1

Introduction to Acids and Bases

The earliest definition was given by Arrhenius:

• An acid contains a hydrogen atom and dissolves in water to form a hydrogen ion, H+.

• A base contains hydroxide and dissolves in water to form OH −.

HCl(g)

acid

H+(aq) + Cl−(aq)

NaOH(s)

base

Na+(aq) + OH −(aq)

2

Introduction to Acids and Bases

• The Arrhenius definition correctly predicts the behavior of many acids and bases.

• However, this definition is limited and sometimes inaccurate.

• For example, H+ does not exist in water. Instead, it reacts with water to form the hydronium ion, H3O+.

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

hydrogen ion:does not really exist

in solution

hydronium ion:actually present inaqueous solution

3

Introduction to Acids and Bases

The Brønsted–Lowry definition is more widely used:

• A Brønsted–Lowry acid is a proton (H+) donor.

• A Brønsted–Lowry base is a proton (H+) acceptor.

H3O+(aq) + Cl−(aq)HCl(g) + H2O(l)

This proton is donated.

• HCl is a Brønsted–Lowry acid because it donates a proton to the solvent water.

• H2O is a Brønsted–Lowry base because it accepts a proton from HCl.

4

Introduction to Acids and Bases Brønsted–Lowry Acids

• A Brønsted–Lowry acid must contain a hydrogen atom.

• Common Brønsted–Lowry acids (HA):

HClhydrochloric acid

HBrhydrobromic acid

H2SO4

sulfuric acid

HNO3

nitric acid

5

Introduction to Acids and Bases Brønsted–Lowry Bases

• A Brønsted–Lowry base is a proton acceptor, so it must be able to form a bond to a proton.

• A base must contain a lone pair of electrons that can be used to form a new bond to the proton.

N

H

HH + H2O(l) N

H

HH

H +

+ OH −(aq)

Brønsted–Lowry base

This e− pair forms a newbond to a H from H2O.

NH3

ammonia

6

Introduction to Acids and Bases Brønsted–Lowry Bases

• Common Brønsted–Lowry Bases(B )

NaOHsodium hydroxide

KOHpotassium hydroxide

Mg(OH)2

magnesium hydroxide

Ca(OH)2

calcium hydroxide

H2Owater

Lone pairs make theseneutral compounds bases.

The OH − is the base in each metal salt.

7

Proton Transfer The Reaction of a Brønsted–Lowry Acid with

a Brønsted–Lowry Base

H A + B A − H B++

gain of H+

acid base

loss of H+

This e− pairstays on A.

This e− pair formsa new bond to H+.

8

Proton Transfer The Reaction of a Brønsted–Lowry Acid with

a Brønsted–Lowry Base

• The product formed by loss of a proton from an acid is called its conjugate base.

H A + B A − H B++

gain of H+

acid base conjugatebase

conjugateacid

loss of H+

• The product formed by gain of a proton by a base is called its conjugate acid.

9

Proton Transfer The Reaction of a Brønsted–Lowry Acid with

a Brønsted–Lowry Base

H Br + +

gain of H+

acid base conjugatebase

conjugateacid

loss of H+

H2O Br− H3O+

• HBr and Br− are a conjugate acid–base pair.

• H2O and H3O+ are a conjugate acid–base pair.

• The net charge must be the same on both sides of the equation.

10

Proton Transfer The Reaction of a Brønsted–Lowry Acid with

a Brønsted–Lowry Base

Amphoteric compound: A compound that contains both a hydrogen atom and a lone pair of e−; it canbe either an acid or a base.

H O H

H2O as a base

add H+

H O H

H +

conjugate acid

H O H

H2O as an acid

remove H+

H O−

conjugate base

11

Acid and Base StrengthRelating Acid and Base Strength

• When an acid dissolves in water, the proton transfer that forms H3O+ is called dissociation.

• When a strong acid dissolves in water, 100% of the acid dissociates into ions.

• Common strong acids are HI, HBr, HCl, H2SO4, and HNO3.

• A single reaction arrow is used, because the product is greatly favored at equilibrium.

H3O+(aq) + Cl−(aq) HCl(g) + H2O(l)

12

Acid and Base StrengthRelating Acid and Base Strength

• When a weak acid dissolves in water, only a small fraction of the acid dissociates into ions.

• Unequal reaction arrows are used, because the reactants are usually favored at equilibrium.

H3O+(aq) + CH3COO−(aq) CH3COOH(l) + H2O(l)

13

Acid and Base StrengthRelating Acid and Base Strength

• When a strong base dissolves in water, 100% of the base dissociates into ions.

