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Acids , Bases and Salts
Concept
(i) Different definitions of acids (ii) Classification of acids
Introduction
For thousands of years people have known that vinegar, lemon juice and many other
foods taste sour. However, it was not until a few hundred years ago that it was discovered
why these things taste sour - because they are all acids. The term acid, in fact, comes
from the Latin term acere, which means "sour".Bases can be thought of as the chemicals
opposite of acids. A reaction between an acid and base is called neutralization. Bases and
acids are seen as opposites because the effect of an acid is reduced or cancelled by a
base. Bases and acids are typically found in aqueous solution. While there are many
slightly different definitions of acids and bases, in this lesson we will introduce the
fundamentals of acid / base chemistry.
Acids
Commonly an acid can be recognized by its following actions :
1) Corrosive ('burns' your skin)
2) Sour taste (e.g. lemons, vinegar)
3) Contains hydrogen ions (H+) when dissolved in water
4) Its solution has pH less than 7
5) Turns blue litmus paper to a red colour
6) Reacts with bases to form salt and water
7) Reacts with metals to form hydrogen gas
8) Reacts with carbonates to form carbon dioxide, water and a salt
Common examples of acids include acetic acid (in vinegar), sulfuric acid (used in car
batteries), and tartaric acid (used in baking). As these three examples show, acids can be
solutions, liquids, or solids, gases such as hydrogen chloride can be acids as well.
Chemicals or substances having the property of an acid are said to be acidic.
Some acids and their natural occurrence are given below.
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Natural source
1) Sting of bees and ants
2) Tamrind, Grapes
3) Gastric juice
4) Apples
5) Fats
6) Sour milk
7) Urine
8) Rancid butter
9) Citrus fruits ( Lemon )
Acid present
Formic acid
Tartaric acid
Hydrochloric acid
Malic acid
Stearic acid
Lactic acid
Uric acid
Butyric acid
Citric acid
Activity 1 - Take 5 ml of lemon juice, tamrind water, vinegar and hydrochloric acid
each in a different test tube. Put one or two drops of each of these solutions on following
indicators as shown in the table. Observe the change in colour on red litmus, blue
litmus, methyl orange and phenolphthalein.
-----------------------------------------------------------------------------------------------------------Sample solution
Red litmus Blue litmus Methyl orange
Phenolphthalein
paper
paper
solution
solution
-----------------------------------------------------------------------------------------------------------1) Lemon juice
2) Tamrind water
3) Vinegar
4) Hydrochloric acid
There are three common definitions for acids: the Arrhenius definition, the Brønsted Lowry definition and the Lewis definition.
The Arrhenius definition states that acids are substances which increase the
concentration of hydronium ions or hydrogen ions (H3O+ or H+ ) in aqueous solution.
The Brønsted-Lowry definition is an expansion of Arrhenius definition.
According to Bronsted - Lowry definition, an acid is a substance
which can act as a proton donor.
Most acids , encountered in everyday life , are aqueous solutions or can be dissolved in
water and then these two definitions are most relevant.
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By the Brønsted-Lowry definition, any compound which can easily be
deprotonated can be considered an acid. Examples include alcohols and amines which
contain O-H or N-H fragments.
According to Lewis definition, acid is a substance which can accept a
lone pair of electrons..
In chemistry, the Lewis definition of acidity is frequently used. Lewis acids are
electron-pair acceptors. Examples of Lewis acids include all metal cations like H+, Fe3+
and electron-deficient molecules such as boron trifluoride and aluminium trichloride.
Hydronium ions are acids according to all three definitions. Interestingly, alcohols and
amines mentioned above as examples of Brønsted - Lowry acids can function as Lewis
bases at the same time.
Arrhenius concept
The Swedish chemist Svante Arrhenius in 1884 , attributed the properties of
acidity to hydrogen ion ( today , more precisely H3O+ ion) . An Arrhenius acid is a
substance that increases the concentration of the hydronium ion, H3O+, when dissolved in
water. For example, HCl, HNO3 increase the concentration of H3O+ in water. This
definition stems from the equilibrium dissociation of water into hydronium (H3O+) and
hydroxide (OH−) ions.
→
H3O+ (aq) + OH−(aq)
HCl (aq) + H2O (l) →
H3O+ (aq) + Cl – (aq)
H2O(l) + H2O(l)
In pure water the majority of molecules exist as H2O but a small number of
molecules are constantly dissociating and re-associating. Pure water is neutral with
respect to acidity or basicity because the concentration of hydroxide ions is always equal
to the concentration of hydronium ions. Note that chemists often write H+(aq) and refer
to the hydrogen ion when describing acid-base reactions but the free hydrogen nucleus, a
proton, does not exist alone in water, it exists as the hydronium ion, H3O+.
Brønsted – Lowry concept
While the Arrhenius concept is useful for describing many reactions, it is also
quite limited in its scope. It requires aqueous medium and describes acid in terms of H+ .
Hence reactions not taking place in aqueous medium are not treated as acid – base
reactions. Similarly, substances not containing H+ ions are not recognized as acids .
In 1923 chemists Johannes Nicolaus Brønsted and Thomas Martin Lowry
independently recognized that acid-base reactions involve the transfer of a proton. A
Brønsted-Lowry acid (or simply Brønsted acid) is a species that donates a proton to a
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Brønsted-Lowry base. For example, a Bronsted acid HCl donates a proton to a Bronsted
base NH3.
Lewis concept
A third concept was proposed by Gilbert N. Lewis which includes reactions
with acid-base characteristics that do not involve a proton transfer. A Lewis acid is a
species that accepts a pair of electrons from another species; in other words, it is an
electron pair acceptor. Brønsted acid-base reactions are proton transfer reactions while
Lewis acid-base reactions are electron pair transfers. All Brønsted acids are also Lewis
acids but not all Lewis acids are Brønsted acids. Contrast the following reactions which
could be described in terms of acid-base chemistry.
In the first reaction a fluoride ion, F−, gives up an electron pair to boron
trifluoride to form the product tetrafluoroborate. Fluoride "loses" a pair of valence
electrons because the electrons shared in the B—F bond are located in the region of space
between the two atomic nuclei and are therefore more distant from the fluoride nucleus
than they are in the lone fluoride ion. BF3 is a Lewis acid because it accepts the electron
pair from fluoride. This reaction cannot be described in terms of Brønsted theory because
there is no proton transfer.
The second reaction can be described using either theory. A proton is
transferred from an unspecified Brønsted acid to ammonia, a Brønsted base.
Alternatively, ammonia acts as a Lewis base and transfers a lone pair of electrons to form
a bond with a hydrogen ion.
The species that gains the electron pair is the Lewis acid; for example, the
oxygen atom in H3O+ gains a pair of electrons when one of the H—O bonds is broken
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and the electrons shared in the bond become localized on oxygen. Depending on the
context, a Lewis acid may also be described as an oxidizer or an electrophile.
The Brønsted-Lowry definition is the most widely used definition; unless otherwise
specified. Acid-base reactions are assumed to involve the transfer of a proton (H+) from
an acid to a base.
Activity 2 – You are give the samples of HCN, AlCl3, CH3COOH, H2O, H3PO4 and
BeF2 . Classify them as acids according to Arrhenius concept, Bronsted – Lowry concept
and Lewis concept.
According to Arrhenius concept –
According to Bronsted – Lowery concept –
According to Lewis concept –
Classification of acids
Inorganic acids are classified into two classes – (i) Hydracids (ii) Oxyacids
(i) Hydracids - These are binary compounds containing hydrogen and a non-metallic
element. For example, HF, HCl, HBr etc.
(ii) Oxyacids - These acids contain oxygen in addition to hydrogen and a non-metal. For
example, HNO3, H2SO4, H3PO4, HClO4 etc.
Activity 3 - Suggest the acids formed by the elements Br, N and Si and classify them
into hydracids and oxyacids.
Test your understanding
1) Give the names of two substances which act as acid according to Bronsted – Lowry
concept as well as Arrhenius concept.
2) Give the names of two substances which act as acid according to Bronsted – Lowry
concept as well as Lewis concept.
3) Name the acids which are found in (i) sting of ant (ii) orange fruit (iii) sour milk
4) Why should curd and sour substances not be kept in brass and copper vessels ?
5) What is the difference between a hydracid and an oxyacid ?
6) Out of the various concepts of acids , which concept is most widely used ?
