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Chapter 15: Acids and Bases

Mrs. Brayfield

Heart Burn

Why do people take milk of magnesia or other

medication for heart burn?

Well, what is heart burn?

It is when the acid from your stomach irritates your esophagus

This is why we take medicine that is basic to “neutralize”

the acid

15.2: The Nature of Acids and Bases

Acids have a few common properties:

Sour taste (citric acid in lemons/limes)

Ability to dissolve metals (remember Mg and HCl lab?)

Turns litmus paper red

Neutralize bases

Bases also have a few common properties:

Bitter taste (antacid tablets)

Slippery feel (like the lye in soap)

Turns litmus paper blue

Neutralizes acids

Nature of Acids and Bases

Acids are much more common in food because of the

bitter taste of bases

This is because organic bases in nature are usually poisonous

Common Acids Common Bases

Hydrochloric Acid HCl Sodium Hydroxide NaOH

Sulfuric Acid H2SO4 Potassium Hydroxide KOH

Nitric Acid HNO3 Sodium Bicarbonate NaHCO3

Acetic Acid HC2H3O2 Sodium Carbonate Na2CO3

Citric Acid H3C6H5O7 Ammonia NH3

Carbonic Acid H2CO3

Hydrofluoric Acid HF

Phosphoric Acid H3PO4 Tables 15.1 & 15.2

15.3: Definitions of Acids and Bases

Swedish chemist Svante Arrhenius (1880s) developed the

first definition for acids and bases:

Acid: A substance that produces H+ ions in aqueous solns

Base: A substance that produces OH- ions in aq solns

HCl is an acid:

𝐻𝐶𝑙 𝑎𝑞 → 𝐻+ 𝑎𝑞 + 𝐶𝑙−(𝑎𝑞)

NaOH is a base:

𝑁𝑎𝑂𝐻 𝑎𝑞 → 𝑁𝑎+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞)

Definitions of Acids and Bases

When acids and bases are put in water they dissociate

(break up) into ions

These ions then interact with water:

As well as:

𝐻+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞) → 𝐻2𝑂(𝑙)


The widely accepted definition of acids and bases is called

the Brønsted-Lowry definition (1923)

Acid: Proton (H+) donor

Base: Proton (H+) acceptor

Well what does this mean?

HCl is an acid because it donates a proton:

𝐻𝐶𝑙 𝑎𝑞 + 𝐻2𝑂 𝑙 → 𝐻3𝑂+ 𝑎𝑞 + 𝐶𝑙−(𝑎𝑞)

NH3 is a base because it accepts a proton:

𝑁𝐻3 𝑎𝑞 + 𝐻2𝑂 𝑙 → 𝑁𝐻4+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞)


According to the Brønsted-Lowry definition, acids and

bases always occur together in the same reaction:

𝐻𝐶𝑙 𝑎𝑞 + 𝐻2𝑂 𝑙 → 𝐻3𝑂+ 𝑎𝑞 + 𝐶𝑙−(𝑎𝑞)

𝑁𝐻3 𝑎𝑞 + 𝐻2𝑂 𝑙 → 𝑁𝐻4+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞)

acid

(proton donor)

acid

(proton donor)

base

(proton acceptor)

base

(proton acceptor)


A substance that can act as a acid or base are called

amphoteric

Look when reactions are reversed:

𝑁𝐻4+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞) → 𝑁𝐻3 𝑎𝑞 + 𝐻2𝑂 𝑙

What was the base (NH3) has become the acid (NH4+)

and visa versa

So we call these pairs conjugate acid-base pairs

Or two substances related to each other by the transfer of a proton

acid

(proton donor)

base

(proton acceptor)


In any acid-base reaction,

A base accepts a proton and becomes a conjugate acid

An acid donates a proton and becomes a conjugate base

Definitions of Acids and Bases Ex 1

Write conjugate acid for the following:

Write the conjugate base for the following:

Base Conjugate Acid

H2O

NH3

CO32-

H2PO4-

Acid Conjugate Base

H2O

NH3

CO32-

H2PO4-

H3O+

NH4+

HCO3-

H3PO4

OH-

NH2-

Cannot be an acid

HPO42-

Definitions of Acids and Bases Ex 2

Identify the Brønsted-Lowry acid, base, conjugate acid,

and conjugate base:

𝐶5𝐻5𝑁 𝑎𝑞 + 𝐻2𝑂 𝑙 → 𝐶5𝐻5𝑁𝐻+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞)

𝐻𝑁𝑂3 𝑎𝑞 + 𝐻2𝑂 𝑙 ↔ 𝐻3𝑂+ 𝑎𝑞 + 𝑁𝑂3

−(𝑎𝑞)

Acid

Acid

Base

Base

Conjugate

Base

Conjugate

Base

Conjugate

Acid

Conjugate

Acid


Which of the following is not a conjugate acid-base pair?

