The document discusses different concepts of acids and bases including: 1. Arrhenius concept which defines acids as substances that produce H+ ions in water and bases as substances that produce OH- ions. 2. Bronsted-Lowry concept which defines acids as proton donors and bases as proton acceptors. 3. Lewis concept which defines acids as electron pair acceptors and bases as electron pair donors and can explain acid-base reactions not involving proton transfer. Hydrogen bonding is also discussed as the electrostatic attraction between hydrogen atoms covalently bonded to electronegative atoms. Hydrogen bonding influences properties like melting/boiling points and aggregation of molecules.

1. Topic: Acids and Bases
1. Arrhenius Concept of Acids and Bases
In 1884, a Swedish chemist Svante Arrhenius framed the first successful concept of acids and
bases. According to Arrhenius concept,
An acid is a substance that contains hydrogen and produces hydrogen ions in the aqueous
solution.
A base is a substance that contains Hydroxyl groups and produces hydroxyl ions in the aqueous
solution.
For example, substances such as HNO3, HCl, CH3COOH are acids whereas substances such as
NaOH, KOH, NH4OH are bases,
1.1 Modern Statement:
An acid is a substance when dissolved in water, increases the concentration of Hydrogen ions.
A base is a substance when dissolved in water, increases the concentration of Hydroxyl ions.
1.2 STRONG AND WEAK ACIDS, BASES
Acids such as HCl, H2SO4 and HNO3, which are almost completely ionized in aqueous solutions
are termed as strong acids whereas acids such as ethanoic acid (CH3COOH) which is weakly
ionized is called weak acid.
Similarly, bases which are almost completely ionized in aqueous solution are called strong bases,
for example, NaOH and KOH. The bases such as NH4OH are only slightly ionized and are called
weak bases.

2. 1.3 Limitations:
i) This concept was limited to aqueous solutions only. It does not explain the reactions taking
place in non aqueous medium such as liquid ammonia.
ii) It is restricted to the compounds containing H and OH ions only. It does not explain the acidic
behavior of SO2 and basic behavior of NH3 which don’t produces H and OH ions.
2. Bronsted-lowry concept of Acids and Bases
In 1923, a Danish Chemist J.H. Bronsted and an English Chemist T.M. Lowry independently
proposed new definitions for acids and bases. The Brønsted-Lowry theory does not go against
the Arrhenius theory in any way - it just adds to it. They proposed that:
An acid is a substance that can donate a proton.
A base is a substance that can accept a proton.
From the above equations, it is obvious that acid base reactions according to Bronsted-Lowry
concept involve transfer of proton from the acid to a base. A substance can act as an acid only if
another substance capable of accepting a proton, is present.
2.1 CONJUGATE ACID-BASE PAIRS
An acid after losing a proton becomes a base whereas a base after accepting the proton becomes
an acid. For example, let us consider the reaction between water and ammonia as represented by
the following equilibrium

3. In the forward reaction, water donates a proton to ammonia (base) and acts as acid. In the reverse
reaction, NH4
+
ions donate a proton to the OH-
ions (base) and act as acid. A base formed by the
loss of proton by an acid is called conjugate base of the acid whereas an acid formed by gain of a
proton by the base is called conjugate acid of the base. In the above example, OH-
is the
conjugate base of H2O and NH4
+
is conjugate acid of NH3 base.
A strong acid would have large tendency to donate a proton. Thus, conjugate base of a strong
acid would be a weak base. Similarly, conjugate base of a weak acid would be a strong base.
Some more conjugated acid-base pairs have been given in the following equations:
It can be noticed that in equation (1) H2O is behaving as a base whereas in equation (2) it is
behaving as an acid. Similarly, HCO3
–
ion in equation (3) acts as an acid and in equation (4) it
acts as a base. Such substances which can act as acids as well as bases are called amphoteric
substances.
2.2 Bronsted concept is superior to Arrhenius
a) Arrhenius concept was restricted to the substances which release H+ and OH-
ion in water
while Bronsted L. concept embraces all molecules that can donate proton and those which can
accept a proton.
b) This concept is not limited to the aq. solutions. It can be extended to even gas phase e.g.
gaseous ammonia reacts with HCl gas to produce ammonium chloride. Here proton is donated by
HCl to ammonia.