7) Name the acid used in cooking of food.
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Concept
(i) Strength of acids
(ii) pH scale
Strength of acids
The strength of an acid refers to its ability or tendency to lose a proton. A
strong acid is one that completely dissociates in water; in other words, one mole of a
strong acid HA dissolves in water yielding one mole of H+ and one mole of A− and none
of the protonated acid HA. In contrast a weak acid only partially dissociates and at
equilibrium both the undissociated acid and the dissociated species are present in
solution. Examples of strong acids are hydrochloric acid (HCl), hydriodic acid (HI),
hydrobromic acid (HBr), perchloric acid (HClO4), nitric acid (HNO3) and sulfuric acid
(H2SO4). In water each of these essentially ionizes 100%. The stronger an acid is, the
more easily it loses a proton, H+. Two key factors that contribute to the ease of
deprotonation are - the polarity of the H—A bond and the size of atom A which
determines the strength of the H—A bond. If the polarity of the H – A bond is more, then
the tendency to liberate a proton in water is more. If the atom A is small, then the
tendency of the H – A bond to break and to produce H+ ion is more.
For an acid HA in gaseous state, as the size of atom A increases, the strength
of the bond decreases, meaning that it is more easily broken, and the strength of the acid
increases. Bond strength is a measure of how much energy it takes to break a bond. In
other words, it takes less energy to break the bond as atom A grows larger and the proton
is more easily removed by a base. This partially explains why hydrofluoric acid is
considered a weak acid while the other hydrohalic acids (HCl, HBr, HI) are strong acids.
Although fluorine is more electronegative than the other halogens, its atomic radius is
also much smaller, so it shares a stronger bond with hydrogen. Moving down a column in
the periodic table, atoms become less electronegative but also significantly larger and the
size of the atom tends to dominate its acidity when sharing a bond to hydrogen. Hydrogen
sulfide, H2S, is a stronger acid than water, even though oxygen is more electronegative
than sulfur. Just as with the halogens, this is because sulfur is larger than oxygen and the
H—S bond is more easily broken than the H—O bond. As sulfur is larger than oxygen,
H—S bond is more easily broken than the H—O bond; like in case of halogens.
Monoprotic acids
Monoprotic acids are those acids that are able to donate one proton per
molecule during the process of dissociation (sometimes called ionization) as shown
below (symbolized by HA). A monoprotic acid is also called a monobasic acid.
.
HA(aq) + H2O(l) → H3O+ (aq) + A−(aq)
Common examples of monoprotic acids in mineral acids include
hydrochloric acid (HCl) and nitric acid (HNO3). On the other hand, for organic acids the
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term mainly indicates the presence of one carboxyl group and sometimes these acids are
known as monocarboxylic acids. Examples in organic acids include formic acid
(HCOOH), acetic acid (CH3COOH) and benzoic acid (C6H5COOH).
Polyprotic acids
Polyprotic acids, also known as polybasic acids, are able to donate more than
one proton per acid molecule, in contrast to monoprotic acids that only donate one proton
per molecule. Specific types of polyprotic acids have more specific names, such as
diprotic acid (two potential protons to donate) and triprotic acid (three potential protons
to donate). A diprotic acid (here symbolized by H2A) can undergo one or two
dissociations.. Each dissociation has its own dissociation constant, Ka1 and Ka2.
H2A(aq) + H2O(l) → H3O+(aq) + HA−(aq)
HA−(aq) + H2O(l) → H3O+(aq) + A2−(aq)
For example, sulfuric acid (H2SO4) can donate one proton to form the bisulfate anion
(HSO4−), then it can donate a second proton to form the sulfate anion (SO42 - )
pH scale
The concentration of hydrogen ions is commonly expressed in terms of the pH
scale. The pH scale takes its name from the German word ‘potenz’ meaning potential or
strength of hydrogen. . It was introduced by a Danish chemist S.P.L. Sorensen in 1909. It
is a scale used to measure the acidity of a dilute solution. The pH scale uses a range from
0 to 14, with 7.0 indicating neutrality.
Concentration of H+ is usually confined to 10 -1 to 10-14 M range. Thus pH
scale contains values falling between 0 and 14. In some rare cases, you may see pH lower
than 0 or higher than 14, when the concentration of H+ takes some extreme values
Mathematically, pH is defined as a negative decimal logarithm of the hydrogen ion
activity in a solution.
pH = - log10 [ H+]
or
pH = log10 1 / [H+ ]
Where [H+] is the activity of hydrogen ion concentration which in a dilute solution is
equivalent to its molar concentration.
Numbers beginning at 7.0 and moving toward 0 indicate acidity, while the
numbers beginning at 7.0 and moving toward 14 indicate alkalinity, so the scale divides
acids from bases ( or alkalies ).
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Activity 4 - Take 10 ml of 1 M HCl and 1 M CH3COOH in two different test tubes. Is
any of these acids stronger than the other ? Why ?
Activity 5 - You have two solutions A and B. The pH of solution A is 5 and that of
solution B is 9. Which solution has more H ion concentration ? Which of these solutions
is acidic ?
Test your understanding
1) Give the name of (i) one strong monobasic acid
(ii) one weak dibasic acid
2) How is the concentration of hydronium ions, H3O+ affected when a solution of an acid
is diluted ?
3) Give the pH of two substances in the environment or in human body ?
4) Fresh milk has pH 6. How do you think the pH will change as it turns into curd ?
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Concept
Methods of preparation of acids
Methods of preparation of acids
Arrhenius acids can be prepared by following general methods.
1) Direct combination or synthesis
Hydrogen reacts with certain non-metals under proper conditions to form the
corresponding acids.
Sunlight
Boil
H2 + Cl2 ------→
2 HCl
; H2 + S -----→
H2S
2) Action of water on acidic oxides ( non-metallic oxides )
Certain non-metallic oxides dissolve in water to form oxy – acids at room temperature.
CO2 + H2O → H2CO3
;
P2O5 + 3 H2O → 2 H3PO4
3) By displacement of more volatile acid from their salts by a less volatile acid
A more volatile acid can be obtained from its salt by treatment with a less volatile
acid.
+
H2SO4
Less volatile
→
NaHSO4
+
HCl
More volatile
NaNO3 +
Salt
H2SO4
Less volatile
→
NaHSO4
+
HNO3
More volatile
NaCl
Salt
4) By oxidation of non-metals
Strong oxidizing agents like concentrated HNO3 and H2SO4 oxidize non-metals like
carbon, phosphorus and iodine.
S + 6 HNO3
2P + 5 H2SO4
→ H2SO4 + 6 NO2 + 2 H2O
→ 2 H3PO4 + 5 SO2 + 2 H2O
Activity 6 - Take small piece of coke and some sulphur powder separately in shallow
iron dishes. Heat them on a strong flame in air. Coke is converted to CO2 while sulphur is
converted to SO2. Collect the gases CO2 and SO2 in different test tubes and add little
water to them. This forms carbonic acid H2CO3 and sulphurous acid H2SO3 solutions
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respectively. Dip a blue litmus in them. It turns red. Put a drop of methyl orange in each
test tube. It turns the solution pink. This shows that non-metal oxides dissolve in water to
form acids.
Test your understanding
1) Give two methods of preparation of following acids - (i) HNO3 (ii) H2SO4
2) Give the names of any two acidic oxides which when dissolved in water form acid.
3) Give any two examples where oxidation of non-metals form oxyacids .
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Concept .
(i) Reactions of acids
(ii) Uses of acids
Reactions of acids
1) Reaction with bases ( oxides and hydroxides ) - Acids react with bases which are
oxides or hydroxides and form salt and water.
HNO3 + KOH → KNO3 + H2O ; H2SO4 + CuO → CuSO4 + H2O
Acid
Base
Salt
Water
Acid
Base
Salt
Water
2) Reaction with metal carbonates and metal bicarbonates – Acids react with metal
carbonates and metal bicarbonates to form salt and evolve carbon dioxide gas.
CaCO3
+
Metal carbonate
NaHCO3
+
Metal bicarbonate
H2SO4 →
Acid
HCl →
Acid
CaSO4 + CO2 ↑ + H2O
Salt
NaCl
Salt
+ CO2 ↑ + H2O
3) Reaction with active metals – Acids react with active metals to form salt and evolve
hydrogen gas.
