1. Acid-Base Theories

1.1.1.Arrhenius Acids and BasesIn 1887, according to Swedish chemist Svante Arrhenius,

acids are compounds containing hydrogen that ionize to produce hydrogen ions (H+) in aqueous solution. On the other hand, bases are compounds that ionize to produce hydroxide ions (OH-) in aqueous solution.

1.1.2.BrØnsted-Lowry Acids and BasesIn 1923, Danish chemist Johannes BrØnsted (1879-1947),

and English chemist Thomas Lowry (1874-1936) independently proposed a new theory. Br Ø nsted-Lowry acid is defined as a hydrogen-ion donor, and a Br Ø nsted-Lowry base is defined as a hydrogen-ion acceptor.

1.1.3.Lewis Acids and BasesThe third theory was proposed by Gilbert Lewis (1875-

1946). He focused on the donation or acceptance of a pair of electrons during a reaction. A Lewis acid is a substance that can accept a pair of electrons to form covalent bond. A Lewis base is a substance that can donate a pair of electrons to form covalent bond.

Table 1. Acid-Base DefinitionsType Acid Base

Arrhenius H+ producer OH- producerBrØnsted-Lowry H+ donor H+ acceptorLewis Electron-pair acceptor Electron-pair donor

2. ArrheniusArrhenius acids are compounds that contain hydrogen that ionize to

yield hydrogen (H+) ions in aqueous solution.

Monoprotic acid is an acid that contains one ionizable hydrogen. Nitric acid (HNO3) is an example. Sulfuric acid (H2SO4) or any acid that contains two ionizable protons is called diprotic acid. A tripotic acid is an acid that contains three ionizable protons like in phosphoric acid (H3PO4).

However, not all compounds containing hydrogen are acids or hydrogens in an acid are released as hydrogen ions. Only hydrogens in very polar bonds are ionizable.

Excemptions are made to methane (CH4) because the four hydrogens are in weakly C-H bonds. Aside from that, ethanoic acid (CH3COOH) is a monoprotic acid though it contains four hydrogens.

H O I IIH-C-C-O-H ethanoic acid (C CH3COOH) I H

The structural formula shows that the three hydrogens are in weakly polar bonds.

Table 2. Some Common AcidsName Formula

Hydrochloric acid HClNitric acid HNO3

Sulfuric acid H2SO4

Phosphoric acid H3PO4

Ethanoic acid CH3COOHCarbonic acid H2Co3

Arrhenius bases are compounds that ionize to yield hydroxide ions (OH-) in aqueous solution.

Table 3. Some Common BasesName Formula Solubility in

waterPotassium hydroxide KOH HighSodium hydroxide NaOH HighCalcium hydroxide Ca(OH)2 Very lowMagnesium hydroxide Mg(OH)2 Very low

Both potassium hydroxide and sodium hydroxide are ionic solids. They dissociate completely into the metal ions and hydroxide ions when dissolved in water. Elements in Group 1A (alkali metals), such as sodium and potassium, react with water to produce alkaline solutions. Calcium hydroxide [Ca(OH)2], and magnesium hydroxide [Mg(OH)2], are both hydroxides of Group 2A metals. They are not very soluble in water. Thus, their solutions are always very dilute, even when saturated. The concentration of hydroxide ions in such solutions is very low.

3. BrØnsted-Lowry

A Bronsted-Lowry acid is defined as a hydrogen-ion donor, while a Bronsted-Lowry base is a hydrogen-ion acceptor.

Let’s have ammonia as an example. When ammonia dissolves in water it acts as a base because it accepts a hydrogen ion from water.

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

In this reaction, ammonia is the Bronsted-Lowry base. Water, on the other hand, is the Bronsted-Lowry acid. Hydrogen ions are transferred from water to ammonia.

Heating an aqueous solution of ammonia drives off ammonia gas. As the ammonia gas moves out of solution, the equilibrium in the ammonia dissolution equation moves to the left. The ammonium ion (NH4

+) reacts with OH- to form NH3 and H2O. When the reaction goes from right to left, NH4

+ gives up a hydrogen ion; it acts as a Bronsted-Lowry acid. The hydroxide ion accepts a H+; it acts as a Bronsted-Lowry base. Then we have two acids and two bases.

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

Base acid conjugate conjugate acid base

A conjugate acid is the particle formed when a base gains a hydrogen ion. A conjugate base is the particle that remains when an acid has donated a hydrogen ion. A conjugate acid-base pair is two substances that are related by the loss or gain of a single hydrogen ion.

