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Chapter 14Acids and Bases

Properties of Acids and Bases

Svante Arrhenius(1859-1927)

First to develop a theory for acids and bases in aqueous

solution

Arrhenius AcidsCompounds which dissolve (dissociate) in water to produce H+ ions

Examples:

HCl(aq) → H+(aq) + Cl-(aq)

HNO3(aq) → H+(aq) + NO3-(aq)

Arrhenius BasesIonic compounds which dissolve (dissociate) in water to produce OH-

ions

Examples:

NaOH(aq) → Na+(aq) + OH-(aq)

Ca(OH)2(aq) → Ca2+(aq) + 2OH-(aq)

Arrhenius acids react with Arrhenius bases to form neutral salt solutions

Ammonia can also react with an Arrhenius acid to form a neutral salt solution!

It’s behaving like a base but contains no OH-!

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A more sophisticated definition of acids and bases is needed…

The Brønsted-Lowry Model

Brønsted-Lowry AcidsCompounds which donate protons (H+) in aqueous solution

Protons cannot exist on their own in aqueous solution. Instead, they combine with water molecules forming the hydronium ion, H3O+:

HA(aq) + H2O(l) → H3O+(aq) + A-(aq)

Brønsted-Lowry BasesCompounds which accept protons (H+) in aqueous solution. These

protons are typically taken from water molecules forming the hydroxide ion, OH-:

B(aq) + H2O(l) → BH+(aq) + OH-(aq)

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Acid-Base EquilibriaWhen an acid (HA) donates H+ to a base (B), the products are A- and BH+:

HA(aq) + B(aq) → A-(aq) + BH+(aq)

Since the products are also an acid and a base, the reverse reaction can occur in which the acid BH+ donates H+ to the base A-:

A-(aq) + BH+(aq) → HA(aq) + B(aq)

We therefore have an equilibrium reaction:

HA(aq) + B(aq) A-(aq) + BH+(aq)

Conjugate Acid-Base PairsTwo substances related by the loss or gain of a single proton, H+

Write the formula and name of the conjugate base for each of the following acids:

a. HCN(aq)

b. NH4+(aq)

c. H2CO3(aq)

d. HSO4-(aq)

e. HClO2(aq)

Write the formula and name of the conjugate acid for each of the following bases:

a. CO32-(aq)

b. SH-(aq)

c. H2PO3-(aq)

d. Br-(aq)

e. NO2-(aq)

Acid Dissociation Constant, Ka

For:

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

Ka = [H3O+][A-]/[HA] = [H+][A-]/[HA]

H+ is normally used to represent the hydrated proton in acid-base reactions rather than H3O+

The size of Ka depends on the strength of the conjugate base, A-, compared to H2O

If A- is a much weaker base than H2O, the equilibrium position will be far to the right (Ka large) with most of the dissolved acid in ionized

form

If A- is a much stronger base than H2O the equilibrium position will be far to the left (Ka small) with most of the dissolved acid in

molecular form

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Acid Strength

Determined by the size of the acid dissociation constant, Ka, for the dissociation (ionization) equilibrium:

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

Strong AcidsThe equilibrium position of its dissociation reaction lies far to the

right:

HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)

Ka large

[H3O+] ~ [HA]0

Conjugate base much weaker base than H2O

Strong acids are excellent conductors of electricity since they almost completely dissociate into ions

Weak AcidsThe equilibrium position of its dissociation reaction lies far to the

left:

CH3CO2H(aq) + H2O(l) H3O+(aq) + CH3CO2-(aq)

Ka small

[H3O+] « [HA]0

Conjugate base much stronger base than H2O

Weak acids are poor conductors of electricity since they hardly dissociate and remain mostly in molecular form

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The Six Common Strong AcidsHCl(aq): hydrochloric acidHBr(aq): hydrobromic acid

HI(aq): hydroiodic acidHNO3(aq): nitric acid

H2SO4(aq): sulfuric acidHClO4(aq): perchloric acid

You can assume that all the others are weak!

