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ACID AND BASE

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Understand different concepts of acid and base

Understand conce t of stron acid base and weakacid/base

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. rr en us e n t on: n ac s a su stance t at ssoc ates n water toproduce H+ ions (protons), and a base is a substance that dissociates in

water to produce OHions (hydroxide); an acid-base reaction involvese reac on o a pro on w e y rox e on o orm wa er

2. BrnstedLowry definition: An acid is any substance that can, ;

acid-base reactions involve two conjugate acid-base pairs and thetransfer of a proton from one substance (the acid) to another (the base)

3. Lewis definition: A Lewis acid is an electron-pair acceptor, and aLewis base is an electron-pair donor

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Arrh ni finiti n

An acid is a substance that, when dissolved in

water, increases the concentration of hydroniumion (H3O

+).

A base is a substance that when dissolved in

water, increases the concentration of hydroxideion OH .

ulooks at acids and bases in aqueous

.

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In the Arrhenius concept, a strong acid is a substance

that ionizes completely in aqueous solution to giveH3O

+ and an anion.

Exam le of stron acids include HCl HBr HI HNO

and H2SO4.

+ aq aq aq

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In the Arrhenius concept, a strong base is a

substance that ionizes completely in aqueoussolution to give OH and a cation.

Exam le of stron bases include LiOH KOH

Ca(OH)2, Sr(OH)2, and Ba(OH)2.

2H O + (aq) (aq) (aq)

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Most other acids and bases that we encounter are

weak. They are not completely ionized and exist in reversible

reaction with the corresponding ions.

Examples are acetic acid and ammonium hydroxide

3 (aq) 2 (l) 3 (aq) 3 (aq)

CH COOH H O H O CH COO+ + +

2H O + 4 (aq) 4 (aq) (aq)

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- An acid is the species donating the proton in a

proton-transfer reaction.

proton-transfer reaction.

In any reversible acid-base reaction, both forward

and reverse reactions involve roton transfer.

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- Consider the reactions:

H Cl H O H O

H

Cl

H H

Acid Base

H

H Cl H N H N ClH H

Acid Base

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- Consider the reaction of NH3 and H2O.

H

H H

H N

H

H

In the forward reaction H O donates a roton to NH .

ase c c ase

The H2O molecule is the acid and NH3 molecule is the base.

In the reverse reaction NH + donates a roton to OH-.

The NH4+ ion is the acid and OH- is the base.

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-A conjugate acid-base pair consists of two species in an

acid-base reaction, one acid and one base, that differby the loss or gain of a proton.

The conjugate base of an acid is a species that

the acid.

e con uga e ac o a ase s a spec es aresults from the addition of one proton to a base.

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H O H OH3C C O H H3C C O

H

Acid

H

Base Acid Base

-:

has been removed from the acid

Conjugate acid of a base: species that results from the addition ofone proton to a base

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HO

H O H OH3C C O H H3C C O

Acid Base Acid Base

Conjugate base of CH3COOH = CH3COO-

Conjugate acid of H2O = H3O+

See example 3.1

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An acid is an electron pair acceptor

A base is an electron air donor.F H F H

BF N H BF N H

acid base

o Boron trifluoride accepts the electron pair, so it is aLewis acid. Ammonia donates the electron air,

so it is the Lewis base.

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(aq) 2 (l) 3 (aq) (aq)HCl H O H O Cl+ + +

3 (aq) 2 (l) 3 (aq) 3 (aq)CH COOH H O H O CH COO+ + +

Weak acid

o e s ronger ac s are ose a ose e rhydrogen ions more easily than other acids.

o Similarly, the stronger bases are those thathold onto hydrogen ions more strongly than

.

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-

Self-ionization is a reaction in which two like

molecules react to give ions. In the case of water, the following equilibrium is

established.

2 (l) 2 (l) 3 (aq) (aq)H O H O H O OH+ + +

The equilibrium-constant expression for this system is:

w 3K [H O ][OH ]+ =

At 25 oC, the value of Kw is 1.0 x 10-14

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-

With Kw value, we can calculate [H+] and [OH] in

pure water These ions are produced in equal numbers in pure

water, so if we let x = [H+] = [OH]

14 o1.0 10 (x)(x) at 25 C =

7

H O OH 1.0 10 M+

= =

+ ,

the Kw expression still holds.

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+

In aqueous solution, [H3O+] and [OH] is related by the

w .

In the solution above, the [H3O+] is approximated to

0.00015 M We can calculate [OH] from w

3

K[OH ]

[H O ]

+

=

111.0 10

6.7 10 M0.00015

= =

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The pH of a solution is defined as the negative of

the common logarithm of [H3O+

].

3pH log[H O ]+=

We can also get pOH as

=

3p w og og .

Kw lo H O lo OH 14+

= =

= + =

pKw pH pOH 14= + =

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In a neutral solution, [H3O+] is 1.0 x 10-7, the pH =

7.00.

+ -7 , 3 . ,

less than 7.00.

Similarly, a basic solution has a pH greater than7.00.

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For a solution in which the hydrogen-ion

concentration is 1.0 x 10-3

, the pH is:

pH log(1.0 10 ) 3.00

= =

o e a e num er o ec ma p aces n e pequals the number of significant figures in thehydrogen-ion concentration.