Na+(aq) + OH−(aq) NaOH(s) + H2O(l)

• Common strong bases are NaOH and KOH.

• When a weak base dissolves in water, only a small fraction of the base dissociates into ions.

NH4+(aq) + OH−(aq) NH3(g) + H2O(l)

14

Acid and Base StrengthRelating Acid and Base Strength

• A strong acid readily donates a proton, forming a weak conjugate base.

HClstrong acid

Cl−

weak conjugate base

• A strong base readily accepts a proton, forming a weak conjugate acid.

OH−

strong baseH2O

weak conjugate acid

15

Dissociation of Water

H O H

base

H O H

H +

conjugate acid

H O H

acid conjugate base

HO−

+ +

loss of H+

gain of H+

Water can behave as both a Brønsted–Lowry acidand a Brønsted–Lowry base. Thus, two water molecules can react together in an acid–base reaction:

16

Dissociation of Water

• Pure water contains an exceedingly low concentration of ions, H3O+ and –OH. Since one H3O+ ion and one –OH ion are formed in each reaction, their concentrations are equal in pure water.

• Multiplying these concentrations together gives the ion-product constant for water, Kw.

Kw = [H3O+][OH−]

ion-productconstant

[H3O+] = [OH−] = 1.0 × 10–7 M at 25 °C.

17

Dissociation of Water

• Substituting the concentrations for H3O+ and –OH into the expression for Kw gives the following result.

Kw = [H3O+] [OH−]

Kw = (1.0 x 10−7) x (1.0 x 10−7)

Kw = 1.0 x 10−14

• Kw is a constant, 1.0 x 10−14, for all aqueous solutions at 25 oC.

18

Dissociation of Water

To calculate [−OH] when [H3O+] is known:

To calculate [H3O+] when [−OH] is known:

Kw = [H3O+][OH−] Kw = [H3O+][OH−]

[OH−] =1 x 10−14

[H3O+]

[OH−] =Kw

[H3O+]

[OH−]=

1 x 10−14

[H3O+]

[OH−]=

Kw[H3O+]

19

Dissociation of Water

If the [H3O+] in a cup of coffee is 1.0 x 10−5 M, then the [−OH] can be calculated as follows:

[OH−] =Kw

[H3O+]=

1 x 10−14

1 x 10−5= 1.0 x 10−9 M

In this cup of coffee, therefore, [H3O+] > [OH−], and the solution is acidic overall.

20

Dissociation of Water

21

The pH ScaleCalculating pH

pH = −log [H3O+]

The lower the pH, the higher the concentration of H3O+.

• Acidic solution: pH < 7 [H3O+] > 1 x 10−7

• Basic solution: pH > 7 [H3O+] < 1 x 10−7

• Neutral solution: pH = 7 [H3O+] = 1 x 10−7

22

The pH ScaleCalculating pH from [H3O+]

• If [H3O+] = 1.0 x 10–5 M for a urine sample, what is its pH?

pH = –log [H3O+] = –log (1.0 x 10–5)

= –(–5.00) = 5.00

• The urine sample is acidic because the pH < 7.

23

The pH ScaleCalculating [H3O+] from pH

• If the pH of seawater is 8.50, what is the [H3O+]?

pH = −log [H3O+]

8.50 = −log [H3O+]

−8.50 = log [H3O+]

antilog (−8.50 ) = [H3O+]

[H3O+] = 3.2 x 10−9 M

• The seawater is basic because [H3O+] > 1 x 10–7 M.

24

The pH ScaleCalculating pH

• A logarithm has the same number of digits to the right of the decimal point as are contained in the coefficient of the original number.

[H3O+] = 3.2 x 10−9 M

two significant figures

pH = 8.50

two digits afterdecimal point

pH = 8.50

25

Focus on the Human BodyThe pH of Body Fluids

26

Common Acid–Base ReactionsReaction of Acids with Hydroxide Bases

Neutralization reaction: An acid-base reaction thatproduces a salt and water as products.

HA(aq) + MOH(aq)

acid baseOH(l) + MA(aq)H

water salt

• The acid HA donates a proton (H+) to the OH− base to form H2O.

• The anion A− from the acid combines with the cation M+ from the base to form the salt MA.

27

Common Acid–Base Reactions

HOW TO Draw a Balanced Equation for a NeutralizationReaction Between HA and MOH

Example Write a balanced equation for the reaction of Mg(OH)2 with HCl.

Step [1] Identify the acid and base in the reactantsand draw H2O as one product.