Summary

Types of Acids Strong acids have weak conjugate bases!The conjugate base of a strong acid has a low attraction for protons

It is a much weaker base than water so the water molecule wins the competition for H+ ions:

HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)

HCl is a strong acidCl- is a weak conjugate base

Weak acids have strong conjugate bases!The conjugate base of a weak acid has a high attraction for protons

It is a much stronger base than water so it wins the competition for H+ ions over water:

CH3CO2H(aq) + H2O(l) H3O+(aq) + CH3CO2-(aq)

CH3CO2H is a weak acidCH3CO2

- is a strong conjugate base

Strengths of Acids and Conjugate Bases

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Amphoteric SubstancesA substance is referred to as amphoteric if it can as act as both an

acid and a base

Water is most common example!

Water accepts H+ in the presence of a stronger acid:

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

Water donates H+ in the presence of a stronger base:

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

Autoionization of WaterSince water is amphoteric it is possible for it to react with itself

forming a conjugate acid-base pair:

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

The equilibrium position is far over to the left so only a very tiny amount of H3O+ and OH- exist in pure water at room temperature

Ion-product Constant for Water, Kw

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

The autoionization of water leads to the equilibrium expression:

Kw = [H3O+][OH-] = [H+][OH-]

where Kw is called the ion-product constant for water

In pure water at 25 ºC, experiments show that:

[H+] = [OH-] = 1.0 x 10-7 M

so

Kw = [H+][OH-] = (1.0 x 10-7)(1.0 x 10-7) = 1.0 x 10-14

Kw is a measure of the strength of water as an acid

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Applications of Kw

Since Kw is an equilibrium constant, in any aqueous solution at 25 C, regardless of what it contains, the product of [H+] and [OH-] must

always be equal to 1.0 x 10-14

Therefore:

If [H+] ↑ then [OH-] ↓

and

If [H+] ↓ then [OH-] ↑

so that [H+][OH-] remains equal to 1.0 x 10-14

Neutral, Acidic and Basic Solutions1. [H+] = [OH-] solution is neutral

2. [H+] > [OH-] solution is acidic

3. [OH-] > [H+] solution is basic

Compositions of Acidic, Basic and Neutral Solutions For all solutions at 25 C, regardless of whether they are acid, basic or neutral:

Kw = [H+][OH-] = 1.0 x 10-14 = constant

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Calculating Ion Concentrations in WaterIf [H+] or [OH-] is given, Kw can be used to calculate the

corresponding [OH-] or [H+] concentration at 25 C:

Kw = [H+][OH-] = 1.0 x 10-14

Since Kw is an equilibrium constant, it varies with temperature and so a different value of Kw has to be used if the temperature is not 25 C

The pH ScalepH is a number representing the [H+] concentration in a solution:

Acidic SolutionpH < 7 [H+] > 1.0 x 10-7 MNeutral SolutionpH = 7 [H+] = 1.0 x 10-7 MBasic SolutionpH > 7 [H+] < 1.0 x 10-7 M

Examples of pH values

Measuring pH

pH meter Litmus paper Indicators

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Calculating pHThe pH of a solution can be calculated if the [H+] concentration is

known from the following formula:

pH = -log[H+]

It is a logarithmic scale in which a pH change of 1 corresponds to a factor of 10 change in [H+] concentration

When calculating pH values we need to remember that the number of decimal places in the pH is equal to the number of

significant figures in the concentration

Example:

[H+] = 1.00 x 10-8 M (3 sf)

pH = -log[H+] = -log(1.00 x 10-8) = 8.000 (3 dp)

Other logarithmic scales:

pOH= -log[OH-]

pK = -log K

Relationship between pH and pOH Scales

Calculating [H+] from pHpH = -log[H+]