A. (CH3)3N; (CH3)3NH+

B. H2SO4; H2SO3

C. HNO2; NO2-

Homework Problems: #1, 2, 4, 6, 8

B – two acids, not pairs

15.4: Acid Strength and Ka

Remember from chapter 4 that the strength of an

electrolyte is determined by the amount of dissociation

The more it dissociates, the stronger the electrolyte is

The same goes for acids:

A strong acid completely dissociates (or ionizes) in solution

A weak acid only partially dissociates in solution

Acid Strength

Hydrochloric acid is a strong acid:

There are only six strong acids (you MUST memorize

these):

Table 15.3: Strong Acids

Hydrochloric Acid HCl Nitric Acid HNO3

Hydrobromic Acid HBr Perchloric Acid HClO4

Hydriodic Acid HI Sulfuric Acid H2SO4

Acid Strength

HF is a weak acid:

For the strength, if the attraction between the conjugate

acid and conjugate base are weak, the acid is strong (the

forward reaction is favored)

Acid Strength

Acid Strength

If you notice there are some acids with multiple

hydrogens

This means that they have more than one ionizable protons

If the acid has two hydrogens (for example H2SO4) we

call that diprotic

If the acid has three hydrogens (for example H3PO4) we

call that triprotic

Whenever the acid undergoes a second (or third) dissociation,

the acid strength greatly decreases (we’ll see in section 6)

Acid Ionization Constant (Ka)

The strengths of weak acids are quantified with the acid

ionization constant (Ka), which is the equilibrium constant

for the ionization reaction of a weak acid:

So for the following reaction:

𝐻𝐴 𝑎𝑞 + 𝐻2𝑂 𝑙 ↔ 𝐻3𝑂+ 𝑎𝑞 + 𝐴− 𝑎𝑞

𝐻𝐴 𝑎𝑞 ↔ 𝐻+ 𝑎𝑞 + 𝐴− 𝑎𝑞

The equilibrium constant is:

𝐾𝑎 =𝐻3𝑂

+ [𝐴−]

[𝐻𝐴]=

𝐻+ [𝐴−]

[𝐻𝐴]

See page 560 (table 15.5) for values

Ka Example

Which acid is stronger:

HF Ka = 3.5x10-4

HClO Ka = 2.9x10-8

Which conjugate base is stronger?

HClO (the stronger the acid, the weaker the conjugate

base)

Homework Problems: #10, 11, 12, 14

HF

15.5: Autoionization of Water and pH

We know from previous sections that water is

amphoteric (can be an acid or a base)

Even in pure water, water acts as an acid and a base with

itself (called ionization):

𝐾𝑤 = 𝐻3𝑂+ 𝑂𝐻− = 𝐻+ [𝑂𝐻−]

Autoionization of Water

Kw is called the dissociation constant for water

Its value at 25°C is 1.0x10-14

In a neutral solution, [H+] = [OH-]

Or when those concentrations equal 1.0x10-7 (pg. 561)

In acidic solution, [H+] > [OH-]

In basic solution, [H+] < [OH-]

Autoionization of Water Example

Calculate [H+] at 25°C for each of the following solutions

and determine if the solution is basic, acidic, or neutral:

1. [OH-] = 1.5 x 10-2 M

2. [OH-] = 1.0 x 10-7 M

3. [OH-] = 8.2 x 10-10 M

[H+] = 6.74 x 10-13 M, basic

[H+] = 1.0 x 10-7 M, neutral

[H+] = 1.2 x 10-5 M, acidic

pH

The pH scale allows us to specify the acidity of a solution

𝑝𝐻 = −log[𝐻3𝑂+]

pH is reported to 2 decimal places

In general:

pH < 7, the solution is acidic

pH > 7, the solution is basic

pH = 7, the solution is neutral

pH Scale

pH scale is from 0-14

The higher the number,

the more basic it is

The pH changes by 1 for

every power of 10 change

in [H+]

pH Example 1

Calculate the pH of each solution and indicate whether

the solution is acidic or basic:

[H3O+] = 9.5 x 10-9 M

𝑝𝐻 = − log 𝐻3𝑂+ = −log(9.5 × 10−9)

𝑝𝐻 = 8.02

[OH-] = 7.1 x 10-3 M

𝐻3𝑂+ 𝑂𝐻− = 𝐾𝑤

𝐻3𝑂+ = 1.41 × 10−12𝑀

𝑝𝐻 = − log 1.41 × 10−12 = 11.85

pH Example 2

Calculate the H3O+ concentration for a solution with a

pH of 8.37.