4. 3. Lewis Acids and Bases
Although Bronsted-Lowry theory was more general than Arrhenius theory of acids and bases but
it failed to explain the acid base reactions which do not involve transfer of proton.
For example, it fails to explain how acidic oxides such as anhydrous CO2, SO2, SO3, etc., (having
no proton) can neutralize basic oxides such as CaO, BaO, etc., even in the absence of solvent.
In 1923G.N. Lewis proposed broader and more general definitions of acids and bases, which do
not require the presence of protons to explain the acid base behavior. According to Lewis
concept,
An acid is a substance which can accept a pair of electrons.
A base is a substance which can donate a pair of electrons.
Acid-base reactions according to this concept involve donation of electron pair by a base to an
acid to form a co-ordinate bond. Lewis bases having one or more unshared pairs of electrons,
such as
or anions such as CN–
, OH–
, CI–
, etc.

5. Classification of Lewis acids
i) All simple cations Cu+2
, Co+2
, Ag+1
, Mg+2
are Lewis acids.
ii) An atom, ion, or molecule with an incomplete octet of electrons can act as an Lewis acid
(e.g., AlCl3 ). AlCl3 it forms three bonds and hence outer shell has 6 electrons. Now Al needs
two more electrons to complete its octet. By definition those which accepts electrons are
called lewis acids. So AlCl3 is a Lewis acid.
iii) Molecules that have multiple bonds between two atoms of different electronegativities (e.g.,
CO2, SO2).
Classification of Lewis bases
i) All the anions can act as Lewis bases because they have tendency to donate electron pairs. e.g.
Oxides, OH-
, Halides X-
, act as Lewis bases.
ii) Substances having lone pair or of electrons are bases e.g. water, ammonia, ether, Ketones.
iii) Compounds containing C=C double bond can also act as base.
Amphoterism
As of now you should know that acids and bases are distinguished as two separate things
however some substances can be both an acid and a base. You may have noticed this with water,
which can act as both an acid or a base. This ability of water to do this makes it an amphoteric
molecule. Water can act as an acid by donating its proton to the base and thus becoming its
conjugate acid, OH-. However, water can also act as a base by accepting a proton from an acid to
become its conjugate base, H3O+
.
• Water acting as an Acid:
H2O + NH3 → NH4
+
+ OH-
• Water acting as a Base:
H2O + HCl → Cl-
+ H3O+

6. The Lewis concept has many important applications because of its great generality. Many
substances which do not fit Arrhenius or Bronstted L. criteria quite logically classified as Lewis
acid or Lewis base. However it has some limitations.
LIMITATIONS OF LEWIS CONCEPT
1. It is too general and includes all the co-ordination compounds and co-ordination reactions.
2. It does not help to assign the relative strengths of acids and bases.
3. It does not explain the behavior of protonic acids such as HCl, H2SO4, etc., which do not
form coordinate bond with bases which is the primary requirement of Lewis theory.
Normally, formation of co-ordination compounds is slow. Therefore, acid-base reactions should
also be slow, but in actual practice, acid-base reactions are extremely fast.
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Topic: Hydrogen Bonding
Interaction between hydrogen atom of one molecule and the electronegative atom of the other
molecule is referred to as hydrogen bond. Thus, hydrogen bond is defined as the electrostatic
force of attraction which exists between the covalently bonded hydrogen atom of one molecule
and the electronegative atom of the other molecule.
The hydrogen bond is represented by dotted line( ….. )
For example, in case of hydrogen fluoride the hydrogen bond exists between H atom of one
molecule and fluorine atom of another molecule as shown.
It may be noted that the hydrogen atom acts as a bridge between the electronegative atoms of
two different molecules, to one atom it is linked through a covalent bond while to the other it is
linked through a hydrogen bond.
STRENGTH OF HYDROGEN BOND
Hydrogen bond is much weaker than covalent bond and its strength ranges between 12.5 kJ/mol-
to 41.5 kJ/mol-
1.
The strength of hydrogen bond is governed by the following factors:

7. (i) Extent of polarization within the molecule.
The molecule must contain highly electronegative atom linked to H-atom. Higher the-
electronegativity of the atom greater is the polarization and larger will be the strength of
hydrogen bond.
(ii) Size of the electronegative atom should be small.
We know electronegativity increases from N to F and atomic radius also decreases in the same
order.
EXAMPLES OF HYDROGEN BONDING
Let us study some examples of the compounds which exhibit hydrogen bonding.
1. Hydrogen fluoride.
In hydrogen fluoride hydrogen atom is bonded to highly electronegative atom fluorine
(electronegativity = 4). It has been found that hydrogen fluoride consists of long zig-zag chains
of H-F molecules associated by H-bonds as shown in Fig.
Fig. Zigzag chains of H-F molecules.
2. Water
In water molecule, oxygen atom is bonded to two hydrogen atoms. Due to large
electronegativity, oxygen atom forms the negative centre whereas each of the hydrogen atoms
acquires a partial positive charge. Each O atom can form two hydrogen bonds as shown in Fig.
Fig. H-bonds in H2O molecules