[H+] = antilog(-pH) = 10-pH

When calculating [H+] from pH values we need to remember that the number of decimal places in the pH value becomes the number of significant figures in the [H+] concentration. Its just the opposite of

calculating pH from [H+]

Example:

pH = 3.45 (2 dp)

[H+] = 10-pH = 10-3.45 = 3.5 x 10-4 M (2 sf)

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Solving Acid-Base Problems

Although acid-base problems vary greatly in complexity, the first things to do for any problem are always to:

1. Write the major species in solution2. Write balanced equations for reactions forming H+

3. Find the species that forms the most H+

Major Species in Solution

Solution components which are present in the largest amounts

Used to determine which components are important and which can be ignored

The first and most important step in solving acid-base (or any aqueous solution) problems!

Composition Major SpeciesPure Water H2OStrong Acid H+, A-, H2OWeak Acid HA, H2O

What are the major species present in 1.0 M solutions of the following:

a. hydrosulfuric acidb. perchloric acid

c. permanganic acidd. hydroiodic acide. sodium chloride

Calculating the pH of Strong Acid Solutions

The major species in solution are: H+, A- and H2O (HA and OH- are present in negligible amounts)

The acid-base equilibria present in a strong acid solution are:

HA(aq) H+(aq) + A-(aq) Ka very large

H2O(l) H+(aq) + OH-(aq) Kw very small

The large amount of H+ from the strong acid will shift the second equilibrium even further to the left so the tiny amount of H+ coming

from the autoionization of water can be safely ignored

The only important source of H+ therefore comes from the strong acid, and assuming [H+] ~ [HA]0, we can calculate the pH of the

solution

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Calculating the pH of Weak Acid Solutions

The major species in solution are: HA and H2O

The acid-base equilibria present in a weak acid solution are:

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

H2O(l) H+(aq) + OH-(aq) Kw very small

The most important source of H+ normally comes from the weak acid, allowing us to calculate [H+] and hence the pH of the solution using

the same method as used in Chapter 13 for systems with small equilibrium constants

Procedure Calculating the pH of Weak Acid Solutions

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Calculating the pH of a Mixture of Weak Acids

If one of the weak acids is much stronger than the other (Ka1 » Ka2) then we can assume that the dominant source of H+ will be from the

strongest of the two and follow the usual method

Percent Dissociation (Ionization)

Used the specify the amount of weak acid dissociated (ionized) at equilibrium:

Calculating % dissociation involves the same calculation as used in checking the 5% rule!

For solutions of any weak acid HA, H+ decreases as [HA]0 decreases, but the percent dissociation increases as [HA]0 decreases:

Percent dissociation can be used to calculate Ka!

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Base Dissociation Constant, Kb

For:

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

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

The position of the equilibrium depends on the competition between the two bases B and OH- for the proton

If B is a much stronger base than OH- the equilibrium position will be far to the right with most of the dissolved base in ionized form

If OH- is a much stronger base than B the equilibrium position will be far to the left with most of the dissolved base in molecular form

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Strong Bases

Hydroxides of Group 1A metals (Li, Na, K, Rb and Cs) are strong bases which completely dissociate in aqueous solution:

LiOH(aq) Li+(aq) + OH-(aq)

Hydroxides of Group 2A metals (Ca, Sr and Ba) are also strong bases forming 2 moles of OH- per mole base:

Ca(OH)2(aq) Ca2+(aq) + 2OH-(aq)

However, Group 2A hydroxides have low solubilities

Calculating the pH of Strong Base Solutions

The major species in solution are: M+, OH- and H2O (MOH and H+ are present in very small amounts)

The equilibria present in a strong base solution are:

MOH(aq) M+(aq) + OH-(aq) Kb very large

H2O(l) H+(aq) + OH-(aq) Kw very small

The large amount of OH- from the strong base will shift the second equilibrium even further to the left so the tiny amount of OH- coming

from the autoionization of water can be safely ignored

The only important source of OH- therefore comes from the strong base, and assuming [OH-] ~ [MOH]0, we can then calculate [H+]

followed by the pH of the solution

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Weak Bases

Most weak bases contain nitrogen atoms with lone pairs onto which protons can attach:

Calculating the pH of Weak Base Solutions

The major species in solution are: B and H2O

The acid-base equilibria present in a weak acid solution are:

B(aq) +H2O(aq) BH+(aq) + OH-(aq) Kb small

H2O(l) H+(aq) + OH-(aq) Kw very small

The most important source of OH- normally comes from the reaction of the weak base with water which allows us to calculate [OH-], [H+] and hence the pH of the solution using a similar method to that used

for weak acids

Polyprotic Acids

Acids that can produce more than one proton

Polyprotic acids always dissociate in a stepwise manner with the conjugate base of one step becoming the acid in the next step

Example for a generic triprotic acid:

H3A(aq) H+(aq) + H2A-(aq) Ka1 = [H+][H2A-]/[H3A]

H2A-(aq) H+(aq) + HA2-(aq) Ka2 = [H+][HA2-]/[H2A-]

HA2-(aq) H+(aq) + A3-(aq) Ka3 = [H+][A3-]/[HA2-]

Typically Ka1 > Ka2 > Ka3

with each acid becoming successively weaker since the increasing negative charge on the acid makes it more difficult to remove the

proton

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Since Ka1 is usually much larger than the others, only the first dissociation step makes a significant contribution to [H+] and hence

the pH

This simplifies the calculations required for most polyprotic acids

Sulfuric Acid (H2SO4)

Sulfuric acid is unique among the polyprotic acids in that it is a strong acid in its first dissociation step and a weak acid in its second step:

All other polyprotic acids are weak acids in each step

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Summary

Acid-Base Properties of Salts

When an ionic compound dissolves in water, it dissociates into ions. These solutions can be acid, basic or neutral depending on the acidic

and basic properties of the cations and anions present

Will a solution of NaNO3 be neutral, acidic or basic?Na+ is a cation of a strong base and has no affinity for H+, nor can it

produce H+ so it have no effect on [H+] and therefore the pH

NO3- is weak conjugate base of a strong acid so it WILL NOT remove H+ from water since it prefers to remain dissociated as NO3

-(aq)

The solution forms neither H+ or OH- and will therefore be neutral with a pH of 7

Salts that produce Neutral SolutionsA solution of a salt containing the cation from a strong base and the

anion from a strong acid will be neutral

Examples:

Alkali metal chlorides, bromides, iodides, nitrates, sulfates and perchlorates

Will a solution of NaC2H3O2 be neutral, acidic or basic?

Na+ is a cation of a strong base and has no affinity for H+, nor can it produce H+ so it have no effect on [H+] and therefore the pH

C2H3O2- is a strong conjugate base of a weak acid and so it will

remove H+ from water forming OH- since it prefers to remain in its molecular form:

C2H3O2-(aq) + H2O(l) HC2H3O2(aq) + OH-(aq)

The solution forms OH- and will therefore be basic with a pH of >7

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Salts that produce Basic SolutionsA solution of a salt containing the cation from a strong base and the

anion from a weak acid will be basic

Examples:

Alkali metal fluorides, cyanides, carbonates, nitrites, acetates, sulfides, phosphates

Will a solution of NH4Cl be neutral, acidic or basic?