𝑝𝐻 = − log 𝐻3𝑂+ = 8.37

𝐻3𝑂+ = 10−8.37 = 4.27 × 10−9𝑀

pOH

The pOH scale is the opposite of the pH scale:

𝑝𝑂𝐻 = −log[𝑂𝐻−]

𝑝𝐻 + 𝑝𝑂𝐻 = 14.00

The sum of pH and pOH is always equal to 14.00 at 25°C

Other p Scales

The other p scale is pKa; which is just another way to

quantify weak acid strength

The smaller pKa is, the stronger the acid

𝑝𝐾𝑎 = −log(𝐾𝑎)


15.6: Finding [H+] and pH of Strong and

Weak Acids

In every solution of acid, we have two sources of H+:


𝐻2𝑂 𝑙 + 𝐻2𝑂 𝑙 ↔ 𝐻3𝑂+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞)

Except in VERY dilute acid solutions, the ionization of

water is negligible

AKA doesn’t matter

So the only contribution of [H+] is from the acid itself

Strong and Weak Acids

Because a strong acid is completely dissociated, we can

say that the concentration of the strong acid is equal to

the concentration of H+

For example:

0.10𝑀𝐻𝐶𝑙 ⇒ 𝐻+ = 0.10𝑀 ⇒ 𝑝𝐻 = − log 0.10 = 1.00

However, we cannot make the same assumption for weak

acids

If we measured the pH of two solutions:

0.10M HCl pH = 1.00

0.10M HC2H3O2 pH = 2.87

We see that the weak acid does not fully dissociate (smaller K)

Weak Acids

To find the concentration of [H+], we must go back to

chapter 14 (equilibrium):


For the reaction above, we can summarize the

concentrations at equilibrium:

𝐾𝑎 =𝐻3𝑂

+ [𝐴−]

[𝐻𝐴]=

𝑥2

0.10 − 𝑥

[HA] [H3O+] [A-]

I 0.10 ≈ 0 0

C -x +x +x

E 0.10-x x x

Weak Acid Example 1

Find the concentration of H3O+ of a 0.250M HF solution

(Ka = 3.5 x 10-4)

𝐻𝐹(𝑎𝑞) + 𝐻2𝑂 𝑙 ↔ 𝐻3𝑂+ 𝑎𝑞 + 𝐹−(𝑎𝑞)

𝐾𝑎 =𝐻3𝑂

+ [𝐹−]

[𝐻𝐹]=

𝑥2

0.250 − 𝑥= 3.5 × 10−4

[HF] [H3O+] [F-]

I 0.250 ≈ 0 0

C -x +x +x

E 0.250 - x x x

Weak Acid Example 2

Find the pH of a 0.0150M HC2H3O2 solution (Ka = 1.8 x

10-5)

𝐻𝐶2𝐻3𝑂2(𝑎𝑞) + 𝐻2𝑂 𝑙 ↔ 𝐻3𝑂+ 𝑎𝑞 + 𝐶2𝐻3𝑂2

−(𝑎𝑞)

𝐾𝑎 =𝐻3𝑂

+ [𝐶2𝐻3𝑂2−]

[𝐻𝐶2𝐻3𝑂2]=

𝑥2

0.0150 − 𝑥= 1.8 × 10−5

[𝐇𝑪𝟐𝑯𝟑𝑶𝟐] [H3O+] [𝑪𝟐𝑯𝟑𝑶𝟐

-]

I 0.0150 ≈ 0 0

C -x +x +x

E 0.0150 - x x x

Weak Acid Example 3

Find the pH of a 0.010M HNO2 solution. (Ka = 4.6 x 10-4)

𝐻𝑁𝑂2(𝑎𝑞) + 𝐻2𝑂 𝑙 ↔ 𝐻3𝑂+ 𝑎𝑞 + 𝑁𝑂2

−(𝑎𝑞)

𝐾𝑎 =𝐻3𝑂

+ [𝑁𝑂2−]

[𝐻𝑁𝑂2]=

𝑥2

0.010 − 𝑥= 4.6 × 10−4

[𝐇𝑵𝑶𝟐] [H3O+] [𝑵𝑶𝟐

-]

I 0.010 ≈ 0 0

C -x +x +x

E 0.010 - x x x

Weak Acid Example 4

A 0.175M weak acid solution has a pH of 3.25. Find Ka for

the acid.