H2SO4 + Zn → ZnSO4 + H2↑
; 2 Al + 3 H2SO4 → Al2 (SO4 )3 + 3 H2 ↑
4) Reaction with metallic sulphites and bisulphites - Acids react with metal
sulphites and metal bisulphites to form salt and evolve sulphur dioxide gas.
CaSO3
+
Metal sulphite
NaHSO3
+
Metal sulphite
H2SO4 →
Acid
HCl
Acid
→
CaSO4 + SO2 ↑ + H2O
Salt
NaCl + SO2 ↑ + H2O
Salt
5) Reaction with metallic sulphides – Acids react with metal sulphides to form salt and
evolve hydrogen sulphide gas.
FeS
+ 2 HCl → FeCl2 + H2S ↑ ; ZnS + H2SO4 → ZnSO4 + H2S ↑
Activity 7 - take a small piece of iron sulphide in a test tube. Add 5 ml of dilute HCl to
it. Heat the test tube for two minutes. You get a smell of rotten eggs due to evolution of
H2S gas. Acid reacts with metal sulphides to form salt and evolve hydrogen sulphide gas.
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Activity 8 - Take one gram copper carbonate in a test tube. Add slowly dilute HCl to it
till the solid dissolves. Stir continuously. Carbon dioxide gas is evolved.. Note the colour
of the solution. The solution becomes blue – green as the copper carbonate dissolves.
The blue – green colour is due to the formation of copper (II) chloride. Acid reacts with
metal carbonates to form salt and evolve carbon dioxide gas.
Uses of acids
Acids find following domestic uses :
1) Many of our food stuffs contain acids. E.g leafy vegetables like ambadi ( deccan
hemp ) , ambat chukka ( country sorrel ) contain acids. Amla, cabbage, leaves of
drumstick and coriander give us vitamin C which contains ascorbic acid. Some fruits
contain acids. E.g. orange, sweet lime and lemon contain citric acid. Tamarind gives
tartaric acid while curd gives lactic acid.
2) Sulphuric acid and nitric acid are used for cleaning gold and silver articles.
Hydrochloric acid is used to clean toilets and other enamel ware.
3) Benzoic acid and acetic acid are used in preservatives.
4) Sulphuric acid is used in car batteries.
Acids find following industrial uses :
1) Sulphuric acid and nitric acid are used to manufacture chemical fertilizers.
2) Nitric acid is used in the manufacture of paints and explosives so also in making
perfumes.
3) Hydrochloric acid is used to prepare glucose from starch and to produce gelatin.
4) Carbonic acid is used to prepare aerated drinks.
5) Tartaric acid is used in making baking powder.
Test your understanding
1) A solution reacts with crushed egg – shells to give a gas that turns lime water milky.
The solution contains - a) NaCl
b) HCl
c) LiCl
d) KCl
2) Effervescence is seen when lemon juice is added to baking soda. Explain.
3) Excess of common salt is added while preparing pickles. Why ?
4) Why is it easy to cook food with the help of baking powder ?
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Concept
Different definitions of bases
Bases
The notion of a base as a concept in chemistry was first introduced by the French
chemist Guillaume François Rouelle in 1754. He noted that acids, which in those days
were mostly volatile liquids (like acetic acid), turned into solid salts only when combined
with specific substances. Rouelle considered that such a substance serves as a base for
the salt, giving the salt a "concrete or solid form”. So the substance was named as a base.
Commonly a base can be recognized by its following actions :
1) Slimy or soapy feel on fingers, due to saponification of the lipids in human skin
2) Concentrated or strong bases are caustic on organic matter and react violently with
acidic substances
3) Furnishes hydroxide ions (OH-) when dissolved in water
4) Aqueous solution or molten base dissociates in ions and conduct electricity.
5) Reacts with acids to form salt and water
6) Reactions with indicators : base turns red ‘litmus paper’ blue, ‘phenolphthalein’ pink,
keeps ‘bromthymol blue’ yellow and turns ‘methyl orange’ yellow.
7) Its solution has pH above 7
Common examples of bases include caustic soda ( used in soaps ), caustic
potash ( used in textiles ), magnesium hydroxide (used in milk of magnesia as antacid ),
calcium hydroxide ( used as lime water ), ammonium hydroxide ( used as a cleaning
agent ) , sodium carbonate ( used as soda ash ) and sodium phosphate ( used in
detergents ).
Chemicals or substances having the property of a base are said to be basic or alkaline.
Activity 9 - Take 5 ml solution of washing soda, lime water, ash formed by burning
magnesium wire in air and caustic potash each in different test tube. Put one or two
drops of each of these solutions on following indicators as shown in the table. Observe
the change in colour on red litmus, blue litmus, methyl orange and phenolphthalein
-----------------------------------------------------------------------------------------------------------Sample solution
Red litmus Blue litmus Methyl orange
Phenolphthalein
paper
paper
solution
solution
-----------------------------------------------------------------------------------------------------------1) Washing soda
2) Lime water
3) Ash of MgO
4) Caustic potash
----------------------------------------------------------------------------------------------------------
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There are three common definitions for bases : the Arrhenius definition, the BrønstedLowry definition and the Lewis definition.
The Arrhenius definition states that bases are substances that react with an acid and
produce salt.
Bases which are soluble in water are called alkalies. Alkalies increase the
concentration of hydroxide ions ( OH - ) in aqueous solution.
The Brønsted-Lowry definition is an expansion of Arrhenius definition.
According to Bronsted - Lowry definition, a base is a substance which can act as a
proton acceptor.
Most bases , encountered in everyday life , are aqueous solutions or can be
dissolved in water and then these two definitions are most relevant.
By the Brønsted-Lowry definition, any species which can easily be protonated
can be considered a base. Examples include ammonia, water, all negative ions etc. .
According to Lewis definition, base is a substance which can donate a
lone pair of electrons..
In chemistry, the Lewis definition of basicity is frequently used. Lewis bases
are electron-pair donors. Examples of Lewis bases include all anions like Cl -, CN - and
electron- rich molecules such as water and ammonia.
Hydroxide ions ( OH - )are bases according to all three definitions.
Interestingly, ammonia , water and negative ions mentioned above, as examples of
Brønsted-Lowry bases, can function as Lewis bases at the same time.
Arrhenius Concept
Arrhenius attributed the properties of basicity to hydroxide ions in aqueous
medium. . An Arrhenius base is a substance that increases the concentration of the
hydroxide ion OH -, when dissolved in water. For example, NaOH, KOH increase the
concentration of OH - in water. This definition stems from the equilibrium dissociation of
water into hydronium and hydroxide (OH−) ions.
H2O(l) + H2O(l) → H3O+(aq) + OH−(aq)
NaOH (aq) + H2O (l) → Na+ (aq) + OH– (aq) + H2O (aq)
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Brønsted – Lowry concept
While the Arrhenius concept is useful for describing many reactions, it is also
quite limited in its scope. It requires aqueous medium and describes a base in terms of
OH – ions . Hence reactions not taking place in aqueous medium are not treated as acid –
base reactions. Similarly, substances not furnishing OH - ions in aqueous medium are
not recognized as bases .
A Brønsted-Lowry base (or simply Brønsted base ) is a species that accepts a
proton from a Brønsted-Lowry acid. For example, a Bronsted base Cl – accepts a proton
from a Bronsted acid NH4 +. .
NH4 +
Acid
Lewis concept
+
Cl Base
→
NH4Cl
Salt
A third concept was proposed by Gilbert N. Lewis which includes reactions
with acid-base characteristics that do not involve a proton transfer. A Lewis base is a
species that donates a pair of electrons to another species. In other words, it is an
electron pair donor.. Brønsted acid-base reactions are proton transfer reactions while
Lewis acid-base reactions are electron pair transfers.
Cu 2+
Acid
+
4 NH3
Base
→
[ Cu(NH3)4 ] 2+
Complex ion of salt
Activity 10 – You are given the samples of KOH, NH4OH, PH3, H2O, C2H5NH2 and
H2S. Classify them as bases according to Arrhenius concept, Bronsted – Lowry concept
and Lewis concept.
According to Arrhenius concept –
According to Bronsted – Lowery concept –
According to Lewis concept –
Test your understanding
1) Give the names of two substances which act as a base according to Bronsted – Lowry
concept as well as Arrhenius concept.
2) Give the names of two substances which act as a base according to Bronsted –
Lowry concept as well as Lewis concept.
3) Do basic solutions also have H+ (aq) ions ? If yes, then why are they basic ?