Since Cl- is a weak conjugate base of a strong acid it will not remove H+ from water since it prefers to remain dissociated as Cl-(aq)

Since NH4+ is a strong conjugate acid of a weak base it will donate H+

to water since it prefers to remain in the form of NH3 molecules:

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

The solution forms H+ and will therefore be acidic with a pH of <7

Salts that produce Acidic SolutionsA solution of a salt containing the cation from a weak base and the

anion from a strong acid will be acidic

Examples:

Ammonium compounds

Relationship between Ka and Kb for Weak Acids

For a weak acid:

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

Ka = [H+][A-]/[HA]

For the conjugate base of this weak acid:

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

Kb = [HA][OH-]/[A-]

Multiplying Ka and Kb:

Ka x Kb = ([H+][A-]/[HA]) x ([HA][OH-]/[A-])= [H+] [OH-] = Kw = 1.0 x 10-14

For any weak base and its conjugate acid:

Ka x Kb = Kw

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Hydrated Cations

Salts that contain small, highly charged metal ions (Al3+, Fe3+, Cr3+) in combination with anions from strong acids will also be acidic since

the hydrated cations formed in water are weak acids:

Al(H2O)63+(aq) Al(OH)(H2O)5

2+ + H+(aq)

Will a solution of NH4F be neutral, acidic or basic?

Since F- is a strong conjugate base of a weak acid it will try to remove H+ from water since it prefers to remain undissociated as

HF(aq)

Since NH4+ is a strong conjugate acid of a weak base it wants to

donate H+ to water since it prefers to remain in the form of NH3(aq) molecules

It is impossible to tell from this information alone since it will depend on the relative strengths of the acid (NH4

+) and the base (F-)

Solutions of salts containing the cations of weak bases and the anions of weak acids can be acidic, basic or depending on the relative size of

Ka for the acidic ion compared to Kb for the basic ion

Their pH’s can be predicted qualitatively as follows:

Quantitative calculations on these salts are complicated!

Predict whether an aqueous solutions of the following salts will be acidic, basic or neutral:

NH4C2H3O2Ka for NH4

+ = 5.6 x 10-10, Kb for C2H3O2- = 5.6 x 10-10

NH4CNKa for NH4

+ = 5.6 x 10-10, Ka for HCN = 6.3 x 10-10

Al2(SO4)3Ka for Al(H2O)6

3+ = 1.4 x 10-5, Ka2 for H2SO4 = 1.2 x 10-2

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This method can in fact be used to determine the qualitative pH of any salt solution!

Classification of Cations and Anions in Solution Summary

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Effect of Structure on Acid-Base Properties

Any molecule containing a hydrogen atom can potentially act as an acid

There are two factors which determine whether a molecule containing an X-H will behave as an acid:

1. Bond strength2. Bond polarity

The stronger the bond and the lower the polarity the lower the probability that the proton will be released

C-H bonds are both strong and non-polar so molecules containing these bonds do not show acidic properties

Hydrogen halides on the other hand have polar bonds:

H-F > H-Cl > H-Br > H-Imost polar least polar

Since the H-F bond is the most polar we might expect that HF would be the strongest acid, but it is in fact a weak acid due the fact

that the H-F bond is very strong

Oxyacids are very common and contain the H-O-X functional group

For these acids, their Ka values increase as the number of electronegative oxygen atoms attached to the central X atoms

increase since they are able to draw electrons away from the O-H bond and thereby weaken it:

This is also true for electronegative atoms other than oxygen:

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Acid-Base Properties of Oxides

When covalent oxides dissolve in water they produce acidic solutions since the strong covalent bonds within the molecule remain intact

while the O-H bonds in water break to produce protons:

These compounds are called acidic oxides

When ionic oxides dissolve in water they produce basic solutions:

This occurs because the oxide ion reacts with water to form the hydroxide ion:

These compounds are called basic oxides

The Lewis Acid-Base Model

G. N. Lewis

The Lewis model is even more general than the Brønsted-LowryModel:

A Lewis Acid is an electron-pair acceptor since it has an empty atomic orbital that it can use to accept (share) from another molecule

which has a lone pair

A Lewis Base is an electron-pair donor since it has a lone pair of electrons which can be donated (shared) with another molecule which

has an empty atomic orbital

The Lewis Model has the advantage that it covers many reactions which do not involve Brønsted-Lowry acids:

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Acid-Base Theories Compared