Shawn McDonald

Linn-Benton Community College

CHEM, 2nd edition

Cengage Learning

Chapter 9

Acid-base reactionsAcids and bases are chemical

compounds that occur regularly in

'everyday life'. These two types of

substances have opposite properties.

They often occur in the foods we eat.

2

Types of Electrolytes

• Salts are water-soluble ionic compounds.

All are strong electrolytes. Example: NaCl

• Acids form H+1 ions in water solution.

• Bases combine with H+1 ions in water solution.

Bases increase the OH-1 concentration of the solution.

May either directly release OH-1 or pull H+1 off H2O

molecule. If the latter, this forms OH- ion and a different

positively charged ion.

3

Properties of Acids• Sour taste. Like biting into a lemon...

• React with “active” metals.

I.e., Al, Zn, Fe, but not Cu, Ag or Au.

2 Al + 6 HCl AlCl3 + 3 H2

Corrosive. To Al (mid-right) and skin!

• React with carbonates, producing CO2.

Marble, baking soda, chalk, limestone.

CaCO3 + 2 HCl CaCl2 + CO2(g) + H2O

• Change color of vegetable dyes.

Blue litmus turns red. Picture at right.

• React with bases to form ionic salts

and water. Called a neutralization.

4

Common AcidsChemical name Formula Uses Strength

Nitric acid HNO3 Explosive, fertilizer, dye, glue Strong

Sulfuric acid H2SO4 Explosive, fertilizer, dye, glue,

batteries Strong

Hydrochloric acid HCl Metal cleaning, food prep, ore

refining, stomach acid Strong

Phosphoric acid H3PO4 Fertilizer, plastics and rubber,

food preservation Moderate

Acetic acid HC2H3O2 Plastics and rubber, food

preservation, vinegar Weak

Hydrofluoric acid HF Metal cleaning, glass etching Weak

Carbonic acid H2CO3 Soda water Weak

Boric acid H3BO3 Eye wash Weak

5

Structures of Acids

• Binary acids have acid

hydrogens attached to a

nonmetal atom. 2 types

of elements only.

HCl, HF

Write the H atom first,

then the nonmetal atom.

Dissociate in water to

form H+ ions and

nonmetal anions (such as

Cl- or F-)

Hydrofluoric acid

6

Structure of Acids• Oxyacids have acid

hydrogens attached to an

oxygen atom.

H2SO4, HNO3

Also write the H atom(s)

first, then the polyatomic

ion group.

Will dissociate when put

into water, to give H+ ions

and a polyatomic anion

(like NO3- or SO4

2-)

7

Structure of Acids

• Carboxylic acids have COOH group.

HC2H3O2, H3C6H5O3

• Only the first H in the formula is acidic.

The H is on the COOH.

Lemons and limes Apples and

wine

Component of vinegar

8

Properties of Bases• Also known as alkalis.

• Taste bitter.Alkaloids = Plant product that is

alkaline.

Often poisonous. Potato and tomato shoots.

• Solutions feel slippery.

• Change color of vegetable dyes.Different color than acid.

Red litmus turns blue. The dye in the paper reacts with OH- from base.

• React with acids to form ionic salts.Neutralization. Negates taste and metal

dissolving power of acids.

Occurs in hemlock bark,

repels beetles.

9

Common BasesChemical

name Formula

Common

name Uses Strength

Sodium

hydroxide NaOH

Lye,

caustic soda

Soap, plastic,

petrol refining Strong

Potassium

hydroxide KOH

Caustic

potash

Soap, cotton,

electroplating Strong

Calcium

hydroxide Ca(OH)2 Slaked lime Cement Strong

Sodium

bicarbonate NaHCO3 Baking soda Cooking, antacid Weak

Magnesium

hydroxide Mg(OH)2

Milk of

magnesia Antacid Weak

Ammonium

hydroxide

NH4OH,

{NH3(aq)}

Ammonia

water

Detergent,

fertilizer,

explosives, fibers

Weak

10

Structure of Bases

• Most ionic bases contain

OH ions.

NaOH, Ca(OH)2

• Some contain CO32- ions.

CaCO3 NaHCO3

• Molecular bases contain

structures that react with

H+.

Mostly amine groups (N

atoms).

Caffeine has three amine type

groups with CH2 group

attached.

11

9-1b What is an acid or a base?

• An acid–base reaction is any reaction in

which an H+ is transferred. Does not have to take place in aqueous solution.

Broader definition than Arrhenius.

