Chemistry-Ch11_Acids and Bases

8/9/2019 Chemistry-Ch11_Acids and Bases

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1111Acids andAcids and

basesbases


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OH H

+ OH H H O

OH H

H

++

11.1 The Brnsted11.1 The BrnstedLowryLowrydefinition of acids and basesdefinition of acids and bases

An acid is a proton donor

A base is a proton acceptor

An acid-base reaction involves thetransfer of a single proton from one

species to another

A BrnstedLowry acid will donate aproton to a BrnstedLowry base

acid base


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For a proton to be measurably acidic, it must

be bound to another atom via an appreciably

acidic bond

Acids tend to contain protons bound to group16 or 17 elements

Basic species require the presence of one or

more lone pairs

Not all species containing lone pairs act as

bases

Bases usually contain group 15 or 16 elements,

the atoms of which are often deprotonated


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Conjugate acid-base pairs For BrnstedLowry acid-base reactions

both the forward and reverse reactions are

acid-base reactions There are always 2 sets of species oneither side of the equation that differ byonly one proton

These are conjugate acid-base pairs

Cl 2 3 Cl


acid acid basebase

conjugate pair

conjugate pair


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11.2 Acid11.2 Acid--base reactionsbase reactions

in waterin water The autoprotolysis of water

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

3

is called the hydronium ion is called the hydroxide ion

This is an autoprotolysis reaction

The extent of the autoprotolysis of water

can be determined using the equilibrium

constant for this process

Kw = [ 3 ][] = 1.0 1014 (25.0 C)

Kw is the autoprotolysis constant of water


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Kw very small, autoprotolysis of water

proceeds only to a small extent

Concentrations of 3and in pure

water at 25 C

Kw = [ 3 ][] = 1.0 1014

If [ 3 ] = [

], then14 7 1

3[ ] [ ] 1.0 10 1.0 10 mol L

! ! v ! v


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Pure water is neutral because it contains

equal concentrations of3

and

[ 3 ] > [] in acidic aqueous solutions

[ 3 ] < [] in basic aqueous solutions

There is an inverse relationship between

[ 3 ] and [ ] in aqueous solutions

As the concentration of one increases, the

other must decrease in order to keep Kwconstant


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in waterin water The concept of p

To avoid working with inconveniently small

numbers we commonly express [3

] in

terms of the p of the solution

p = log[ 3 ]

p = log[ ]

log means log10, as opposed to naturallogarithms (ln or loge)

[ 3 ] = 10p

[

] = 10

p


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in waterin water The concept of p Pure water at 25 C

p = log[ 3 ] = log(1.0 107) = 7.00

p = log[

] = log(1.0 107

) = 7.00 Acidic solutions

[ 3 ] > []

At 25 C p < 7

Basic solutions [ 3 ] < [

]

At 25 C p > 7

At 25 C p p = pKw

p p = 14


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in waterin water The strength of acids and bases A strong acid reacts completely with water

to give quantitative formation of 3 A weak acid reacts incompletely with

water to form less than stoichiometricamounts of 3

A strong base reacts completely withwater to give quantitative formation of

A weak base reacts incompletely withwater to form less than stoichiometric

amounts of


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in waterin water The strength

of acids andbases


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11.3 Strong acids and bases11.3 Strong acids and bases

The conjugate base of a strong acid is very weak

3 is the strongest acid and is the strongest

base that

can exist inaqueous

solution


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11.3 Strong acids and11.3 Strong acids and

basesbases

p calculations in solutions of strong

acids and bases

Relatively straightforward exercise instoichiometry as strong acids and bases

react with water almost completely

A 0.010 M solution of Cl contains 0.010

mol L1

of 3 .The p of this solution is:

p = log[ 3 ] = log(1.0 102) = 2.00

Cl(aq) 2 (l) H3 (aq) Cl(aq)


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basesbases

contd A 0.010 M solution of Ba( H)2 contains

0.020 M of H.

The p H of this solution is:

p H = log[ H] = log(2.0 102) = 1.70

pH = 14.00 1.70 = 12.30

Ba( H)2(aq) H2 (l) Ba2 (aq) 2 H(aq)