4) Alkalies are normally not kept exposed to air. Why ?
5) Choose from the following substances which act as (i) only bases (ii) bases as well
as alkalies.
(i) NaOH
(ii) CaO (iii) Fe(OH)2
6) Name two bases which are not alkalies.
(iv) NH4OH (v) K2CO3 (vi) ZnO
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Concept
(i) Strength of bases
(ii) pOH scale
Strength of bases
Usually, Arrhenius concept is used to define strength of acids and bases. The
strength of a base refers to its ability or tendency to furnish hydroxide ions in aqueous
medium . A strong base is one that completely dissociates in water. In other words, one
mole of a strong base BOH dissolves in water yielding one mole of OH- ions and one
mole of B+ ions and none of the undissociated BOH . In contrast a weak base only
partially dissociates and at equilibrium both the undissociated base and the dissociated
species are present in solution. Examples of strong bases are sodium hydroxide (NaOH )
and potassium hydroxide (KOH ). In water, each of these essentially ionizes 100%. The
stronger a base is, the more easily it gives a hydroxide ion OH - . Examples of weak
bases include ammonium hydroxide ( NH4OH ) , calcium hydroxide ( Ca(OH)2 ). In
water they dissociate only partially.
Monoacidic base
Monoacidic bases are those bases that are able to furnish one hydroxide ion
per molecule during the process of dissociation (sometimes called ionization) as shown
below (symbolized by BOH).
H2O (l)
BOH (aq) ------→ B+ (aq) + OH – (aq)
Common examples of monoacidic bases include sodium hydroxide ( NaOH )
and potassium hydroxide ( KOH ) . On the other hand, for organic bases this term
mainly indicates the presence of amino groups called amines and nitrogen containing
compounds such as imines or amides. Examples of organic bases include ethyl amine,
diethyl amine etc.
Polyacidic bases
Polyacidic bases are able to furnish more than one hydroxide ion per
molecule of the base e.g. calcium hydroxide Ca(OH)2 is a diacidc base, aluminium
hydroxide Al(OH)3 is a triacidic base .
A diacidic base (here symbolized by B(OH)2 ) can undergo dissociation in
one or two steps . Each dissociation has its own dissociation constant, Kb1 and Kb2.
B(OH)2 (aq) → B2+ (aq) + 2 OH – (aq)
B(OH)3 (aq) → B3+ (aq) + 3 OH – (aq)
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pOH scale
The concentration of hydroxide ions is commonly expressed in terms of the
pOH scale. The pOH scale takes its name from the German word ‘potenz’ meaning
potential or strength of hydroxide ions . It is a scale used to measure the basicity of a
dilute solution. The pOH scale uses a range from 0 to 14, with 7.0 indicating neutrality.
Concentration of OH- is usually confined to 10 -1 to 10-14 M range. Thus pOH scale
contains values falling between 0 and 14. In some rare cases, you may see pOH lower
than 0 or higher than 14, when the concentration of OH- takes some extreme values.
Mathematically, pOH is defined as a negative decimal logarithm of the
hydroxide ion activity in a solution.
pOH = - log10 [ OH-]
or pOH = log10 1 / [OH- ]
Where [OH-] is the activity of hydroxide ion concentration which in a dilute solution is
equivalent to its molar concentration.
Activity 11 – You are given three test tubes. One contains distilled water, one contains an
acidic solution and the remaining one contains an alkaline solution. If you are given only
red litmus paper, how will you identify the contents of each test tube ?
Activity 12 - Take 10 ml of 1 M KOH and 1 M NH4OH in two different test tubes. Is
any of these alkalies stronger than the other ? Why ?
Activity 13 - You have two solutions A and B. The pH of solution A is 2 and that of
solution B is 10. Which solution has more OH - ion concentration ? Which of these
solutions is basic ?
Test your understanding
1) Classify the following into strong and weak bases .
(i) KOH
(ii) NH4OH
(iii) C2H5NH2
(iv) Na2CO3
2) Which one has a higher pOH - an acid or base ? Why ?
3) Give one example each of a monoacidic base and diacidic base .
18
Concept
Methods of preparation of bases
Methods of preparation of bases
Bases can be prepared by following general methods.
1) Direct combination of metals with oxygen
Most of the metals including mercury react with oxygen, when heated, to form oxides
which act as bases.
2 Ca + O2 → 2 CaO
2 Hg + O2 →
;
2 HgO
2) Action of water on soluble metallic oxides
Oxides of metals like sodium, potassium and calcium dissolve in water to form bases
( being soluble in water, they are called alkalies )
K2O + H2O →
2 KOH
CaO + H2O → Ca(OH)2
;
3) Action of water on ammonia
Ammonia dissolves in water to form a weak alkali NH4OH.
NH3 + H2O → NH4OH
4) Action of cold water or steam on metals
Reactive metals like sodium, potassium and calcium react with cold water to form
their hydroxides which behave as bases and alkalies both.
2 Na + 2 H2O → 2 NaOH + H2
:
2 K + 2 H2O → 2 KOH + H2
Slightly less reactive metals like zinc, magnesium and aluminium react with hot water
or steam to form their oxides which behave as bases.
Zn + H2O → ZnO + H2
;
4 Al + 3 H2O → 2 Al2O3 + 3 H2
19
5) Precipitation of metal hydroxides
Many hydroxides ( exception – NaOH, KOH and NH4OH ) are precipitated from
their salt solutions when treated with sodium hydroxide.
CuSO4 + 2 NaOH → Cu (OH)2  + Na2SO4
FeSO4 + 2 NaOH → Fe(OH)2  + Na2SO4
6) Thermal decomposition of carbonates and nitrates
Many metal carbonates like zinc carbonate, barium carbonate decompose on
heating to give metal oxides which behave as bases.
ZnCO3 → ZnO + CO2 ↑ ;
BaCO3 → BaO + CO2 ↑
Similarly, many metal nitrates like copper nitrate, calcium nitrate decompose on
heating to give metal oxides which behave as bases.
2 Cu(NO3)2 → 2 CuO + 4 NO2 ↑+ O2 ↑ ; 2 Ca(NO3)2 → 2 CaO + 4 NO2↑ + O2 ↑
Activity 14 –Burn a magnesium wire and collect the ash of MgO thus formed in a test
tube. Put 10 ml water in it and stir the solution. Take a little calcium carbonate in second
test tube and heat it strongly. It is converted to calcium oxide. Put 10 ml water in it and
stir the solution. Take a small piece of sodium in a third test tube. Put 10 ml water to it.
( There is a chance of small explosion, so add small amount of water at a time carefully.
Do this under the supervision of a teacher ). In all the three test tubes a base i.e.
hydroxide is formed. Dip a red litmus paper in each test tube. It turns blue. Put a drop of
methyl orange in each test tube. It turns the solution yellow. This shows that direct
combination of elements ( Mg + O ), decomposition of metal carbonate ( CaCO3) and
reaction of a metal with cold water ( Na + H2O) produce an alkali / base .
Test your understanding
1) Starting from a metal, give two methods of preparation of bases.
2) ‘One base can be prepared by using another base’. Justify this statement by giving an
example.
3) Give two examples to show that bases can be obtained by decomposition of salts.
20
Concept .
(i) Reactions of bases ( and alkalies )
(ii) Uses of bases (and alkalies )
Reactions of bases ( and alkalies )
1) Reaction with acidic oxides - Bases react with acidic oxides like CO2 , SO3 and P2O5
to form corresponding salts.
CaO + CO 2 → CaCO3
;
2 NaOH + SO3 → Na2SO4 + H2O
2) Reaction with acids - Bases react with acids and neutralize them to form salt and
water.
ZnO + H2SO4 → ZnSO4 + H2O ;
KOH + HNO3 → KNO3 + H2O
3) Reaction with ammonium salts - On heating, alkalies react with ammonium salts to
evolve ammonia gas .
Ca(OH)2 + (NH4)2SO4 → CaSO4 + 2 NH3 ↑ + 2 H2O
NaOH + NH4Cl
→ NaCl + NH3 ↑ + H2O
4) Reaction with solution of salt of heavy metals - Alkalies like NaOH, KOH and
NH4OH react with aqueous solution of salt of heavy metals like Cu, Fe, Zn and Pb to
form an insoluble metal hydroxide.
