15 Acids and Bases Contents

15-1 The Bronsted-Lowry Definitions15-2 The Ion Product of Water, Kw15-3 The pH and Other “p” Scales15-4 Concentrations of Hydrogen Ions in Aqueous Solutions of Acids 15-7 The Common Ion Effect15-8 Buffer Solutions15-9 How Indicators Work

Acids and Bases Arrhenius Definitions － Acids produce hydrogen ions in

aqueous solution, and bases produce hydroxide ions in aqueous

solution.

Brønsted-Lowry Definitions － An acid is a proton (H+) donor,

and a base is a proton acceptor.

Lewis Acid-Base Definitions － An acid as an electron pair

acceptor and a base as an electron donor.

A rrh en iu s B ron s ted /L ow ery L ew is

A c id -B ase C on cep ts

strong acid formula

hydrochloric HCl hydrobromic HBr hydroiodic HI Nitric HNO3

chloric HClO3

perchloric HClO4

sulfuric H2SO4

Sulfuric acid is the only strong acid that is diprotic, meaning it has two protons to donate.

Acid H+ donor

Base H+ acceptor

15-1 The Bronsted-Lowry Definitions

Terms to know: proton - H+ ion

hydronium ion - H3O+; results from water reacting with

H+

H2O + H+ --> H3O+

conjugate base - whatever is left from an acid after a

proton has been donated

conjugate acid - whatever has been formed when a

proton has been accepted by a base

Acids are hydrogen ion (H+) donorsBases are hydrogen ion (H+) acceptors

HCl + H2O H3O+ + Cl-

donor, acid acceptor, base

+ +

HCl/Cl- ; H3O+ / H2O :Conjugate Acid-Base Pair

An acid-base reaction consists of the transfer of a proton for an acid to a base.

In acidic solutions, the protons (H+) that are

released into solution will not remain alone due to

their large positive charge density and small size.

They are attracted to the negatively charged

electrons on the oxygen atoms in water, and form

hydronium ions.

H+(aq) + H2O(l) = H3O+

(l)[H+] = [H3O+]

OH H

.. ..

H atom and H+ ion atomic structure? Use symbols H+ and H3O+ to represent the same thing

Conjugate Acid-Base Pair

• Conjugate acid-base pair consists of two substances

related to each other by the donating and accepting of

a single proton.

• Acids and bases that are related by loss or gain of H+.

• NH3+HCl NH4++ Cl-

• NH4+/NH3, HCl/ Cl-

A species like water that can react either as an acid or as a base is said to be amphoteric.

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

Acid 1 Base 2 Acid 2 Base 1

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

Base 1 Acid 2 Acid 1 Base 2

Summary

1. Acids and Bases can be molecules or ions

2. Not just defined in aqueous solutions

3. Conjugate pairs differ by H+

4. Some substances are both, AMPHOTERIC

depending on the reaction.

• Autoionization: A reaction in which a substance reacts with itself

to form ions.

• Water is the common amphoteric substance.

• HF(aq) + H2O(l) H3O+(aq) + F-(aq) H2O as base.

• NH3(aq) + H2O(l) NH4++ OH- (aq) H2O as acid.

• Kw=[H3O+][OH-]=1×10-14 ion-product constant (25 )℃

• pH=-log[H+]

2H2O(l) H3O+(aq)+OH-(aq)

15-2 The Ion Product of Water, Kw

• Water

– The most common amphoteric substance

• water can act as both an acid and a base

• water can autoionize:

– H2O + H2O <==> H3O+ + OH-

– one water molecule acts as an acid (H+ donor), the other acts

as an acid (an H+ acceptor)

– Kw = [H3O+][OH-] = 1.00 x 10-14 @ 25oC

• dissociation or ion product constant for water

• Kw

– In any aqueous solution at 25oC, the product of [H+] and

[OH-] will be 1.0 x 10-14

– So if you know the [H+], you can figure out the [OH-] and

vice versa

– If [H+] = [OH-], the solution is neutral

– If [H+] > [OH-], the solution is acidic

– If [H+] < [OH-], the solution is basic

• pH scale

– Because the [H+] in any solution is generally quite small, it

is easier to use the pH scale to represent a solution’s

acidity.