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basesbases

Suppression of the autoprotolysis of

water

Previously we assumed that all of the

H3 or H in solutions of strong acid

or base comes from the acid or base

respectively

There is another potential source of bothH3 and H in aqueous solution,

namely water

[H3 ]total = [H3 ]from acid [H3 ]from H2


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basesbases


water

Except in very dilute solutions of strong

acids, [H3 ]from water is negligible compared

to [H3 ]from acid

In any solution of an acid, the autoprotolysis

of water is suppressed

Add another source ofH3 to water,

instantaneously increase Qw, so that

Qw > Kw, resulting in lower amounts of both

H3and H from the autoprotolysis of

water


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basesbases


water

Calculate the pH of a 1.0

1010 M solution of

HCl(aq)

[H3 ]from HCl = 1.0 1010 M

[H3 ]from water= 1.0 107 M

[H3 ]total = 1.0 107 1.0 1010 M

= 1.0 107 M

pH = log[H3 ]

= log(1.0

107) = 7.00


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11.4 Weak acids and11.4 Weak acids and

basesbases

Weak acids and bases react

incompletely with water

Quantify the extent to which thesereactions occur by looking at the values

of their equilibrium constants

HA(aq) H2

(l) H3

(aq) A(aq)

Ka is called the acidity constant

? A3

a

H O

HK

- - !


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basesbasesB(aq) H2 (l) BH

+(aq) + H(aq)

Kb is called the basicity constant

Ka and Kb values are generally

significantly less than 1 pKa = logKa Ka = 10

pKa

pKb = logKb

Kb = 10pKb

? A

+H OHK

- - !


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basesbases

The strength of a weak acid is

determined by its Ka value

The larger the Ka

the stronger the acid

The smaller the pKa the stronger the acid

The strength of a weak base is

determined by its Kb value The larger the Kb the stronger the base

The smaller the pKb the stronger the

base


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basesbases

KaKb = Kw

pKa

+ pKb

= pKw

These relationships hold for any acid-

base conjugate pair, regardless of the

strength of the acid or base

? A

+ -

3

a

H O HCOO=

HCOOHK

- -

? A -

-

H OOH OH

H OOK

-

-

HCOOH(aq) +H2O(l) H3O+(aq) +HCOO(aq)

HCOO(aq) +H2O(l) HCOOH(aq) +OH(aq)

H2O(l) +H2O(l) H3O+

(aq) +OH

(aq)+ -

3= H O OHwK - -


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11.4 Weak acids and bases11.4 Weak acids and bases

There is an

inverse

relationship

between thestrengths of

the acid and

base

members of aconjugate pair


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basesbases

pH calculations in solutions of weakacids and bases

Must use the value ofKa orKb to

determine [H3O+] or [OH] and hence pH

The pH of a 1.0 M solution of acetic acid, CH3COOHis:

CH3

COOH(aq) +H2

O(l) H3

O+(aq) + CH3

COO(aq)

Cannot solve this equation for [H3O+] as there are

two unknowns: [CH3COO

] and [CH3COOH]

? A

+ -

3 3 5

a

3

H O H OO.8

H OOHK

- - ! v


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basesbases pH calculations in solutions of weak

acids and basesH2O + CH3COOH H3O

+ + CH3COO

Initial concentration 1.0 0 0

Change in concentration x +x +xEquilibrium concentration 1.0x x x

? A

+ -

3 3 5

a

3

H O H OO.8

H OOH .

x xK

x

- - ! ! v

2

51.8 101.0x ! v

2 5 51.0 1. 10 1. 10x ! v ! v5 31. 10 4.2 10 Mx ! v ! v

3pH log 4.2 10 2.3! v !


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basesbases

pH calculations in solutions of salts of

weak acids and bases

What is the pH of a 0.010 M solution of NaOCl?ForHOCl, Ka = 3.0 10

OCl(aq) +H2O(l) HOCl(aq) +OH(aq)

? A -

-

HO l OH

O lK

-

-

147

8

1.0 103.3 10

3.0 10w

b

a

KK

K

v! ! ! v

v


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basesbases pH calculations in solutions of salts of

weak acids and bases

H2O + OCl HOCl + OH

Initial concentration 0.10 0 0Change in concentration x +x +x

Equilibrium concentration 0.10x x x

[HOCl] = [OH] = x

? A

- 2

7b -

HOCl OH= 3.3 10

0.10 0.10OCl

x x xK

x

- ! ! ! v -

2 7 80.10 3.3 10 3.3 10x ! v ! v8 43.3 10 1.8 10 Mx ! v ! v


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pH calculations in solutions of salts ofweak acids and bases


basesbases

-

4

pOH l [OH ]

pOH l 1.8 10

pOH 3.74

!