CuCl2 + 2 NaOH → Cu(OH)2  + 2 NaCl ; FeCl2 + 2 NaOH → Fe(OH)2 + 2 NaCl
ZnCl2 + 2 NH4OH → Zn(OH)2 + 2 NH4Cl
Pb(NO3)2 + 2 NaOH → Pb(OH)2 + 2 NaNO3
Activity 15 - Take one gram NH4Cl in a test tube. Add 5 ml dilute NaOH solution to it.
Warm the test tube. Hold a moist turmeric paper near the mouth of the test tube. Soon it
becomes reddish brown showing the evolution of ammonia gas. Alkalies react with
ammonium salts to evolve ammonia gas on heating.
Activity 16 - Take 5 ml dilute solution of ferrous sulphate in water. Add drop by drop
ammonium hydroxide solution to it. At one stage, you see a dirty green precipitate of
Fe(OH)2 . Alkali like NH4OH reacts with aqueous solution of salt of heavy metals like
Fe to form an insoluble metal hydroxide.
21
Uses of bases ( and alkalies )
1) Sodium hydroxide is used in the manufacture of soap and paper.
2) Oxides of chromium, manganese, cobalt , iron and copper are used in the manufacture
of paints.
3) Potassium hydroxide is used in alkaline batteries.
4) Magnesium hydroxide is used in antacids to neutralize the acidity in the stomach.
5) Aluminium hydroxide is used as a foaming agent in fire extinguishers.
6) Calcium hydroxide is used in the manufacture of bleaching powder, in neutralizing
acidity of soil and for softening water.
Test your understanding
1) 10 ml of a solution of NaOH is found to be neutralized by 16 ml of a given solution of
HCl. If we take 40 ml of the same solution of HCl , the amount of NaOH solution
required to neutralize it will be :
a) 12 .5 ml
b) 8 ml
c) 25 ml
d) 16 ml
2) Alkalies should not be left exposed to air. Why ?
3) Which of the following salts will form a precipitate with NaOH solution ?
a) NH4Cl
b) CuCl2
c) NaCl
d) Al(OH)3
4) Name the alkalies used in the –
a) manufacture of soap
b) preparing bleaching powder
c) antacid
22
Concept
(i) Acid - base indicators
(ii) Buffer solutions
Acid – Base Indicators
Acid - Base indicators (also known as pH indicators) are substances which
change colour with change of pH . An acid-base indicator is a dye which is a weak acid
or a weak base. The undissociated form ( also called acid form ) of the indicator has a
different color than the ionic form ( also called the base form ) of the indicator. An
Indicator shows color change over a range of hydrogen ion concentration . This range is
termed the color change interval or transition range of the indicator. . It is expressed as a
pH range. An indicator shows one colour in acidic medium and a different colour in
alkaline medium.
Litmus is a weak acid. It is a common acid-base indicator. The un-ionized litmus is red,
whereas the ion is blue in colour.
Litmus solution is a purple dye which is extracted from lichen, a plant belonging to the
division Thallophyta and is commonly used as an indicator. When the litmus solution is
neither acidic nor basic, its colour is purple. When it is acidic, it is blue and when it is
alkaline it is red. There are many other natural materials like red cabbage leaves,
turmeric, coloured petals of some flowers such as Hydrangea, Petunia and Geranium
which indicate the presence of acid or base in a solution. These are called acid – base
indicators or sometimes simply indicators.
Universal indicator which is actually a mixture of several indicators displays
a variety of colours over a wide range of pH, so it can be used to determine an
approximate pH of a solution but is not used for titrations .
Acid-base indicators are commonly employed to mark the end of an acid-base
titration or to measure the existing pH of a solution. . Litmus is not used in titrations
because the pH range over which it changes colour is too great . There is no visual way to
see the completion of the acid – base reaction. So an indicator which shows a sharp
colour change over a small pH range is used to find the end of acid - base reaction also
called neutralization reaction.
Some common acid – base indicators, their colour changes and pH ranges are
given below.
23
Indicator
Low pH color
Transition pH
range
High pH
color
Gentian violet (Methyl violet 10B) yellow
0.0–2.0
blue-violet
Leucomalachite green (first
transition)
yellow
0.0–2.0
green
Leucomalachite green (second
transition)
green
11.6–14
colorless
Thymol blue (first transition)
red
1.2–2.8
yellow
Thymol blue (second transition)
yellow
8.0–9.6
blue
Methyl yellow
red
2.9–4.0
yellow
Bromophenol blue
yellow
3.0–4.6
purple
Congo red
blue-violet
3.0–5.0
red
Methyl orange
red
3.1–4.4
orange
Bromocresol green
yellow
3.8–5.4
blue
Methyl red
red
4.4–6.2
yellow
Methyl red
red
4.5–5.2
green
Azolitmin
red
4.5–8.3
blue
Bromocresol purple
yellow
5.2–6.8
purple
Bromothymol blue
yellow
6.0–7.6
blue
Phenol red
yellow
6.8–8.4
red
Neutral red
red
6.8–8.0
yellow
Naphtholphthalein
colorless to
reddish
7.3–8.7
greenish to
blue
Cresol Red
yellow
7.2–8.8
reddishpurple
Phenolphthalein
colorless
8.3–10.0
fuchsia
Thymolphthalein
colorless
9.3–10.5
blue
Alizarine Yellow R
yellow
10.2–12.0
red
Litmus
red
4.5-8.3
blue
Activity 17 - Take a beet root and crush in a mortar. Add sufficient water to obtain the
extract. Filter the extract. Collect the filtrate to test the substances you may have tested
earlier. Arrange four test tubes in a test tube stand. Label them as A, B, C and D. Pour 2
ml. each of lemon juice solution, soda water, vinegar and baking soda solution in them
respectively. Put 2 to 3 drops of the beet-root extract in each test tube and note the colour
change ,if any. Write your observation in a table.
24
Buffer solution
A buffer solution is an aqueous solution consisting of a mixture of a weak
acid and its conjugate base ( or its salt with a strong base ) or a weak base and its
conjugate acid ( or its salt with a strong acid ). The former is called an acidic buffer
which has pH less than 7 while the latter is called a basic buffer which has pH more than
7. For example, CH3COONa and CH3COOH form an acidic buffer while NH4Cl and
NH4OH form a basic buffer. Buffer has the property that the pH of the buffer solution
changes very little when a small amount of strong acid or base is added to it. Buffer
solutions are used as a means of keeping pH at a nearly constant value in a wide variety
of chemical applications.
In a human body, many reactions take place within a narrow range of pH .
The digestive enzymes in the stomach require an acidic medium to digest the proteins. In
the intestine, the digestive enzymes need an alkaline medium. The body produces
chemicals called buffers that neutralize any changes in the acidity or alkalinity and keep
the pH of the blood constant.
Two natural buffers – carbonate and bicarbonate – play an important role in regulating
pH levels of human body. When blood becomes too acidic, the body produces
bicarbonate to balance out the acidity. When the blood becomes too alkaline, the kidneys
introduce carbonic acid ( or carbonate ) into the blood to bring down the excess
alkalinity.
The phosphate buffering system is used between cells in the urine to take acid out of the
body. Phosphate literally holds onto the acidic hydrogen ions in the urine to ensure that
they are eliminated and not reabsorbed back into the body.
Activity 18 - Take 50 ml. o.1 M acetic acid and 50 ml 0.1 M sodium acetate solution and
mix them in a beaker. Find out the pH of this solution with the help of pH meter or pH
paper. Add 2 to 3 drops of 0.1 M concentrated HCl ( a strong acid ) to it. Find out the pH
of this solution. Is there any change in the pH ? Report your observation to your teacher.
Test your understanding
1) What is the use of an indicator ? How does an indicator function ?
2) What is the chemical composition of an acid – base indicator ?
3) What is the use of (i) universal indicator (ii) buffer
4) Illustrate the use of acid – base indicator in a titration ?
5) Give any two applications of buffers.
25
Concept
(i) Acid – Base Reactions – Different concepts
(ii) Practical applications of neutralization
Acid – Base reactions
Acid-base reaction is a chemical reaction that occurs between an acid and a base .
(i) Arrhenius concept
As defined by Arrhenius, acid-base reactions are characterized by Arrhenius acids
which dissociate in aqueous solution to form hydrogen ions (H+) and Arrhenius bases
which form hydroxide (OH−) ions.
