acid that dissociates can be

expressed as a percent

• In general %-dissociation increases with the

value of Ka

• For any weak acid, HA, %-dissociation

increases with dilution

%- Dissociation = ___________ x 100[HA]dissociated

[HA]initial

POLYPROTIC ACIDSPOLYPROTIC ACIDS

• Polyprotic acids: Acids that provide more than

one hydronium ion in solution

e.g. H2SO4, H3PO4

• Polyprotic acids ionize in a stepwise manner

• Consider sulfuric acid, H2SO4

H2SO4(aq) + H2O(l) H3O+(aq) + HSO4

- (aq)

HSO4- (aq) + H2O(l) H3O

+(aq) + SO42- (aq)

• In general, Ka1 > Ka2

WEAK BASESWEAK BASES

• Weak bases undergo equilibria in water

• For a general weak base, B

• Kb is called the base-dissociation constant

NH3(aq) + H2O(l) NH4+(aq) + OH- (aq)

B(aq) + H2O(l) BH+(aq) + OH- (aq)

[BH+][OH- ]Kb = [B]

• For conjugate acid-base pairs the product of

their equilibrium constants is the ionization

constant for water

Ka x Kb = Kw

HN

OH

H

Hydroxylamine

HN

CH3

HMethylamine

FACTORS AFFECTING ACID STRENGTH

• The strength of an acid depends on the polarity

of the H-E bond

• The polarity of the bond is related to the bond

strength of H-E

• The weaker the H-E bond the stronger the

acid

• Take the hydrohalic acids:

HX; X = F, Cl, Br, I

HF << HCl < HBr < HI

• For binary acid in the same group H-A bond

strength determines acid strength

• Same applies to other groups:

H2O < H2S < H2Se

• For binary acids in the same row polarity of the

H-E bond determines acid strength

:- acid strength increase with the

electronegativity of E

• The oxoacids are also important

e.g. HNO3, H2SO4, HClO4

• The acidity is dictated by factors that affect

the O-H bond strength

1. The electronegativity of the central element

2. The oxidation number of the central element

SO

OO

O

H

HO N

O

O

H

H O I H O Br H O Cl< <hypoiodous acid 2.5

hypobromous acid 2.8

hypochlorous acid 3.0

HClO4 > HClO3 > HClO2 > HClO

+7 +5 +3 +1

• Some small, highly charged metal ions are quite

acidic

:- they are hydrated and transfers a proton

to a water molecule

SALTS: ACID-BASE PROPERTIESSALTS: ACID-BASE PROPERTIES

• Salts: an ionic compound that is formed when

an acid neutralizes a base

NaOH(aq) + HCl(aq) NaCl(aq) + H2O

• Aqueous solutions of salts can be neutral, acidic

or basic

Strong acid + strong base neutral solutions

Strong acid + weak base acidic solutions

Strong base + weak acid basic solutions

• When a salt dissolves in water, its constituent

ions may react with water – reaction called

hydrolysis

NEUTRAL SOLUTIONS

• Salts of strong acids and strong bases

e.g. NaCl

• Because the ions do not hydrolyze

• Cl- is the conjugate base of HCl – it is a weak

base

• The same argument is made for Na+

• Essentially [H3O+]/[OH-] ratio does not

change

• {Group 1A metals, Ca2+, Sr2+, Ba2+} and {I-,

Br-, Cl-, NO3-, ClO4

-} does not hydrolyze

Cl-(aq) + H2O(l) HCl(aq) + OH- (aq)

Cannot occur

Acidic Solutions

NH3(aq) + HCl(aq) NH4Cl(aq)

• Salts from weak bases and strong acids

• In aqueous solution NH4+ undergo hydrolysis

- the chloride ion does not

• The generation of H3O+ from the reaction

makes these solutions acidic

NH4+(aq) + H2O(l) NH3(l) + H3O

+(aq)

Basic Solutions

• Salts from strong bases and weak acids give

basic solutions

• This basic anions of weak acids hydrolyze to

form hydroxide ions

NaCH3COO CH3COO- + H3O+

CH3COO- (aq) + H2O(l) CH3COOH(aq) + OH- (aq)

SALTS FROM WEAK ACID AND BASESSALTS FROM WEAK ACID AND BASES

e.g. NH4CH3COO, ammonium acetate

• Aqueous solution of the salts may be basic,

acidic or neutral

• The pH depends on the relative Ka and Kb of

the parent acid and base

• Consider the case when Ka = Kb

NH4CH3COO(aq) NH4+(aq) + CH3COO-(aq)

• Both ions can undergo hydrolysis

• Hydrolysis constant for the acetate ion (Kbh) is

equal to the hydrolysis constant for the

ammonium ion (Kah)

• The same concentration of H3O+ as OH- is

produced solution is NEUTRAL


+(aq)

CH3COO- (aq) + H2O(l) CH3COOH(aq) + OH- (aq)

• If the parent Ka > Kb: solutions acidic

e.g. NH4F

Ka(HF): 7.2 x 10-4 > Kb (NH3): 1.8 x 10-5

Kb(F-): 1.4 x 10-11 < Ka(NH4+): 5.6 x 10-10

• So NH4+ hydrolyzes to a greater extent than F-

more H3O+ is produced than OH-


+(aq)

F- (aq) + H2O(l) HF(aq) + OH- (aq)

• If the parent Ka < Kb: solutions basic

e.g. NH4CN

Ka(HCN): 4.9 x 10-10 < Kb (NH3): 1.8 x 10-5

Kb(CN-): 2.0 x 10-10 > Ka(NH4+): 5.6 x 10-10

• So NH4+ hydrolyzes to a greater extent than F-

more OH- is produced than H3O+


+(aq)

CN-(aq) + H2O(l) HCN(aq) + OH-(aq)

LEWIS THEORYLEWIS THEORY

• The Lewis acid-base theory is the most

complete that were are going to see

• In this theory:

- an acid is any species that can accept an

electron pair

- a base is any species that can donate an

electron pair

• The Lewis acid typically have vacant orbitals

• Common Lewis acids: metal ions and electron

deficient molecules such as AlCl3, BF3, etc.

• A typical Lewis base would be ammonia, NH3

as well as the common anions

• Acid-base reactions:

BF3 + NH3 → [F3B←NH3]

acid base acid-base adduct

AlCl3 + Cl- →[AlCl4]-

Acid-Base Equilibria

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