𝑝𝐻 = − log 𝐻3𝑂+ = 3.25 ⇒ 𝐻3𝑂

+ = 5.62 × 10−4𝑀

𝐾𝑎 =𝑥2

0.175 − 𝑥=

(5.62 × 10−4)2

0.175 − 5.62 × 10−4= 1.8 × 10−6

[HA] [H3O+] [A-]

I 0.175 ≈ 0 0

C -x +x +x

E 0.175 - x x x

Polyprotic Acids

Remember that polyprotic acids are acids with more than

one ionizable protons

In general a polyprotic acid ionizes in successive steps,

each with its own Ka; for example:

𝐻2𝑆𝑂4 𝑎𝑞 ↔ 𝐻+ 𝑎𝑞 + 𝐻𝑆𝑂4− 𝑎𝑞 𝐾𝑎1 = 1.6 × 10−2

𝐻𝑆𝑂4− 𝑎𝑞 ↔ 𝐻+ 𝑎𝑞 + 𝑆𝑂4

2− 𝑎𝑞 𝐾𝑎2 = 6.4 × 10−8

Notice that Ka2 is smaller than Ka1

This is because it is harder to remove a proton from a charged

molecule

Polyprotic Acids

Homework Problems: #24, 26, 28, 30, 32, 34, 36, 38, 42, 44,

46

15.7: Base Solutions

Just like with strong acids, strong bases also completely

dissociates in solution

There are only 6 strong bases:

As you may notice, these are group 1A and 2A metal

hydroxides

Table 15.8 Strong Bases

Lithium hydroxide LiOH Strontium hydroxide Sr(OH)2

Sodium hydroxide NaOH Calcium hydroxide Ca(OH)2

Potassium hydroxide KOH Barium hydroxide Ba(OH)2

Base Solutions

Unlike diprotic acids, bases with more than one OH-

dissociate in one step:

𝑆𝑟(𝑂𝐻)2 𝑎𝑞 → 𝑆𝑟2+ 𝑎𝑞 + 2𝑂𝐻−(𝑎𝑞)

Weak Bases

We can write a weak base dissociation just like a weak

acid dissociation:

𝐵 𝑎𝑞 + 𝐻2𝑂 𝑙 ↔ 𝐵𝐻+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞)

𝐾𝑏 =𝐵𝐻+ [𝑂𝐻−]

[𝐵]

The p scale also applies here:

𝑝𝐾𝑏 = −log𝐾𝑏

See table 15.9 (page 575) for common Kb values

pH From Strong Base Example

Find the [OH-] and pH of a 0.010M Ba(OH)2 solution.

𝐵𝑎 𝑂𝐻 2 𝑎𝑞 → 𝐵𝑎2+ 𝑎𝑞 + 2𝑂𝐻−(𝑎𝑞)

𝑂𝐻− = 2 0.010 = 0.020𝑀

𝑝𝑂𝐻 = − log 0.020 = 1.70

𝑝𝐻 = 14.00 − 𝑝𝑂𝐻 = 12.30

pH of a Weak Base Example

Find the [OH-] and pH of a 0.33M methylamine solution.

(Kb = 4.4 x 10-4)

𝐶𝐻3𝑁𝐻2 𝑎𝑞 + 𝐻2𝑂 𝑙 ↔ 𝐶𝐻3𝑁𝐻3+ 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞)

𝐾𝑏 =𝐶𝐻3𝑁𝐻3

+ [𝑂𝐻−]

[𝐶𝐻3𝑁𝐻2]=

𝑥2

.33 − 𝑥= 4.4 × 10−4

[𝑪𝑯𝟑𝑵𝑯𝟐] [𝑪𝑯𝟑𝑵𝑯𝟑+] [OH-]

I 0.33 ≈ 0 0

C -x +x +x

E 0.33 - x x x

15.8: Acid-Base Properties of Salts

We know that ions can act as an acid or base

For example:

𝐻𝐶𝑂3− 𝑎𝑞 + 𝐻2𝑂 𝑙 ↔ 𝐻2𝐶𝑂3 𝑎𝑞 + 𝑂𝐻−(𝑎𝑞)

We also know that ions are not stable by themselves

This is why we would see the bicarbonate ion bonded to

sodium in a salt:

𝑁𝑎𝐻𝐶𝑂3 𝑠 ↔ 𝑁𝑎+ 𝑎𝑞 + 𝐻𝐶𝑂3−(𝑎𝑞)