The universal idea of aqueous acid-base reaction of the Arrhenius concept is
described as the formation of water from hydrogen and hydroxide ions.
H+ (aq) + OH− (aq) → H2O ( l )
More recent IUPAC recommendations now suggest the newer term
"hydronium" be used in favor of the older accepted term "oxonium” in place of
hydrogen ion ( H+) .
( In modern times, the use of H+ is regarded as a shorthand notation for H3O+. In reality,
it is known that a bare proton H+ does not exist as a free species in solution.) .
H3O+ (aq) + OH− (aq) → 2H 2O ( l )
This also leads to an expectation that according to Arrhenius concept, the acid –
base reaction always leads to formation of salt and water. This is also called a
neutralization reaction.
acid+ + base− → salt + water
The positive ion from a base forms a salt with the negative ion from an acid.
For example, two moles of the base sodium hydroxide (NaOH) can combine with one
mole of sulfuric acid (H2SO4) to form two moles of water and one mole of salt - sodium
sulfate.
2 NaOH + H2SO4 → 2 H2O + Na2SO4
26
The Arrhenius definitions of acidity and alkalinity are restricted to aqueous
solutions and refer to the increase in concentration of the solvent ions i.e. H+ and OH –
ions. . Under this definition, pure H2SO4 or HCl dissolved in toluene are not acidic, and
molten KOH and solutions of sodium amide in liquid ammonia are not alkaline.
(ii) Brønsted-Lowry concept
The Brønsted-Lowry definition is based upon the idea of protonation of bases
and the de-protonation of acids — that is, the ability of acids to "donate" hydrogen ions
(H+) or protons to bases which "accept" them. Unlike the previous definitions, the
Brønsted-Lowry definition does not refer to the formation of salt and solvent but instead
to the formation of conjugate acids and conjugate bases, produced by the transfer of a
proton from the acid to the base. In this approach, acids and bases are fundamentally
different in behavior from salts which are seen as electrolytes. An acid and a base react
not to produce a salt and a solvent but to form a new acid and a new base. The concept
of neutralization is thus absent.
According to Brønsted-Lowry definition, an acid is a compound or species
that can donate a proton and a base is a compound or species that can receive a proton.
An acid-base reaction is, thus, the removal of a hydrogen ion from the acid and its
addition to the base. This does not refer to the removal of a proton from the nucleus of an
atom which would require levels of energy not attainable through the simple dissociation
of acids but to removal of a hydrogen ion (H+).
According to this theory, in an acid – base reaction, a new acid-base pair is
formed which is called ‘conjugate acid- base pair’.
HCl
Acid 1
+
NH3 →
Base 2
NH4+
+
Conjugate acid 2
Cl –
Conjugate base 1
For example, when HCl ( acid ) and NH3 ( base ) react, they form a conjugate
acid ( NH4+) and conjugate base ( Cl - ) pair. According to Bronsted – Lowry concept,
the acid-base pair which differs in a proton is called a conjugate acid- base pair.
When HCl and NH3 form one acid – base pair, then they produce Cl – as the conjugate
base of the acid HCl and NH4+ as the conjugate acid of the base NH3.
It can be explained by taking another example. The removal of a proton
(hydrogen ion) from an acid produces its conjugate base (which is the acid with a
hydrogen ion removed) and the reception of a proton by a base produces its conjugate
acid ( which is the base with a hydrogen ion added.) . For example, the removal of H+
from hydrochloric acid (HCl) produces the chloride ion (Cl−), as the conjugate base of the
acid HCl.
HCl
→
H+
+
Cl−
Acid
Conjugate base
27
The addition of H+ to the hydroxide ion (OH−), a base, produces water (H2O), as
its conjugate acid.
H+
+
OH−
→
H2O
Base
Conjugate acid
Brønsted-Lowry acid-base theory has several advantages over Arrhenius theory.
Consider the following reactions of acetic acid (CH3COOH) .
Both theories easily describe the first reaction : CH3COOH acts as an
Arrhenius acid because it acts as a source of H3O+ when dissolved in water and it acts as
a Brønsted acid by donating a proton to water. In the second example CH3COOH
undergoes the same transformation, in this case donating a proton to ammonia (NH3) but
cannot be described using the Arrhenius definition of an acid because the reaction does
not produce hydronium ion. . Brønsted-Lowry theory can also be used to describe
molecular compounds like NH3 whereas Arrhenius acids must be ionic compounds.
Hydrogen chloride (HCl) and ammonia combine under several different
conditions to form ammonium chloride, NH4Cl and exists as hydronium and chloride
ions. Following reactions illustrate the limitations of Arrhenius's definition i.e. they can
not be explained by Arrhenius concept.
1. H3O+(aq) + Cl−(aq) + NH3 (aq) →
Cl−(aq) + NH4+(aq)
2. HCl(benzene) + NH3(benzene) → NH4Cl(s)
3. HCl (g) + NH3(g) → NH4Cl(s)
As with the acetic acid reactions, both definitions work for the first
example, where water is the solvent and hydronium ion is formed on the left hand side.
The next two reactions do not involve the formation of ions but are still proton transfer
reactions. In the second reaction hydrogen chloride and ammonia (dissolved in benzene)
react to form solid ammonium chloride in a benzene solvent and in the third gaseous HCl
and NH3 combine to form the solid NH4Cl .
28
The reaction of ammonia, a base, with acetic acid in absence of water can be
described to give ammonium cation, an acid, and acetate anion, a base .
CH3COOH + NH3 → NH 4 + + CH3COO−
Acid 1
Base2
Acid 2
Base 1
Bronsted – Lowry definition also explains the dissociation of water into low
concentrations of hydronium and hydroxide ions.
H2O + H2O →
Acid 1 Base 2
H3O+ + OH−
Acid 2
Base 1
Water, being amphoteric, can act as both an acid and a base. Here, one
molecule of water acts as an acid, donating a H+ ion and forming the conjugate base, OH−
and a second molecule of water acts as a base, accepting the H+ ion and forming the
conjugate acid, H3O+.
Acid dissociation and acid hydrolysis are seen to be entirely similar phenomena .
NH3 ( base ) + H2O (acid ) → NH4+ ( acid ) + OH – (base )
HCl (acid ) + H2O (base) → H3O+ (acid) + Cl− (base)
(iii) Lewis Concept
The hydrogen requirement of Arrhenius and Brønsted-Lowry was removed
by the Lewis definition of acid-base reactions. In this system, an acid does not exchange
atoms with a base but combines with it. For example, consider this classical aqueous
acid-base reaction.
HCl (aq) + NaOH (aq) → H2O (l) + NaCl (aq)
The Lewis definition does not regard this reaction as the formation of salt
and water or the transfer of H+ from HCl to OH−. Instead, it regards the acid to be the H+
ion itself and the base to be the OH− ion which has an unshared electron pair. Therefore,
the acid-base reaction here, according to the Lewis definition, is the donation of the
electron pair from OH− to the H+ ion. This forms a coordinate covalent bond between H+
and OH−, thus producing water (H2O).
By treating acid-base reactions in terms of electron pairs instead of specific
substances, the Lewis definition can be applied to reactions that do not fall under other
definitions of acid-base reactions. For example, a silver cation behaves as an acid with
respect to ammonia, which behaves as a base, in the following reaction:
Ag+ + 2 :NH3 → [H3N : Ag : NH3]+
The result of this reaction is the formation of an ammonia-silver adduct.
29
In reactions between Lewis acids and bases, there is the formation of an adduct when the
molecule, such as NH3 , donates lone pair of electrons to the electron-deficient molecule
such as BF3 through the formation of a co-ordinate covalent bond. The product of a
Lewis acid-base reaction, is a neutral, dipolar or charged complex which may be a stable
covalent molecule.
In highly-polar molecules, such as boron trifluoride (BF3),[ the most
electronegative element ( say F here ) pulls electrons towards its own orbitals, providing
a more positive charge on the less-electronegative element ( say B here ), thus making the
boron atom more electron deficient. In this case, it becomes more easy for the donor
( NH3) to donate the electron pair to the electron deficient molecule ( BF3).
Another example showing Lewis acid-base reaction is described below.