8. 3. Ammonia.
In ammonia molecule, nitrogen, an electronegative atom is bonded to three hydrogen atoms. The
nitrogen atom forms a negative site of the molecule whereas each of three H-atoms acquires a
partial positive charge. The ammonia molecules are associated by H-bonds as shown in Fig.
4. Ethanol (C2 H5 OH) and Ethanoic acid (CH3COOH):
These molecules contain the highly electronegative oxygen atom linked to H-atom and H atom
form associated molecules as shown in Fig.
Fig. H-bonds in C2H5OH and CH3COOH .

9. EFFECT OF HYDROGEN BONDING ON THE PROPERTIES
Hydrogen bond has a marked influence on the properties of various substances as follow:
(i) Association. The hydrogen bonds link up molecules of the same substance to form
large aggregates. This is called association of molecules. For example, H-F molecules are
associated with one another by hydrogen bonds. The formula of hydrogen fluoride can be written
as (HF)n. Similarly enthanoic acid exists as dimer and pertains to formula (CH3COOH)2 The
association of molecule through hydrogen bonding results in the unexpected larger values of
many physical properties such as melting points, boiling points.
(ii) Melting and boiling points. The compounds whose molecules are associated with one
another by hydrogen bonds have abnormally high melting and boiling points. It is due to fact that
a large amount of energy is needed to overcome intermolecular hydrogen bonds and to separate
the molecules.
The molecules of glucose, sugar honey, carboxylic acids have polar -OH groups in their
molecules. Hence, their solubility, in water is also attributed to the ability of their molecules to
form H-bonds with water molecules.
TYPES OF HYDROGEN BONDS
There are two types of hydrogen bonds intermolecular hydrogen bonds and intramolecular
hydrogen bonds
1. Intermolecular Hydrogen Bonding. Intermolecular hydrogen bond is formed between two
different molecules of the same or different substances. For example:
(i) Hydrogen bond between the molecules of hydrogen fluoride.
(ii) Hydrogen bond between alcohol and water molecules. Intermolecular hydrogen bond results
into association of molecules. Hence, it usually increases the melting point, boiling point,
viscosity, surface tension, solubility, etc.
2. Intramolecular Hydrogen Bond. Intramolecular hydrogen bond is formed between the
hydrogen atom and highly electronegative atom (F, 0 or N) present in different bonds within the
same molecule.
Intramolecular hydrogen bond results in the cyclization of the molecules and prevents their
association. Consequently, the effect of intramolecular hydrogen bond on the physical properties
is negligible.

10. A few examples of molecules which form intramolecular hydrogen bonds are given below:
AGGREGATION OF WATER MOLECULES IN LIQUID STATE
In gaseous state, the individual covalent molecules H2O exist as such. However, in liquid state,
large aggregates of varying number of H20 units are formed because of their association through
intermolecular hydrogen bonds.
In ice, a solid state of water, each H2O molecule is tetrahedrally surrounded by four
neighbouring H2O molecules. There are four H atoms around each 0 atom. Two of the four H
atoms are bonded by covalent bonds whereas the other two are linked through hydrogen bonds as
shown in Fig. This gives highly ordered three dimensional structure having large vacant spaces
which may be compared to open cage.

11. Open cage like structure of ice
Due to open cage-like structure, ice has a relatively larger volume for a given mass of liquid
water. Consequently, density of ice is less than water and it floats over water.
HYDROGEN BONDS AND STRUCTURE OF MACROMOLECULES
Macromolecular structures such as proteins, carbohydrates, nucleic acids, etc., are joined
together by the formation of H-bond through their active groups. Some examples are being
discussed as follows:
(i) Proteins are polymers of a-amino acid and they have C=0 and –N-H groups in their
molecules. These groups form both intermolecular and intramolecular hydrogen bonds.
The alpha -helix structure (spiral structure) of proteins is due to intramolecular hydrogen
bonding. Further, the protein chains are held together by intermolecular hydrogen bonds to form
fibres as in hair, silk and wool, etc. It may be noted that curly nature of hair is because of S-S
cross links in addition to H-bonding. This cross linking occurs between one part of alpha-helix
chain over another.
Fig .Alpha helix structure of protein.