• An acid is a H+ donor; A base is a H+

acceptor. Either can be a molecule or an ion. Since H+ is a proton, acid is a proton donor and

base is a proton acceptor.

Base structure must contain an atom with an

unshared pair of electrons to bond to H+.

• In the reaction, the acid molecule gives an H+

to the base molecule.

H–A + :B :A– + H–B+

12

Amphoteric Substances• Amphoteric substances can act as either an acid or

a base.They have both a transferable H atom and an atom with a

lone pair.

• HCl(aq) is acidic because HCl transfers an H+ to H2O, forming H3O

+ ions.Water acts as base, accepting H+.

HCl(aq) + H2O(l) → Cl–(aq) + H3O

+(aq)• NH3(aq) is basic because NH3 accepts an H

+ from H2O, forming OH

–(aq).Water acts as acid, donating H+.

NH3(aq) + H2O(l) NH4+(aq) + OH–(aq)

Thus water is amphoteric, it can act as a base with an acid, or as an acid with a base. Its nature is the opposite of the

compound with which it is interacting.

13

An example acid-base reaction

In the reaction H2O + NH3 HO– + NH4

+:

water ammonia hydroxide ion ammonium

ion

H2O and HO– constitute an

acid/conjugate–base pair. If

hydroxide ion accepts a proton it

will revert to a water molecule.

NH3 and NH4+ constitute a

base/conjugate–acid pair. If

the ammonium ion donates

a proton to a base, it will

revert to the ammonia

molecule.

14

Example—Identify the Brønsted–Lowry acids and bases and their conjugates in this reaction.

C. HNO3(aq) + H2O(l) H3O+(aq) + NO3

-(aq)

15

Neutralization Reactions• H+ + OH- H2O net ionic eqn.

• Acid + base salt + water

• Double-displacement reactions.Salt = cation from base + anion from

acid. Sometimes the salt in insoluble.

Cation and anion charges stay constant.

H2SO4 + Ca(OH)2 → CaSO4 + 2 H2O

• Some neutralization reactions are gas evolving, where H2CO3 (carbonic acid) decomposes into CO2 and H2O.

H2SO4 + 2 NaHCO3 → Na2SO4 + 2 H2O + 2 CO2sulfuric acid sodium bicarbonate sodium sulfate water carbon dioxide gas

16

Example — Write the equation for the reaction of

aqueous perchloric acid with strontium hydroxide.

1. Write the formulas of the reactants.

HClO4(aq) + Sr(OH)2(aq)

2. Determine the ions present when each reactant

dissociates.

(H+ + ClO4−) + (Sr2+ + OH−)

3. Exchange the ions. H+1 combines with OH-1 to

make H2O(l). Like other double displacements.

(H+ + ClO4−) + (Sr2+ + OH−) (Sr2+ + ClO4

−) + H2O(l)

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4. Write the formulas of the products.

Cross charges and reduce subscripts if possible.

HClO4(aq) + Sr(OH)2(aq) Sr(ClO4)2 + H2O(l)

5. Balance the equation. Each atom and group on

left vs. right side of the equation.

May be quickly balanced by matching the numbers of

H and OH to make H2O.

Coefficient of the salt is always 1.

2 HClO4(aq) + Sr(OH)2(aq) Sr(ClO4)2 + 2 H2O(l)

Write the equation for reaction of aqueous

perchloric acid with strontium hydroxide,

Continued.

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6. Determine the solubility of the salt.

Sr(ClO4)2 is soluble (look up in solubility table).

7. Write an (s) after the insoluble products and an

(aq) after the soluble products.

2 HClO4(aq) + Sr(OH)2(aq) Sr(ClO4)2(aq) + 2 H2O(l)

Write the Equation for reaction of aqueous

perchloric acid with strontium

hydroxide.Continued

The reaction occurs since one of the products formed is water,

and water molecules mainly stay as molecules and don't ionize

very much.

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9-2 Strong or Weak

• A strong acid is a strong electrolyte.Practically all the acid molecules ionize, →. completely

• A strong base is a strong electrolyte.Practically all the base molecules form OH– ions, either

through dissociation or reaction with water, →. completely

• A weak acid is a weak electrolyte.Only a small percentage of the molecules ionize, .

• A weak base is a weak electrolyte.Only a small percentage of the base molecules form OH–

ions, either through dissociation or reaction with water, .

20

Strong Acids• The stronger the acid, the

more willing it is to donate H.

Use water as the standard base.