– pH comes from the Danish…potenz or strength of the H+

ion

– The pH of a solution is usually defined as the negative of

the base 10 logarithm.

– pH = - log[H+]

– pOH = - log [OH-]

15-3 The pH and Other “p” Scales

Acids and Bases• pH is a log scale

– when the pH changes by one, the [H+] concentration changes by a power of 10.

• A solution with a pH of 3 has 10 times more H+ than a solution with a pH of 4, and 100 times more H+ than a solution with a pH of 5.

– As pH decreases, the [H+] increases.

– Rule for significant figures for logarithms - the number of places after the decimal point is equal to the number of significant figures in the original number

• pH = - log 1.0 x 10-9 M (2 significant figures in 1.0 x 10-9)

• pH = 9.00 ( 2 places after the decimal point for significant figures)

Methods for Measuring the pH of an Aqueous Solution

(a) pH paper (b) Electrodes of a pH meter

Strong and Weak Acids

Strong:Strong: 100% dissociation100% dissociation

good Hgood H++ donor donor

equilibrium lies far to right (HNOequilibrium lies far to right (HNO33))

generates weak base (NOgenerates weak base (NO33--))

Weak:Weak: <100% dissociation<100% dissociation

not-as-good Hnot-as-good H++ donor donor

equilibrium lies far to left (CHequilibrium lies far to left (CH33COOH)COOH)

generates strong base (CHgenerates strong base (CH33COOCOO--))

15-4 Concentrations of Hydrogen Ions in Aqueous Solutions of Acids

Strong acids:

For dense and moderately dilute solutions (1.0×10-6

to 0.01M) of strong acids that have only one ionizable

hydrogen, [H+] = stoichiometric concentration of the

strong acid.

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

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

acid base conjugate acid

conjugatebase

3[ ][ ] [ ][ ]

[ ] [ ]

: acid dissociation constant

a

a

H O A H AK

HA HA

K

HA(aq) H+ （ aq ） +A- （ aq）

Weak acids:

P584 on the textbook: Table 15.3

Ionization Constants of some Common Weak Acids at 25 ℃

Ka I

ncre

ases

Strength vs. Ka

Calculating the pH of Weak Acid Solutions

HF H++F-

1.0MBefore 0 0Equilibrium (1-x)M xM xM

24 2 4

2 4 4 2

2

7.2 10 7.2 10 ( 1)1

7.2 10 7.2 10 2.7 10

log(2.7 10 ) 1.57

a

xK x if x

x

x x

pH

C/K<500

The 5% Rule

• The validity of an approximation should always be checked.

• The 5% rule should be used to evaluate which approximation are reasonable.

%.%.

.For HF,

%%[HA]

x

-

72100001

1072

5100

2

0

The approximation is acceptable.

The pH of a Mixture of Weak Acids

HNO2 is assumed to be the dominant producer

of H+.

HCN(aq) H+(aq)+CN-(aq) Ka=6.2X10-10

HNO2(aq) H+(aq)+NO2-(aq) Ka=4.0X10-4

H2O(l) H+(aq)+OH-(aq) Kw=1.0X10-14

4 2

2

[ ][ ]4 10

[ ]a

H NOK

HNO

HNO2(aq) H+(aq)+NO2-(aq)

5 0 0

5-x x x2 2

4 2

2

2 0

4 10 4.5 105 5

Using the 5% rule

4.5 10100% 100% 0.9%

[ ] 5

a

x xK x

x

x

HNO

2

-

- 2 -10

- -8

[ ] 4.5 10 1.35

To calculate [CN ]

[ ][CN ] (4.5 10 )[CN ]6.2 10

[HCN] 1

[CN ]=1.4 10 ( )

a

H pH

HK

M

Percent Dissociation

amount dissociationPercent Dissociation= 100%

initial concentration

HC3H5O3 H++C3H5O3-

0.1 0 00.1-x x x

3

3 34

3.7% 100% 3.7 10 ( )0.1

(3.7 10 )(3.7 10 )1.4 10

0.1a

xx M

K

15.7 The Common Ion Effect

• Common ion effect: The shift in equilibrium position

that occurs because of the addition of an ion already

involved in the equilibrium reaction.

The common ion effect is the shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance.