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The pH of human arterial blood is 7.40. What is

-

( pH)+

( 7.40) 8[H ] 10 4.0 10 M+ = =

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An ammonia solution has a hydroxide-ion

concentration of 1.9 x 10-3

M. What is the pH ofthe solution?

We irst ca cu ate t e pOH:

3p og . .= =

en e p s:

H 14.00 2.72 11.28= =

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a

(aq) 2 (l) 3 (aq) (aq)HA H O H O A+ + +

(aq) (aq) (aq)HA H A+ +

the acid ionization constant (Ka) is

3H O A

K

+

=

H A

K

+ =

oraHA HA

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(aq) 2 (l) (aq) (aq)B H O BH OH

+

+ +

the base ionization constant (Kb) is

BH OH

K

+

= B

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a

We can use Ka and Kb values to compare the

strengths of acids and bases. The larger Ka, the stronger is the acid and the

greater the [H+] at equilibrium.

e arger b, e s ronger s e ase an egreater the [OH-] at equilibrium.

K lo K =

Low values of pKa and pKb

corres ond to lar e values

b bpK log K=

of Ka and Kb. Thus, thestronger acid, the lower isthe pKa

See table 3.2

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The relationship between conjugate acid-base

on zat on constants

-

a b w

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The degree of ionization is the fraction of the acid

molecules that ionize. Percent ionization gives the proportion of ionized

molecules on a ercenta e basis.

3[H O ] derived from HA+

initial [HA]

=

3[H O ] derived from HAPercent ionization 100+

=

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In a solution of a strong acid we can normally

ignore the self-ionization of water as a source ofH+.

e s usua y e erm ne y e s rong ac

concentration.

o However, the self-ionization still exists and is

res onsible for a small concentration of OH- ion.

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Example, calculate the concentration of OH- ion in

0.10 M HCl.HCl(aq) H (aq) Cl (aq)+ +

Because we started with 0.10 M HCl ( a strong acid ) thereaction will produce 0.10 M H+(aq).

Substituting [H+]=0.10 into the ion-product expression, we get:

. . =-14

-.[OH ] 1.0 10 M0.10

= =

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Similarly, in a solution of a strong base we can

normally ignore the self-ionization of water as asource of OH.

e s usua y e erm ne y e s rong ase

concentration.

o However, the self-ionization still exists and isresponsible for a small concentration of H+ ion.

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Example, calculate the concentration of H+ ion in

0.010 M NaOH.2H ONaOH(s) Na (aq) OH (aq)+ +

Because we started with 0.010 M NaOH ( a strong base ) the

reaction will produce 0.010 M OH-(aq).

Su stituting OH = . 1 into t e ion-pro uct expression, we get:

. . =-14 -[H ] 1.0 10 M

0.010

= =

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For weak acid and base, the calculation will involve

acid-base equilibrium. The acid-base equilibrium calculations are much like

those of the chemical e uilibrium cha ter.

The ICE table is always useful in calculation equilibriumroblems.

See examples 3.7 3.10

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Acid that can release more than one proton (H )

Example:

We can see that Ka1 >> Ka2

See example 3.11

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A i -B Pr rti f lt l ti n

Salts are strong electrolytes that completely

dissociate in water The pH of the solution depends on the hydrolysis of

ions.

The term salt hydrolysis describes the reaction of an, , .

Salt hydrolysis usually affects the pH of a solution.

Concentration of hydrogen ion or hydroxide ion hasbeen changed.

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When dissolve in water will produce neutral

solutions No hydrolysis

2

(s) (aq) (aq)NaCl Na Cl

+

+

(aq) 2 (l)a no reac on

+

(aq) 2 (l)

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When dissolve in water will produce basic solutions

Hydrolysisof conjugate base of the weak acid.H O +

(s) (aq) (aq)a a

(aq) 2 (l) (aq) (aq)CN H O HCN OH

+ +

Hydroxide ionincreases

See example 3.12

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When dissolve in water will produce acedic solutions

Hydrolysisof conjugate acid of the weak base.2

4 (s) 4 (aq) (aq)

NH Cl NH Cl +

4 (aq) 2 (l) 3(aq) 3 (aq)NH H O NH H O+ ++ +

Hydroniumionincreases

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A buffer solution is a solution of a weak acid or a

weak base and its salt. The solution has the ability to resist changes in

H u on the addition of small amounts of either

acid or base.

conjugate base/acid.

e so um ace a e-ace c ac u er sys em can ewritten as CH3COONa/CH3COOH or CH3COO/CH3COOH.

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Buffer solution consists of weak acid (HA) and its

conjugate base (A

)When acid is added, it is consumed by reacting with A-

HA A- A- + H3O+ HA + OH-

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Buffer solution consists of weak acid (HA) and its

conjugate base (A

)When base is added, it is consumed by reaction with HA

HA A- HA + OH-A- + H2O

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Adding small amount of acid or base does not

change [H3O+

] of the solution. Since pH = -log [H3O

+] , then the pH of the solution

is almost unchan ed.

However, there is a limit in which the buffer can.

This is called buffer capacity.

u er capacity epen s on t e amount o aci anconjugate base in the solution.

C l l f b ff l

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Calculation H of buffer solution

Consider bu er solution o acetate ion acetic acid

3 2 3 3CH COOH H O H O CH COO+ +

[ ]3 3

3Ka CH COOH=

3CH COO

pH pKa log

= +

Henderson-

3

con u ate base

[ ]

p p a og

weak acid

= +

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In general, the pH range over which a buffer

solution is most effective is about one pH unit oneither side of pH = pKa.

This corres onds to the ratios

[ ]conjugate base

[ ] . .

weak acid

[ ]1.0 log 1.0

weak acid