Acid-Base Properties of Salts

We can think of any anion as the conjugate base of an

acid:

And any cation as the conjugate acid of a base:

This anion is the conjugate base of this acid

Cl– HCl

F– HF

NO3– HNO3

C2H3O2– HC2H3O2

This cation is the conjugate acid of this base

NH4+ NH3

C2H5NH3+ C2H5NH2

CH3NH3+ CH3NH2

Acid-Base Properties of Salts Example

Classify the following as a weak base, weak acid, or pH-

neutral:

1. CHO2 –

2. CH3NH3+

3. F –

4. Na+

Weak base

Weak acid

Weak base

pH-neutral


Every anion may not act as a base – it depends on the

strength of the corresponding acid

An anion that is the cb of a weak acid is itself a weak base

An anion that is the cb of a strong acid is pH-neutral

Every cation may not act as a acid – it depends on the

strength of the corresponding base

A cation that is the ca of a weak base is itself a weak acid

A cation that is the ca of a strong base is pH-neutral

The stronger the starting compd, the weaker the conj. is


There are 4 rules to follow when determining if a salt is

acidic, basic, or neutral:

1. Salts in which neither the cation nor the anion acts as an acid

or base form pH-neutral solutions

2. Salts in which the cation does not act as an acid and the

anion acts as a base form basic solutions

3. Salts in which the cation acts as an acid and the anion does

not acts as a base form acidic solutions

4. Salts in which the cation acts as an acid and the anion acts as

a base form solutions in which the pH depends on the

relative strengths of the acid and the base

1. Salts with pH-neutral solutions

Some examples include:

NaCl

Na+ is neutral

Cl – is from a strong acid, so it is neutral

Ca(NO3)2

Ca2+ is neutral

NO3 – is from a strong acid, so it is neutral

KBr

K+ is neutral

Br – is from a strong acid, so it is neutral

2. Salts that form basic solutions


NaF

Na+ is neutral

F – is from a weak acid, so it is basic

Ca(C2H3O2)2

Ca2+ is neutral

C2H3O2 – is from a weak acid, so it is basic

KNO2

K+ is neutral

NO2 – is from a weak acid, so it is basic

3. Salts that form acidic solutions


FeCl3

Fe3+ is from a weak base, so it is acidic

Cl – is from a strong acid, so it is neutral

Al(NO3)3

Al3+ is from a weak base, so it is acidic

NO3 – is from a strong acid, so it is neutral

NH4Br

NH4+ is from a weak base, so it is acidic

Br – is from a strong acid, so it is neutral

*Small highly charged metal cations (like Al3+ and Fe3+) form weakly acidic

solutions

4. Salts in which it depends on strength


FeF3

Fe3+ is weakly acidic

F – is weakly basic

Al(C2H3O2)3

Al3+ is weakly acidic

C2H3O2 – is weakly basic

NH4NO2

NH4+ is weakly acidic

NO2 – is weakly basic

So are these acidic or basic?????

Strength Dependent

For those salts, we need to compare the Ka of the acid to

the Kb of the base

The ion with the higher K value dominates and determines if

the solution is acidic or basic

Salts: Acidic, Basic, or Neutral Example

Determine whether the solutions formed by each of the

following will be acidic, basic, or neutral:

NaHCO3

Basic

CH3CH2NH3Cl

Acidic

KNO3

Neutral

Fe(NO3)3

Acidic

NH4NO2

Can’t tell….

Homework Problems: #60, 62, 64, 65, 66, 68, 70

Section 9 is beyond the scope of this course

15.10: Lewis Acids and Bases

Under the Brønsted-Lowry definition of acids and bases,

ammonia (NH3) is a base because it accepts a proton

There is a better definition of acids and bases, called

Lewis acids and bases

A Lewis acid is an electron pair acceptor

A Lewis base is an electron pair donator

Lewis Acids and Bases

This new definition is why dry ice, when placed in water,

will acidify water:

A Lewis acid has an empty orbital (or can rearrange

electrons to create an empty orbital) that can accept an

electron pair

Homework Problems: #80, 82

Review Problems: #88, 94(a-c), 99

http://www.youtube.com/watch?v=ANi709MYnWg&index

=9&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr

http://www.youtube.com/watch?v=LS67vS10O5Y&index=

31&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr

http://www.youtube.com/watch?v=ANi709MYnWg&index=9&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr



http://www.youtube.com/watch?v=LS67vS10O5Y&index=31&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr



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