Water can either donate or accept hydrogen ion. A substance that can act both is called amphoteric.

4. Common Acid & Base

Table 4. Common Acids and Bases

Name Formula Locations

Acids

Acetic acid HC2H3O2 Vinegar(aqueous solution)

Acetylsalicylic acid HC9H7O4 Aspirin

Ascorbic acid H2C6H6O6

Vitamin C

Citric acid H3C6H5O Lemon juice, citrus fruits

7

Hydrochloric acid HCl Gastric juices (digestive fluid in stomach)

Sulfuric acid H2SO4 Batteries

Bases

Ammonia NH3 Household cleaners(aqueous solution)

Calcium hydroxide Ca(OH)2 Slaked lime (used in mortar for construction)

Magnesium hydroxide Mg(OH)2 Milk of magnesia(antacid and laxative)

Potassium hydroxide(also called caustic potash)

KOH Soft soap

Sodium hydroxide NaOH Drain and oven cleaners

5. Strength of Acids and BaseAcids are classified as strong or weak depending on the degree to

which they ionize in water.

Strong acids are completely ionized in aqueous solutions Examples are hydrochloric acid and sulfuric acid. Weak acids ionize only slightly in aqueous solution.

An acid-base equilibrium always lies in the direction of the weaker acid.

Acid Base

Thus, the strongest acids have the weakest conjugate bases, and the strongest bases have the weakest conjugate acids.

6. Ionization of H2OWater shows a very small conduction. That was the result from self-

ionization or autoionization, a reaction in which two like molecules react to give ions.

In case of water, a proton from one H2O molecule is transferred to another H2O molecule, leaving behind an OH- ion and forming a hydronium ion, H3O+ (aq):

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

The self-ionization is: H2O (l) → H+ (aq) + OH- (aq)

We can see the slight extent to which self-ionization occurs by the small value of the equilibrium constant, Kc. An equilibrium constant is the constant value of the equilibrium-constant value. This expression is obtained by multiplying the concentrations of the product and dividing by the concentrations of the reactants, raising each concentration term to a power equal to the coefficient of that substance in the chemical equation.

Kc = [H + ][OH - ] [H20]

The value of Kc at room temperature is 1.8 x 10 -16. In water, its concentration remains 56 M at 25º C. Replacing [H2O] by Kc, then the ionic product [H+] [OH-] equals a constant:

[H2O]K = [H+][OH-], and the value of the constant is called the ion-product constant for water (Kw). At 25º C, the value of Kw is 1.00 x 10-14.

Acidic solution: [H+] > 1.00 x 10-7 MNeutral solution: [H+] = 1.00 x 10-7 MBasic solution: [H+] < 1.00 x 10-7 M

7. pH (acid) and pOH (base) concepts

In 1909, Danish scientist Soren Sorensen (1868-1939) proposed the pH scale. On the pH scale, neutral solutions have a pH of 7.0. A pH is strongly acidic. The pH of a solution is the negative logarithm of the hydrogen-ion concentration.

pH = -log[H+]

In a neutral solution [H+] = 1.0 x 10-7 mol/L. The pH of a neutral is 7.

pH = -log (1.0 x 10-7 mol/L)pH = -(log 1 + log 10-7)pH = -[0.0 + (-7)] = 7.0

Meanwhile, the pOH of a solution equals the negative logarithm of the hydroxide-ion concentration.

pOH = -log[OH-]

A neutral solution has a pOH of 7.0. A solution with a pOH less than 7.0 is basic, and higher than 7.0 is acidic.

pH + pOH = 14pH = 14 – pOHpOH = 14- pH

Neutral solution: pH = 7.0; [H+] = 1 x 10-7 m/LAcidicl solution: pH < 7.0; [H+] > 1 x 10-7 m/LBasic solution: pH > 7.0; [H+] < 1 x 10-7 m/L

8. IndicatorsAn indicator (In) is a weak acid or base that undergoes dissociation

in a known pH range. In this range, the acid or base is a different color from its conjugate bases or acid.