• Strong acids donate practically all their H’s.

100% ionized in water.

Strong electrolyte.

• [H3O+] = [strong acid].

[ ] = molarity.

HCl H+ + Cl-

HCl + H2O H3O+ + Cl-

No HCl is left after you dissolve

the compound in water.

21

Strong Acids, Continued

Hydrochloric acid HCl

Hydrobromic acid HBr

Hydroiodic acid HI

Nitric acid HNO3

Perchloric acid HClO4

Sulfuric acid H2SO4

22

Strong Acids, Continued

Pure waterHCl solution

HCl is a strong electrolyte. Its

solution will conduct current.Water is not an electrolyte and

will not conduct current.

23

Weak Acids

• Weak acids donate a small

fraction of their Hs.

Most of the weak acid

molecules do not donate H

to water.

Often less than 1% ionized

in water.

• [H3O+]

24

Weak Acids, Continued

Hydrofluoric acid HF

Acetic acid HC2H3O2

Formic acid HCHO2

Sulfurous acid H2SO3

Carbonic acid H2CO3

Phosphoric acid H3PO4

Stronger attraction between H and F than between H+ and water

25

Weak Acids, Continued

Pure water HF solution

HF in water

is a weak

electrolyte

and only

conducts

electricity

poorly – note

dimness of

light bulb.

26

Degree of Ionization

• The extent to which an acid ionizes in water

depends in part on the strength of the bond

between the acid H+ and anion compared

to the strength of the bond between the

acid H+ and the O of water.

HA(aq) + H2O(l) A−(aq) + H3O

+(aq)

In other words, which is stronger? the H-A bond or the

H-O bond in the hydronium ion. If the former, then the

acid will only ionize slightly. Fluorine for example

forms a stronger bond with H than does Cl, thus HF is a

weak acid and HCl is a strong acid.

27

Strong Bases• The stronger the base, the more

willing it is to accept H.

Use water as the standard acid.

• Strong bases, practically all molecules are dissociated into OH– or accept Hs.

Strong electrolyte.

Multi-OH bases completely dissociated.

• [HO–] = [strong base] x (# OH).

• Molarity will be discussed shortly....

NaOH Na+ + OH-

All the dissolved NaOH has

dissociated into sodium and

hydroxide ions.

28

Strong Bases, Continued

Lithium hydroxide LiOH

Sodium hydroxide NaOH

Potassium hydroxide KOH

Calcium hydroxide Ca(OH)2

Strontium hydroxide Sr(OH)2

Barium hydroxide Ba(OH)2

29

Weak Bases

• In weak bases, only a small

fraction of molecules accept

Hs.

Weak electrolyte.

Most of the weak base molecules

do not take H from water.

Much less than 1% ionization in

water.

• [OH–]

30

Weak Bases, Continued

Ammonia NH3(aq) + H2O(l) NH4+(aq) + OH−(aq)

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

Methyl amine CH3NH2(aq) + H2O(l) CH3NH3+(aq) + OH−(aq)

Ethyl amine C2H5NH2(aq) + H2O(l) C2H5NH3+(aq) + OH−(aq)

Bicarbonate HCO3−(aq) + H2O(l) H2CO3 (aq) + OH

−(aq)

Most of the base molecules or species stay in the unprotonated form.

31

Autoionization of Water• Water is actually an extremely weak electrolyte.

Therefore, there must be a few ions present.

• About 1 out of every 10 million water molecules form

ions through a process called autoionization.

H2O + H2O H3O+ + OH–

• All aqueous solutions contain both H3O+ and OH–.

The concentration of H3O+ and OH– are equal in DI water.

[H3O+] = [OH–] = 1 x 10-7M at 25 °C in pure water. These are

important concentrations to remember (related to pH that will

we study shortly).

32

Ion Product of Water• The product of the H3O

+ and OH–

concentrations is always the same number for solutions at 25 Celsius.

• The number is called the ion product of water and has the symbol Kw.

• [H3O+] x [OH–] = 1 x 10-14 = Kw.

• As [H3O+] increases, the [OH–] must

decrease so the product stays constant.

Inversely proportional.

33

Acidic and Basic Solutions

• Neutral solutions have equal [H3O+] and [OH–].

[H3O+] = [OH–] = 1 x 10-7

• Acidic solutions have a larger [H3O+] than [OH–].

[H3O+] > 1 x 10-7; [OH–] < 1 x 10-7

• Basic solutions have a larger [OH–] than [H3O+].