! v

!

pH = 14.00 - pOH

pH = 14.00 - 3.74

pH = 10.26

The pH f this s luti n is 10.26


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Solutions that contain the salt of a

weak acid and a weak base

When both the cation and anion in a

single salt are able to affect the pH the

net effect is dependent on the relative

strengths of its ions, one functioning as

an acid and the other as a base

NH4HCOO

Ka of NH4+ is 5.6 x 1010

Kb ofHCOO is 5.6 x 1011

Solution of NH4HCOO is slightly acidic


basesbases


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basesbases

Situations where simplifyingassumptions do not work

When a weak acid HA reacts with water

[HA]equilibrium = [HA]initial x

Previously assumed

[HA]equilibrium = [HA]initial

This is only the case if [HA]initial is greaterthan, or equal to, 400 times the value ofKa

If this assumption is not valid, we must

solve a quadratic equation


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basesbases

Situations where simplifying

assumptions do not work

Previously assumed that we could safelyneglect any contribution of [H3O

+] or

[OH] from the autoprotolysis of water

This is only the case if the value of[H3O+] or [OH] from the acid or base is

greater than 1 105


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11.5 The molecular basis11.5 The molecular basis

of acid strengthof acid strength

Binary acids

An acid containing H and only one other

element, generally nonmetal

HF


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11.5 The molecular basis of11.5 The molecular basis of

acid strengthacid strength

Binary acids

Relative acidity increases as we proceeddown any group in the periodic table

Due primarily to the decreasing bonddissociation enthalpy

Across a period the acidity is determinedprimarily by the electronegativity of the

nonmetal X As HX bonds become more polar +

charge on the H atom increases, making iteasier for the proton to separate and bond

to a base


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Oxoacids

Acids composed of hydrogen, oxygenand some other element

A feature common to the structures of alloxoacids is the presence ofOH groupsbonded to some central atom

Reaction of oxoacid with water involves

breaking of an OH bond

O O

O

O

H H SeO O

O

O

H H

sulfuric acid selenic acid


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Oxoacids

The more polar the bond, the stronger

the acid

Two principal factors determine how the

polarity of the OH bond is affected

The electronegativity of the central atom in

the oxoacid

The number of oxygens attached to the

central atom

G-O-H- +


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Oxoacids

As the electronegativity of the centralatom increases, the oxoacid becomes a

better proton donorHIO4 < HBrO4 < HClO4

H3PO4 < H2SO4 < HClO4 As the number of lone oxygens increases,

the oxoacid becomes a better protondonor

O HCl ClO O HO HClO

OCl O HO

O

O

< < 7 as conjugate base formed by addition

ofOH to weak acid

Calculate pH using method for the

determination of pH of a salt of a weak acid


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11.7 Acid11.7 Acid--base titrationsbase titrations

Weak acid strong base titration

The alkaline region

Simply adding excess OH

pH controlled by excess OH


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Acid-base indicators

Indicators used in acid-base titration are

weak acids

HIn

Acid form ofHIn must have a different

colour from the conjugate base form In-

HIn(aq) +H2O(l) H3O

+

(aq) +In

-

(aq)acid form base form

colour 1 colour 2


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Typically as we pass

the equivalence point

there is a suddenand large change

in pH.

This causes a

sudden shift in theposition of the

equilibrium for

the indicator


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The pKIn of an

indicator should be

equal or close to thepH at the

equivalence point

When performing a

titration we want touse as little indicator

as possible



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Lewis acids and bases

A Lewis acid is an electron-pair acceptor

A Lewis base is an electron-pair donor

Strength of Lewis acids and bases is not

as readily quantified as BrnstedLowry

counterparts

B

F

F FN

H

H HB N

F H

F HHF

Lewis acid Lewis base adduct


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Chemistry-Ch11_Acids and Bases

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