An electron deficient aluminium atom in AlCl3 binds to a chlorine atom of the
other molecule which donates one of its three lone pairs of electrons. The aluminium
atom then achieves the electron configuration of a noble gas – argon. Because this
sharing is unilateral ( chlorine atom contributes and donates both the electrons ) , chlorine
gets a formal positive charge and aluminium gets a formal negative charge. This is shown
in the following reaction. If the carbon – chlorine bond breaks, with both binding
electrons remaining with more electronegative chlorine atom, the carbon atom assumes
formal positive charge . We call such a carbon species as carbocation. All carbocations
are Lewis acids.
Many carbocations (but not all) may also function as Brønsted acids.
Following reaction illustrates this dual behavior. Lewis acidic site is colored red and
three of the nine acidic hydrogen atoms are colored orange. In its Brønsted acid role the
carbocation donates a proton to the base (hydroxide anion) and is converted to a stable
neutral molecule butene , having a carbon-carbon double bond. In the Lewis acid role,
the carbocation accepts a lone pair of electrons from the nucleophile OH – to form
tertiary butyl alcohol.
30
A terminology related to the Lewis acid-base nomenclature is often used by organic
chemists. Here the term electrophile corresponds to a Lewis acid and nucleophile
corresponds to a Lewis base.
( Electrophile is an electron deficient atom, ion or molecule that has an affinity for an
electron pair and will bond to a base or nucleophile. Nucleophile is an atom, ion or
molecule that has an electron pair that may be donated in bonding to an electrophile or
Lewis acid ).
All Bronsted - Lowry acid - base reactions are also Lewis acid - base reactions.
One advantage of the Lewis theory is the way it complements the model of
oxidation-reduction reactions. Oxidation-reduction reactions involve a transfer of
electrons from one atom to another, with a net change in the oxidation number of one or
more atoms.
The Lewis theory suggests that acids react with bases to share a pair of
electrons, with development of formal positive and negative charges on two atoms. Many
chemical reactions can be sorted into one or the other of these classes. Either electrons
are transferred from one atom to another, or the atoms come together to share a pair of
electrons.
The principal advantage of the Lewis theory is the way it expands the
number of acids and therefore the number of acid-base reactions. In the Lewis theory, an
acid is any ion or molecule that can accept a pair of nonbonding valence electrons. For
example, Al3+ ions form bonds to six water molecules to give a complex ion.
Al3+(aq) + 6 H2O(l)
[Al(H2O)6] 3+(aq)
This is an example of a Lewis acid-base reaction. The Lewis structure of
water suggests that this molecule has nonbonding pairs of valence electrons and can
therefore act as a Lewis base. The electron configuration of the Al3+ ion suggests that this
31
ion has empty orbitals that can be used to hold pairs of nonbonding electrons donated by
neighboring water molecules.
Thus, the [Al(H2O)6]3+ ion is formed when an Al3+ ion acting as a Lewis acid
picks up six pairs of electrons from neighboring water molecules acting as Lewis bases to
give an acid-base complex, or complex ion .
The Lewis acid-base theory can also be used to explain why nonmetal oxides
such as CO2 dissolve in water to form acids, such as carbonic acid H2CO3.
CO2(g) + H2O(l)
H2CO3 (aq)
In the course of this reaction, the water molecule acts as an electron-pair
donor or Lewis base. The electron-pair acceptor is the carbon atom in CO2. When the
carbon atom picks up a pair of electrons from the water molecule, it no longer needs to
form double bonds with both of the other oxygen atoms as shown in the figure below
One of the oxygen atoms in the intermediate formed, when water is added to
CO2 , carries a positive charge; another carries a negative charge. After an H+ ion has
been transferred from one of these oxygen atoms to the other, all of the oxygen atoms in
the compound are electrically neutral. The net result of the reaction between CO2 and
water is, therefore, formation of carbonic acid, H2CO3
Practical applications of neutralization
1) Lithium hydroxide, an alkali, is used by astronauts to neutralize the dangerous levels
of carbon dioxide exhaled.
2) The sting of bees and ants releases formic acid which causes pains and irritation. This
acid can be neutralized by rubbing soap which contains alkali ( NaOH ) .
3) The effect of acid rain on the soil can be neutralized by adding slaked lime to it.
4) The sting of wasp contains an alkali which can be neutralized by rubbing an acid like
acetic acid which is present in vinegar.
5) Persons who suffer from acidity , drink cold milk. Stomach secretes gastric juice
which contains HCl. This is neutralized by the cold milk which is alkaline.
32
Nature provides neutralization
Nettle ( Urtica dioica ) is a herbaceous plant . It is a common wild plant. It has leaves
with stinging hair. They cause painful stings when touched accidentally due to the
secretion of methanoic acid. A traditional remedy for this sting is rubbing the area with
the leaf of dock plant (Rumex ) which often grows side by side of nettle plant in the
wild. It is alkaline in nature.
Activity 19 - Take 25 ml of 0.1 M NaOH solution in a conical flask.. Add 3 drops of
phenolphthalein indicator to it. The solution becomes pink. Add drop by drop 0.1 M HCl
to it by a burette and swirl the solution . At one stage, the solution becomes colourless.
This shows that the alkali has fully reacted with the acid and has formed salt and water.
Note the burette reading. This reaction is called the neutralization reaction. This is a
titration.
Test your understanding
1) What is neutralization ? Is aqueous medium necessary for neutralization ?
:
2) Identify the conjugate acid-base pair in the following reactions .
(i) HCO3 - + OH - → CO3 2 - + H2O
(ii) NH4NO3 + NaNH2 → NaNO3 + 2 NH3
3) Classify the following acid – base reactions according to :
(i) Arrhenius concept (ii) Bronsted – Lowry concept (iii) Lewis concept
a) H2C2O4 + 2 NH4OH → (NH4)2C2O4 + 2 H2O
b) CH3COOH + NH3 → CH3COONH4
c) BCl3 + Cl - → BCl4 –
4) State any two applications of neutralization reaction in day to day life.
5) A person suffering from acidity of stomach is advised to take ‘milk of
magnesia’ .Explain.
33
Concept
(i) Salts
(ii) Classification of salts
Salts
In chemistry, salts are ionic compounds that can result from the neutralization
reaction of an acid and a base according to Arrhenius concept. . They are composed of
cations (positively charged ions) and anions (negatively charged ions) so that the product
is electrically neutral (without a net charge). These component ions can be inorganic such
as chloride (Cl−), as well as organic such as acetate (CH3COO−) and monatomic ions such
as fluoride (F−) as well as polyatomic ions such as sulfate (SO42−).
Classification of salts
Salts are classified as follows.
(i) Normal salt - The salt which does not contain any ionizable or replaceable hydrogen
atom in its molecule is called a normal salt. It is produced by complete replacement of
all the obtainable hydrogen ions by metallic or ammonium ions e.g. NaCl, K2SO4 .
NaOH + HCl → NaCl + H2O ;
2 KOH + H2SO4 → K2SO4 + 2 H2O
(ii) Acid salts - The salt which contains one or more replaceable hydrogen ions in its
molecule is called an acid salt. It is produced by the partial replacement of the
obtainable hydrogen ions by metallic or ammonium ions. Acid salts ionise in water
giving hydronium ions ( H3O) +. e.g. Na2HPO4 , KHSO4
2 NaOH + H3PO4 → Na2HPO4 + 2 H2O ; KOH + H2SO4 → KHSO4 + H2O
An acid salt can react with a base to form a normal salt.
(iii) Basic salts – The salt which contains one or more replaceable hydroxide ions in its
molecule is called a basic salt. It is produced by the partial replacement of the
obtainable hydroxide ions by acidic radicals. Basic salts ionize in water to give OH –
ions. e.g. Mg(OH)Cl, Pb(OH) NO3
Mg(OH)2 + HCl → Mg(OH) Cl + H2O ; Pb(OH)2 + HNO3 → Pb(OH)NO3 + H2O
A basic salt can react with an acid to form a normal salt.
(iv) Mixed salt – A salt which contains more than one acidic or basic radicals, other than
hydrogen or hydroxyl ion, in its molecule is called a mixed salt. e.g. Ca(OCl) Cl,
NaKSO4 . Mixed salts are prepared by different methods.
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Ca(OH)2 + Cl2 → Ca(OCl) Cl + H2O ; NaOH + KOH + H2SO4 → NaKSO4 + 2 H2O
(v) Double salt - A salt which is formed by the combination of two simple salts is called
a double salt. Such a salt is formed when the two salts are slowly crystallized together
from a mixture of their saturated salt solution. e.g. saturated solutions of KCl and
MgCl2 on crystallization form a double salt KCl,MgCl2,6H2O. Similarly, saturated
solutions of FeSO4 and (NH4)2SO4 on crystallization form a double salt
FeSO4 (NH4)2SO4. 6 H2O.