The presence of a common ion suppresses the ionization of a weak acid or a weak base.

Consider mixture of CH3COONa (strong electrolyte) and CH3COOH (weak acid).

CH3COONa (s) Na+ (aq) + CH3COO- (aq)

CH3COOH (aq) H+ (aq) + CH3COO- (aq)

common ion

16.2

A buffer solution is a solution of:

1. A weak acid or a weak base and

2. The salt of the weak acid or weak base

Both must be present!

A buffer solution has the ability to resist changes in pH upon the addition of small amounts of either acid or base.

16.3

CH3COOH (aq) H+ (aq) + CH3COO- (aq)

Consider an equal molar mixture of CH3COOH and CH3COONa

Adding more acid creates a shift left IF enough acetate ions are present

15.8 Buffer Solutions

• A Buffer Solution (or buffered solution) is one that

resists a change in pH when either hydroxide ions or

protons are added.

• A buffered solution may contain a weak acid and its salt

or a weak base and its salt. (HF+NaF, NH3+NH4Cl)

How Do the H+/OH- Ions Work in Buffered Solutions

• The equilibrium concentration of H+ and the pH are

determined by the ratio [HA]/[A-].

NaA(s) Na+(aq)+A-(aq)H2O(l)

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

][

][][

][

]][[

A

HAKH

HA

AHK aa

The Effect of Added Bases

When OH- are added, HA is converted to A- , causing

the ratio [HA]/[A-] to decrease. If the amount of HA

and A- originally present are very large compared

with the amount OH- added, the change in [HA]/[A-]

ratio is small.

NaA(s) Na+(aq)+A-(aq)H2O(l)

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

The Effect of Added Acids

When protons are added to a buffered solution, the added

H+ ions react with A- to form the weak acid. If [HA]

and [A-] are large compared with the [H+] added, only

a slight change in the pH occurs.

NaA(s) Na+(aq)+A-(aq)H2O(l)

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

Unique Properties of Buffer Solution

• The Effect of Dilution

－ The pH of a buffer solution remains essentially

independent of dilution.

• The Effect of Added Acids and Bases

－ A buffer solution resists pH change after addition of

small amounts of strong acids or bases

• Buffer Capacity

Which of the following are buffer systems? (a) KF/HF (b) KCl/HCl, (c) Na2CO3/NaHCO3

(a) KF is a weak acid and F- is its conjugate basebuffer solution

(b) HCl is a strong acidnot a buffer solution

(c) CO32- is a weak base and HCO3

- is it conjugate acidbuffer solution

16.3

What is the pH of a solution containing 0.30 M HCOOH and 0.52 M HCOOK?

HCOOH (aq) H+ (aq) + HCOO- (aq)

Initial (M)

Change (M)

Equilibrium (M)

0.30 0.00

-x +x

0.30 - x

0.52

+x

x 0.52 + x

16.2

Mixture of weak acid and conjugate base!

KKaa for HCOOH = 1.8 x 10 for HCOOH = 1.8 x 10 -4-4

[H+] [HCOO-]Ka = [HCOOH]

x = 1.038 X 10 -4

pH = 3.98

30.0

52.0108.1

)30.0(

)52.0(

][

]][[ 4 x

x

xx

HCOOH

HCOOHKa

NaF(s) Na+(aq)+F-(aq)H2O(l)

HF(aq) H+(aq)+F-(aq)

1 0 01-x x x

1 0 01 10

4

[ ][ ] (1 ) ( )(1)

[ ] 1 1

7.2 10

a

H F x x xK

HF x

x

15.9 How Acid-Base Indicators Work

HIn(aq) H+(aq)+In-(aq)

][

][

][][

]][[

HIn

In

H

K

HIn

InHK a

a

Add a few drops of the phenolphthalein indicator to a acidic solution. (pH=1)

10000000

1

][

][

101

101

][ 1

8

HIn

In

H

Ka

The ratio shows that the predominant form of the indicator is HIn, resulting in a red solution.

As OH- is added to this solution, [H+] decreases and the

equilibrium shift to right, changing HIn to In-. A color

change from red to reddish purple will occur.

For most indicators, about 1/10 of the initial form must be

converted to the other form before a new color is

apparent.

Indicators