[H3O+] < 1 x 10-7; [OH–] > 1 x 10-7

We can measure the concentration of hydronium ions in a

solution by using a pH meter. These devices are used in

the general chemistry courses to study acid-base behavior

and reactions. pH paper also works to give us the rough

pH of a given aqueous solution.

34

Example—Determine the [H3O+] for a 0.00020 M

Ba(OH)2 solution and Determine Whether the

Solution Is Acidic, Basic, or Neutral.

Ba(OH)2 = Ba2+ + 2 OH– therefore:

[OH–] = 2 x 0.00020 = 0.00040 = 4.0 x 10−4 M

4

14

3

3

100.4

101

OHOH

OHOH

w

w

K

K

[H3O+] = 2.5 x 10-11 M.

Since [H3O+] < 1 x 10−7, the solution is basic.

35

9-3a The pH scale

• The acidity/basicity of a solution is often

expressed as pH.

• pH = ─log[H3O+], [H3O

+] = 10−pH

The exponent on 10, but with a positive sign.

pHwater = −log[10-7] = 7.

Need to know the [H3O+] concentration to find pH.

• 3 cases:

case 1: pH < 7 is acidic; case 2: pH > 7 is basic;

case 3: pH = 7 is neutral.

36

pH, Continued• The lower the pH, the more acidic the solution; the

higher the pH, the more basic the solution.1 pH unit corresponds to a factor of 10 fold

difference in acidity of that solution.

• Normal range is 0 to 14.pH 0 is [H+] = 1 M, pH 14 is [OH–] = 1 M.

pH can be negative (very acidic) or larger than 14

(very alkaline, at high concentration).

37

pH of Common SubstancesSubstance pH

1.0 M HCl 0.0

0.1 M HCl 1.0

Stomach acid 1.0 to 3.0

Lemons 2.2 to 2.4

Soft drinks 2.0 to 4.0

Plums 2.8 to 3.0

Apples 2.9 to 3.3

Cherries 3.2 to 4.0

Unpolluted rainwater 5.6

Human blood 7.3 to 7.4

Egg whites 7.6 to 8.0

Milk of magnesia (saturated Mg(OH)2) 10.5

Household ammonia 10.5 to 11.5

1.0 M NaOH 14

38

Example—Calculate the pH of a 0.0010 M

Ba(OH)2 Solution and Determine if It Is

Acidic, Basic, or Neutral.

[H3O+] =

1 x 10-14

2.0 x 10-3= 5.0 x 10-12M

pH > 7 therefore, basic.

Ba(OH)2 = Ba2+ + 2 OH− therefore,

[OH-] = 2 x 0.0010 = 0.0020 = 2.0 x 10-3 M.

pH = −log [H3O+] = −log (5.0 x 10-12)

pH = 11.3

39

Practice—Calculate the pH of the

Following Strong Acid or Base

Solutions.• 0.0020 M HCl

• 0.0050 M Ca(OH)2

• 0.25 M HNO3

40

[H3O+] = 1 x 10

−14

1 x 10−2= 1.0 x 10−12

Practice—Calculate the pH of the

Following Strong Acid or Base

Solutions, Continued.

• 0.0020 M HCl strong acid therefore, [H3O+] =

0.0020 M.

• 0.0050 M Ca(OH)2 strong base, [OH–] = 0.010 M.

• 0.25 M HNO3 a strong acid, therefore, [H3O+] =

0.25 M.

pH = −log (2.0 x 10-3) = 2.70 acidic

pH = −log (1.0 x 10−12) = 12.00 basic

pH = −log (2.5 x 10−1) = 0.60 acidic

41

9-4 Acid-base buffers

• Buffers are solutions that resist changing pH when small amounts of acid or base are added.

• They resist changing pH by neutralizingadded acid or base.

• Buffers are made by mixing together a weak acid and its conjugate base.

Or weak base and its conjugate acid. For example, a mixture of ammonia and ammonium chloride is a buffer.

42

How Buffers Work

• The weak acid present in the buffer mixture

can neutralize added base.

• The conjugate base present in the buffer

mixture can neutralize added acid.

• The net result is little to no change in the

solution pH. As long as you don’t over

come the capacity of the buffer. Then the

pH will change drastically. Can change by

several pH units after capacity is overcome.

43

Acetic Acid/Acetate Buffer

The conjugate base

neutralizes any added

acid…. makes weak acid

(acetic acid)

The weak acid neutralizes

the added base, forms H2O.