HCl(aq) + Mg(OH)2(aq)

acid base

H2O(l) +water

salt

28

Common Acid–Base Reactions

HOW TO Draw a Balanced Equation for a NeutralizationReaction between HA and MOH

Step [2] Determine the structure of the salt.

• The salt is formed from the parts of the acid and base that are not used to form H2O.

HCl

H+

reacts to form H2O

Cl−

used to form salt

Mg(OH)2

Mg2+

used to form salt

2 OH−

react to form water

Mg2+ and Cl− combine to form MgCl2.

29

Common Acid–Base Reactions

HOW TO Draw a Balanced Equation for a NeutralizationReaction between HA and MOH

Step [3] Balance the equation.

HCl(aq) + Mg(OH)2(aq)

acid base

H2O(l) +

water salt

MgCl22

Place a 2 tobalance Cl.

Place a 2 tobalance O and H.

2

30

Common Acid–Base ReactionsReaction of Acids with Hydroxide Bases

A net ionic equation contains only the species involved in a reaction.

HCl(aq) + NaOH(aq) H—OH(l) + NaCl(aq)

• Written as individual ions:

H+(aq) + Cl−(aq) + Na+(aq) + OH−(aq)

H—OH(l) + Na+(aq) + Cl−(aq)

• Omit the spectator ions, Na+ and Cl–.

H+(aq) + −OH(aq) H—OH(l)

• What remains is the net ionic equation:

31

Common Acid–Base ReactionsReaction of Acids with Bicarbonate Bases

• A bicarbonate base, HCO3−, reacts with one H+ to

form carbonic acid, H2CO3.

H+(aq) + HCO3−(aq)

• Carbonic acid then decomposes into H2O and CO2.

H2O(l) + CO2(g)

H2CO3(aq)

HCl(aq) + NaHCO3(aq)

H2O(l) + CO2(g)

NaCl(aq) + H2CO3(aq)

• For example:

32

Common Acid–Base ReactionsReaction of Acids with Bicarbonate Bases

• A carbonate base, CO32–, reacts with two H+ to

form carbonic acid, H2CO3.

2 H+(aq) + CO32–(aq)

H2O(l) + CO2(g)

H2CO3(aq)

2 HCl(aq) + CaCO3(aq)

H2O(l) + CO2(g)

2 CaCl2(aq) + H2CO3(aq)

• For example:

33

Buffers

A buffer is a solution whose pH changes very littlewhen acid or base is added.

Most buffers are solutions composed of roughlyequal amounts of

• a weak acid.

• the salt of its conjugate base.

The buffer resists change in pH because

• added base, −OH, reacts with the weak acid.

• added acid, H3O+, reacts with the conjugate base.

34

BuffersGeneral Characteristics of a Buffer

CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO−(aq)

weak acid conjugatebase

If an acid is added to the following buffer equilibrium,

Adding moreproduct…

…drives the reaction to the left.

then the excess acid reacts with the conjugate base, so the overall pH does not change much.

35

BuffersGeneral Characteristics of a Buffer

CH3COOH(aq) + −OH(aq) H2O(l) + CH3COO−(aq)

weak acid conjugatebase

If a base is added to the following buffer equilibrium,

Adding morereactant…

…drives the reaction to the right.

then the excess base reacts with the weak acid, so the overall pH does not change much.

36

Focus on the Human BodyBuffers in the Blood

• Normal blood pH is between 7.35 and 7.45.

• The principle buffer in the blood is carbonic acid/ bicarbonate (H2CO3/HCO3

−).

CO2(g) + H2O(l) H2CO3(aq)H2O

H3O+(aq) + HCO3−(aq)

• CO2 is constantly produced by metabolic processes in the body.

• The amount of CO2 is related to the pH of the blood.

37

Focus on the Human BodyBuffers in the Blood

Respiratory acidosis results when the body fails toeliminate enough CO2, due to lung disease or failure.

CO2(g) + 2 H2O(g) H3O+(aq) + HCO3−(aq)

A lower respiratory rate increases [CO2].

This drives the reaction to the right,increasing [H3O+].

Blood then has a higher [H3O+] and a lower pH.

38

Focus on the Human BodyBuffers in the Blood

Respiratory alkalosis is caused by hyperventilating;very little CO2 is produced by the body.

CO2(g) + 2 H2O(g) H3O+(aq) + HCO3−(aq)

A higher respiratory rate decreases [CO2].

This drives the reaction to the left,decreasing [H3O+].

Blood then has a lower [H3O+] and a higher pH.