Acidneutral

Base 

[H+](mol/L)

100 10-1 10-2 10-3 10-4 10-5 10-6 10-7 10-8 10-9 10-10 10-11 10-12 10-13 10-14  

pH 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH range

Universal indicator

red red orange-red

orange pale orange

orange-yellow

pale yellow

green-yellow

green dark-green

blue blue blue blue blue  

cyanidin(red

cabbage water)

red red red cerise purple blue blue blue aqua-marine

emerald-green

lime lime yellow yellow yellow  

blue litmus indicator

red red red red red red red blue blue blue blue blue blue blue blue 5.0 - 8.0

red litmus red red red red red red red red blue blue blue blue blue blue blue 5.0 - 8.0

indicator

phenol-phthalein indicator

colour-less

colour-less

colour-less

colour-less

colour-less

colour-less

colour-less

colour-less pink pink pink pink pink pink pink

8.3 - 10.0

thymol blue indicator

yellow yellow yellow yellow yellow yellow yellow yellow yellow blue blue blue blue blue blue 8.0 - 9.6

phenol red indicator

yellow yellow yellow yellow yellow yellow yellow yellow red red red red red red red 6.8 - 8.4

bromo-thymol blue

indicatoryellow yellow yellow yellow yellow yellow yellow blue blue blue blue blue blue blue blue 6.2 - 7.6

methyl red indicator

pink pink pink pink pink pink yellow yellow yellow yellow yellow yellow yellow yellow yellow 4.4 - 6.0

bromo-cresol green

indicator

yellow yellow yellow yellow yellowpale blue-green

blue-green

blue-green

blue-green

blue-green

blue-green

blue-green

blue-green

blue-green

blue-green

3.8 - 5.4

methyl orange

indicatorred red red red yellow yellow yellow yellow yellow yellow yellow yellow yellow yellow yellow 3.1 - 4.4

bromo-phenol blue

yellow yellow yellow yellow blue blue blue blue blue blue blue blue blue blue blue 3.0 - 4.6

cresol red red red red yellow yellow yellow yellow yellow yellow yellow yellow yellow yellow yellow yellow 0.2 - 1.8

pH 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH

9. BuffersBuffer is a solution characterized by the ability to withstand changes

in pH when limited amounts of acid or base are added to it.

Most buffers contain a weak acid and its conjugate base, or a weak base and its conjugate acid.

For an example, biological fluid such as blood, are usually buffer solutions because the control of pH is vital to their proper functioning.

Say, a buffer contains approximately equal molar amounts of a weak acid HA and its conjugate base, A-. If a strong acid is added to the buffer, it supplies hydrogen ions that react with the base A-.

H+ (aq) + A- (aq) → HA (aq)

On the other hand, if a strong base is added to the buffer, it supplies hydroxide ions. These ions react with the acid HA:

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

Thus, a buffer solution resists changes in pH by the ability to combine with both H+ and OH- ions.

Two important characters of a buffer are the pH and the buffer capacity, which is the amount of acid or base the buffer can react with

before given a significant pH change. Buffer capacity depends on the amount of acid and conjugate base in the solution.

10. NeutralizationNeutralization is a reaction in which an acid and a base react in an

aqueous solution to produce a salt and water.

Examples:HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l)H2SO4 (aq) + 2KOH (aq) → K2SO4 (aq) + 2H2O (l)

1.1.1.TitrationThe amount of acid or base in a solution can be determined

by carrying out a neutralization reaction. An appropriate acid-base indicator is used to show when neutralization has occurred. Phenolpthalein is often the indicator for acid-base neutralization reactions. Solutions that contain this indicator turn from colorless to deep pink as the pH of the solution is changed from acidic to basic. Neutral solutions are very faintly pink.

Steps in Titration1. A measured volume of the unknown concentration is added

to the flask.2. The indicator is added to the solution.3. Measured volumes of a base known concentration are

mixed into the acid until the indicator barely changes color.

Standard solution is the solution of known concentration. The point at which the indicator changes color is the end point of the titration.

Thus, titration is the addition of a known amount of solution of known concentration to determine the concentration of another solution.

1.1.2.EquivalentsOne equivalent is the amount of an acid (or base) that will

give one mole of hydrogen (or hydroxide ions).

1.1.3.NormalityAn older unit sometimes used to express the equivalents of an

acid and base is normality. Normality of a solution is the concentration expressed as the number of equivalents of solute in 1 L solution.

A solution containing 1.0 equiv of an acid or base per liter has a normality of 1.0

Normality (N) = equiv/L

The numerical values of normality and molarity are equal for acids and bases that give 1 equiv of H+ or OH- per mole.

The number of equivalents of an acid or base in a known volume of a solution of known normality can be calculated.

Number of equivalents of solution = volume (L) of solution x normality [Equiv = V (L) x N]

Solutions of known morality can be made less concentrated by diluting them with water. The formula to calculate the changes in concentration is:

N1V1=N2V2

Titration calculations can be done in terms of morality instead of molarity since normality allows number of ionizable hydrogens in an acid whereas molarity does not. In a titration, the point of neutralization is called the equivalence point. In conclusion, it is possible to calculate the number of equivalents of acid or base in an unknown sample.

