includes Henderson hasselbach equations, common ion effect, strength of acid and bases, le chatlier principle.basics of acid base titrations

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RELATIVESTRENGTHS OF ACIDS AND BASES
LAW OF MASS ACTION
COMMON-ION EFFECTS
HENDERSON – HASSELBALCH EQUATION
NEHLA P
MOULANA COLLEGE OF PHARMACY

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RELATIVE STRENGTHS OF ACIDS AND BASES
•The stronger acids are those that lose their protons more easily than
others. Similarly the stronger bases are those that hold onto protons
more strongly than others.
•An acid is strong if it completely ionizes in water. In the reaction of
Hydrogen chloride with water, for e.g., water acts as a base,
accepting the proton from HCI,
HCl(aq) + H2O(aq)  Cl- + H3O+
acid base base acid
The reverse reaction occurs only to an extremely small extent.
Because the reaction goes almost completely to the right, we say
that HCI is strong acid.

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Relative strength of acids
•The strength of an acid depends on its ability to transfer its proton
(H+) to a base to form its conjugate base.
•When a monoprotic acid (HA) dissolves in water, it transfers its proton
to water (a Bronsted base) to form hydronium ion (H3O+) and
conjugate base.
HA + H2O  H3O+ +A-
Acid conjugate base
For simplifying our discussion, we take
H3O+ = H+ Thus we can write the equilibrium reaction (1) as
HA + H2O  H+ +A- ….(2) This equation represents the
dissociation of the acid HA into H+ion and A- ion.

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Applying the law of mass action to the acid dissociation equilibrium,
we can write
Ka = [H+][A-]
--------------
[HA] (3)
[HA] where, Ka is called the acid dissociation constant.
It is evident that the concentration of H+ depends on the value of Ka.
Therefore the value of Ka for a particular acid is the measure of its acid
strength or acidity.

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Relative strength of Bases
According to the Arrhenius concept, a base is a substance which
produces OH ions in aqueous solution. The basic properties of a such
a substance are due to these hydroxyl ions. Let us consider a base
BOH whose dissociation can be represented as
BOH  B+ + OH-
Applying the law of mass action to the above equilibrium we can
write the equilibrium expression as
Kb = [B+][OH-]
----------------- (2)
[BOH]
Kb is called the base dissociation or base ionisation constant.

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•The strength of a base is defined as the concentration OH ions in its
aqueous solution at a given temperature.
•From the equilibrium expression (2), it is evident that the
concentration of OH ions (OH) depends on the value of Kb
Therefore, the value of Kb for a certain base is a measure of the base
strength.

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LAW OF MASS ACTION

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LAW OF MASS ACTION
Two Norwegian chemists, Guldberg and Waage, In 1867, postulated
generalisation called the Law of Mass Action. It states that,”The rate of
chemical reaction is proportional to the active masses on the reactants”.
By the term “active mass” is meant the molar concentration,. It is
expressed by enclosing the formula of the substances in square
brackets.
E.g., Consider the reaction.
CO(g) + 3H2  CH4 + H2O(g)
Kc = [CH4] [H2O]
-----------------
[CO] [H2]3
Kc = Equilibrium constant.

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Consider the general reaction
aA + bB--- cC + dD
where
A, B, C and D denote reactants and products
a, b, c and d are coefficients in the balanced chemical equation.
For the general reaction, we have
Kc = [C]c [D]d Products concentrations
------------
[A]a[B]c Reactants concentrations

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0
Hence the law of mass action is a relation that states that the values of
the equilibrium-constant expression K, are constant for a particular
reaction at a given temperature, whatever equilibrium concentration are
substituted.

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COMMON-ION EFFECTS
•The concentration of a particular ion in an ionic reaction can be
increased by the addition of a compound which produces that ion upon
dissociation.
•The particular ion is thus derived from the compound already in solution
and also from the added reagent, hence the name common ion.
•In the solution of an electrolyte in water, these exists an equilibrium
between the ions and the undissociated molecules to which the law of
mass action can be applied. Considering the dissociation of an
electrolyte, we have
AB ---- A+ + B- and
[A] +[B]-/[AB] = k (dissociation constant)

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• If now another electrolyte or soluble salt is added to the above
solution containing a common ion (A+B), it will result in the increase
of concentration of the ions A+ or B- and in order that k may remain
the same.
•The concentration of (AB) must eventually increase. In other words
the degree of dissociation of an electrolyte is suppressed by the
addition of an another electrolyte (or salt) containing common ion.
• Hence, The common ion affect is the shift in an ionic equilibrium
caused by the addition of a solute that provides an ion that takes part in
the equilibrium.
Thus the dissociation of an electrolyte NH4OH is diminished by the
addition of NH4Cl which furnishes the common NH4 ion.

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THE PH OF A SOLUTION
•The pH is defined as the negative of the logarithm of the molar
hydrogen ion concentration. Mathematically it may be expressed as.
pH = -log [H+]
where [H+] is the concentration of hydrogen ions in moles per litre.
•A neutral solution, whose hydrogen-ion concentration at 25°C is 1.0 x
10-7M, has a pH of 7.00.
•For an acidic solution, the Hydrogen ion concentration is greater than
1.0 x 10-7M, so the pH is less than 7.00.
•Similarly, a basic solution has a pH greater than 7.00.

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Determination of pH of a solution
The pH of a solution can be measured accurately by
1. Electrometric method (Potentiometric method)
2. Colorimetric method
1. Potentiometric method
The potentiometric method is the more accurate and it is the ultimate
standard of Hydrogen ion concentration. In practice, it is applied of
means of potentiometric measurements and with a pH meter.
The Hydrogen ion concentration of a solution can be determined
electrometrically by measuring the potential difference set up between
an electrode immersed in the solution and a standard electrode of
known potential e.g. standard hydrogen electrode.

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Colorimetric Determination of pH
•Colorimetric methods are based on the use of reagents which alter
colour in accordance with hydrogen ion concentration. These reagents
such as litmus (which contains Azolitimic acid, with red undissociated
molecules yielding blue anions) are acid base indicators.

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HENDERSON – HASSELBALCH EQUATION
The pH of a buffer solution is related to the pKa of the buffer acid and
the log of the ratio of buffering species. This relationship is described by
the Henderson-Hasselbalch equation, which can be derived from the
ionization constant expression for the weak acid.
pH = pKa + log [Conj. Base]
----------------
[Acid]
This is an equation relating the pH of a buffer for different
concentrations of conjugate acid and base, it is known as the Henderson
– Hasselbalch equation.

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The weak acid is only slightly dissociated and its dissociation is further
depressed by the addition of the salt (N+ A-) which provides A- ions
(common ion affect).
As a result the equilibrium concentration of the unionised acid is nearly
equal to the initial concentration of the acid. The equilibrium
concentration [A-] is presumed to be equal to the initial concentration
of the salt added since it is completely dissociated. Thus we can write
the equation as
[H+] = Ka [acid]/[salt] (2)
where [acid] is the initial concentration of the added acid and [salt] that
of the salt used.
Taking negative logs of both sides of the equation (2) we have
– log [H+] =- log Ka – log[acid]/[salt] (3)

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But, – log [H*]=pH and log Ka = pKa, Thus from (3) we have,
pH = pKa – log[acid]/[salt]
=pKa + log [salt]/[acid]
Hence pH = pKa + log[salt]/[acid]
This relationship is called the Henderson-Hasselbalch
equation or simply Henderson equation.

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In a similar way, the Henderson equation for a basic buffer
can be derived. This can be stated as,
POH= pka + log[salt]/[base]