H2O
KCl + MgCl2. 6 H2O -------→
FeSO4 7 H2O + (NH4)2SO4
KCl. MgCl2. 6 H2O
H2O
----------→ FeSO4. (NH4)2SO4 . 6 H2O
(vi) Complex salt - The salt which contains a simple ion and a complex ion is called a
complex salt. It is formed by the combination of two simple salts under proper
experimental conditions. Examples of complex salts include [Co(NH3)6] Cl3 ,
K4 [Fe(CN)6].
CoCl3 + 6 NH3 → [Co(NH3)6] Cl3 ; Fe(CN)2 + 4 KCN → K4 [Fe(CN)6]
This salt dissociates in water to give a simple ion and a complex ion. The ion which
contains both metallic and non-metallic ions together which do not dissociate in water is
called a complex ion.
Aq
Aq
3+
[Co(NH3)6] Cl3 → [Co(NH3)6] + 3 Cl ; K4 [Fe(CN)6] → 4 K+ + [Fe(CN)6] 4 –
Comlex ion
Simple ion
Simple ion Complex ion
In aqueous medium, only the simple ion gives its characteristic tests.
Acidic salts and basic salts are defined in an alternative way also.
1) Basic salts - Salts that hydrolyze to produce hydroxide ions when dissolved in water
are called basic salts . These salts react with water to produce alkalinity in the solution.
Such salts are prepared by the reaction between a strong base and a weak acid.
For example, sodium carbonate is a salt of strong base NaOH and weak acid
H2CO3 . This salt hydrolyses in water to produce OH – ions. Hence though it does not
contain OH – ion, sodium carbonate is a basic salt.
Na2CO3 + 2 H2O → 2 Na+ + 2 OH - + H2CO3
Salt
Weak acid
Other examples of such salts include KCN, CH3COONa, Na3PO4 .
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2) Acidic salts - Salts that hydrolyze to produce H + ions when dissolved in water are
called acidic salts. These salts react with water to produce acidity in the solution. Such
salts are made by the reaction between a strong acid and weak base.
For example, ammonium chloride is a salt of strong acid HCl and a weak
base NH4OH. This salt hydrolyzes in water to produce H + ions . Hence though it
does not contain H + ion, ammonium chloride is an acidic salt.
NH4Cl + H2O → NH4OH + H + + Cl –
Salt
Weak base
Other examples of such salts include FeCl3 , Al2(SO4)3, NH4NO3 .
3) Neutral salts – Salts that do not hydrolyze when dissolved in water are called neutral
salts. All normal salts are neutral salts. Such salts are made by the reaction between a
strong acid and a strong base. For example, sodium chloride is a salt of strong base
NaOH and strong acid HCl. On dissolution in water, the base and acid fully ionize and
produce equal number of H + and OH – ions. Hence this salt solution is neither acidic
nor basic i.e. it is neutral.
NaOH + HCl → Na+ + OH - + H + + Cl –
Other examples of such salts include KCl , K2SO4, NaNO3 .
Activity 20 - Prepare salt ( NaCl ) and water as mentioned in activity 15. Evaporate the
salt solution to dryness in an evaporating dish .You have prepared a normal salt , NaCl,
by the reaction between an acid and an alkali.
Activity 21 – Take about 5 g NaCl in a test tube. Carefully add 5 ml concentrated H2SO4
to it. Heat the test tube with caution for 2 minutes on a moderate flame. This heating
raises the temperature between 2000 to 3000 C. The salt NaCl reacts with H2SO4 to form
HCl according to following equation.
NaCl + H2SO4
→
NaHSO4 + HCl ↑
Heat the substance in the test tube, till it becomes dry. You have prepared an acidic salt
NaHSO4.
If the salt NaCl is heated with H2SO4 on a strong flame for a long time , the temperature
goes above 5000 C and the reaction takes place as follows.
2 NaCl + H2SO4
→
Na2SO4 + 2 HCl ↑
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.
Uses of salts
Many salts find use in day to day life.
1) Potassium permanganate and potash alum are used to purify water.
2) Magnesium sulphate ( Epsom salt ) is used in purgatives.
3) Baking soda is used in cooking so also as a mild antacid.
4) Common salt is added to add taste to our food.
5) Plaster of paris is used for making statues and for false ceiling.
6) Ammonium sulphate, ammonium nitrate and some phosphates are used as fertilizers.
7) Silver bromide is used in photography.
8) Calcium carbonate is used in making chalk powder.
Test your understanding
1) Classify the following salts in different categories.
(i) Pb(OH)Cl,
(ii) CaMg(CO3)2 , (iii) KNO3 .
(v) [Cu(NH3)4] SO4 , (vi) FeCl3 , (vii) KCN ,
(iv) K2SO4.Al2(SO4)3 .24 H2O,
(viii) MgNH4 PO4
2) Which of the following salts is used in kitchen for cooking ?
Baking soda, Washing soda, Bleaching powder
3) Name the sodium salt which is used to soften hard water.
4) Explain the following :
(i) A milkman adds a little salt – NaHCO3 to milk to protect it from spoiling .
(ii) Sodium carbonate is a normal salt yet its solution has pH greater than 7.
(iii) Ferric chloride solution is stored in air tight bottles.
5) Give the name of the acid – salt found in ‘ Health salts’.
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Concept
Typical properties of some water soluble salts
Typical properties of some water soluble salts
(i) Water of crystallization - Several salts contain definite amount of water molecules
loosely combined in their structure. This chemically combined water present in the salt
accounts for the crystalline nature of the salt. Hence this water is called ‘water of
crystallization’. Salts containing water of crystallization are called hydrated salts. The
examples of hydrated salts include FeSO4. 7 H2O, CuSO4. 5 H2O, Na2CO3. 10 H2O.
(ii) Efflorescence - Certain hydrated crystals, when exposed to the atmosphere at
ordinary temperature, lose their water of crystallization and are transformed into a
powder. Such substances are called efflorescent substances and the property is called
efflorescence. Effloresce is maximum when the temperature is high and the atmosphere is
dry. The examples of efflorescent substances include Na2SO4. 10H2O, MgSO4.7 H2O.
(iii) Deliquescence - Certain substances, when exposed to the atmosphere at ordinary
temperature, absorb moisture from the atmosphere or air . They become moist, lose their
crystalline form and ultimately dissolve in the absorbed water forming a saturated
solution. Such substances are called deliquescent substances and the property is called
deliquescence. Deliquescence is maximum when temperature is low and atmosphere is
humid. The examples of deliquescent substances include CaCl2. 6 H2O, Cu(NO3)2. 3 H2O,
MgCl2.6 H2O.
(iv) Hygroscopy - Certain substances, when exposed to the atmosphere at ordinary
temperature, absorb moisture from the atmosphere without dissolving in it . Such
substances are called hygroscopic substances and the property is called hygroscopy.
Hygroscopic substances are generally anhydrous solids or liquids. They are used as
drying agents for drying gases. The examples of hygroscopic substances include CaCl2,
CaO, Silica gel, Ethyl alcohol.
Activity 22 - Take about 2 g of solid sodium hydroxide flakes. Keep them in open air for
2 to 3 days. You will find that the solid has turned moist . Sodium hydroxide is a
deliquescent substance. It absorbs moisture from air and tries to dissolve in it.
Activity 23 - Take about 2 g ordinary common salt in a dish in a rainy season. Keep it
open to atmosphere for about 2 hours. You will find it has become wet. Ordinary
common salt has an impurity of magnesium chloride which is highly deliquescent . In
rainy season or in humid conditions, MgCl2 absorbs moisture from air and dissolves in it.
This makes the salt wet.
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Test your understanding
1) What do you understand by ‘water of crystallization’ of a salt ?
2) Taking sodium carbonate as an example, give the meaning of the following terms :
a) anhydrous salt
b) efflorescence
3) Explain the following :
a) Aanhydrous calcium chloride used in a desiccator
b) Copper sulphate crumbles down to white powder after long exposure to air.
c) Copper nitrate crystals turn to a liquid on keeping for a long time.
d) On exposure to air, Glauber’s salt ( Na2SO4. 10 H2O ) looses weight while quick
lime ( CaO ) gains weight.
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.