Equivalents of acid = NAVA

Equivalents of base = NBVB

1.1.4.Salt HydrolysisThe cations or anions of the dissociated salt accept hydrogen

ions from water or donate hydrogen ions to water. Hydrolyzing salts are usually derived from a strong acid and a weak base or from a weak acid and a strong base.

1.1.5.Solubility Product ConstantMost “insoluble” salts dissolve to some extent in water. These

salts are said to be slightly or sparingly soluble in water.

Solubility product constant, Ksp, is equal to the product of the concentration terms each raised to the power of the coefficient of the substance in the dissociation equation.

ActetatesAgC2H3O2 -- 2 x 10-3

BromidesAgBr -- 5 x 10-13

PbBr2 -- 5 x 10-6

CarbonatesBaCO3 -- 2 x 10-9

CaCO3 -- 5 x 10-9

MgCO3 -- 2 x 10-8

ChloridesAgCl -- 1.6 x 10-10

Hg2Cl2 -- 1 x 10-18

PbCl2 -- 1.7 x 10-5

ChromatesAg2CrO4 -- 2 x 10-12

BaCrO4 -- 2 x 10-10

PbCrO4 -- 1 x 10-16

SrCrO4 -- 4 x 10-5

FluoridesBaF2 -- 2 x 10-6

CaF2 -- 2 x 10-10

PbF2 -- 4 x 10-8

HydroxidesAl(OH)3 -- 5 x 10-33

Cr(OH)3 -- 4 x 10-38

Fe(OH)2 -- 1 x 10-15

Fe(OH)3 -- 5 x 10-38

Mg(OH)2 -- 1 x 10-11

Zn(OH)2 -- 5 x 10-17

IodidesAgI -- 1 x 10-16

PbI2 -- 1 x 10-8

SulfatesBaSO4 -- 1.4 x 10-9

CaSO4 -- 3 x 10-5

PbSO4 -- 1 x 10-8

SulfidesAg2S -- 1 x 10-49

CdS -- 1 x 10-26

CoS -- 1 x 10-20

CuS -- 1 x 10-35

FeS -- 1 x 10-17

HgS -- 1 x 10-52

MnS -- 1 x 10-15

NiS -- 1 x 10-19

PbS -- 1 x 10-27

ZnS -- 1 x 10-20

1.1.6.Common Ion EffectIn a saturated solution of silver chloride equilibrium is

established between the solid silver chloride and its ions.

AgCl (s) → Ag+ (aq) + Cl-(aq) Ksp = 1.8 x 10-10

If you would add silver nitrate, the product of the [Ag+] and [Cl-] would be greater than Ksp. Applying Le Chatelier’s principle, the stress of the additional Ag+ can be relieved if the reaction shifts to the left. In our example, silver ion is the common ion.

Common ion is an ion that is common to both salts. The lowering of the solution of a substance by the addition of a common ion is called common ion effect.

11. SaltsSalts are often prepared by neutralization of the appropriate acid

and base.

Properties of SaltsSalts are characterized by ionic bonds, relatively high melting

points, electrical conductivity when melted or when in solution, and a crystalline structure in solid state.

Table 11.1 Common salts and their applicationsNAME FORMULA APPLICATIONSAmmonium sulfate (NH4)SO4 FertilizerBarium sulfate BaSo4 Gastrointestinal

studies; white pigment

Calcium chloride CaCl2 Deicing roadways and sidewalks

Calcium sulfate dihydrate (gypsum)

CaSO4∙2H2O Plasterboard

Copper sulfate pentahydrate (blue vitriol)

CuSO4∙5H2O Dyeing, fungicide

Calcium sulfate sesquihydrate

CaSO4∙1/2H2O Plaster casts

Potassium chloride KCl Sodium-free salt substitute

Potassium permanganate

KMnO4 Disinfectant and fungicide

Silver nitrate AgNO3 Cauterizing agentSilver bromide AgBr Photographic

emulsionsSodium hydrogen carbonate (baking soda)

NaHCO3 Antacid

Sodium carbonate decahydrate (washing soda)

Na2CO3∙10H2O Glass manufacture; water softener

Sodium chloride (table salt)

NaCl Body electrolyte; chlorine manufacture

Sodium thiosulfate (hypo)

Na2S2O3 Fixing agent in photographic process