Acids and Bases

Welcome to our presentation on Acid and Base Equilibria

 OUR TEEM MEMBERS

 Seamul Islam (190525)

 MD. Awoal Hossain (190526)
 Sagor Hossain (190527)
 Nazmus Sakib (190528)
 Abdullah Al Mamun (190530)
 MD. Zubair Haque (190531)
 Nadia Akhter Aroni (190533)
 Basic concept of acid
Three concept of acid
 (a) Arrhenius concept
 (b) Bronsted-Lowry concept
 (c) Lewis concept
ARRHENIUS CONCEPT
 an acid is a compound that releases 𝐻 + ions in water; and a base is
a compound that releases 𝑂𝐻− ions in water.
 𝐻𝐶𝑙(𝑎𝑞) 𝐻 +𝑎𝑞 + 𝐶𝑙(𝑎𝑞)
−

_ −
 𝑁𝑎𝑂𝐻(𝑎𝑞) 𝑂𝐻(𝑎𝑞) + 𝑁𝑎(𝑎𝑞)
Limitations of Arrhenius
Concept
 Free 𝐻+ and 𝑂𝐻− ions do not exist in water.

Step:1 Step: 2
 Limited to water only.
−
 Some bases do not contain OH . Ex- 𝑁𝐻3, 𝐶𝑎𝑂
BRONSTED-LOWRY CONCEPT
 an acid is any molecule or ion that can donate a proton (H+ ) a base is
any molecule or ion that can accept a proton.
Examples of Bronsted acids and
bases
 HCl gas and H2 O.

 HCl and Ammonia, N𝐻3

Bronsted-Lowry concept is superior to
Arrhenius concept
 The scope of Bronsted-Lowry concept is wider than Arrhenius
concept .
 Bronsted-Lowry concept is not limited to aqueous solutions as
Arrhenius concept.

(Gas) (Gas)
Conjugate Acid-Base pairs
The acid (HA) and the conjugate base (A− ) that are related to each
other by donating and accepting a single proton, are said to
constitute a conjugate Acid-Base pair.
Relative strength of conjugate acid
and bases
A weak base has strong conjugate acid .
A weak acid has a strong conjugate base.
LEWIS CONCEPT OF ACIDS AND BASES
An acid is an electron-pair acceptor a base is an electron-pair donor.
Examples of Lewis reactions
 Reaction between 𝑂2− and 𝑆𝑂3

 Relation between 𝐹 − and 𝐵𝐹3

Water can act both as an acid and
a base
Water is an amphoteric substance. It can behave either as an acid or
a base.

Kk

(When substance can act as both an acid and a base they are called amphoteric.)
RELATIVE STRENGTH OF ACIDS
The strength of an acid is defined as the concentration of H+ ions in its
aqueous solution at a given temperature.
When a monoprotic acid (HA) dissolve in water then-
+
HA + H2 O H3 O + A−
+
For simplifying our discussion, we take H3 O = H+
Thus we can write the equilibrium reaction as
HA + H2 O H+ + A−
Applying the Law of Mass action to the acid dissociation equilibrium,
we can write

Therefore, the value of 𝐾𝑎 for a particular acid is a measure of its acid

strength or acidity.
VALUES OF 𝒌𝒂 FOR SOME COMMON MONOPROTIC ACIDS
RELATIVE STRENGTH OF BASES
According to the Arrhenius concept, a base is a substance which produces
−
OH ions in aqueous solution.
VALUES OF Kb FOR SOME COMMON WEAK
BASES

 The strength of base is relatively related with the value of 𝐾𝑏

BASICITY OF ACID ACIDITY OF BASE
Basicity of an acid is the number of hydrogen ions which can be
produced by one molecule of acid.

The number of inoizable hydroxide (oH-) ions present in one molecule

of base is called acidity of bases.
Base Acidity

NaOH 1

Ba(𝑂𝐻)2 2
THE pH OF SOLUTIONS

The negative of the base-10 logarithm (log) of the H+ concentration.

pH = -log[H+ ]

where [H+] is the concentration of hydrogen ions in moles per litre.

Alternative and more useful forms of pH definition are:
1
pH = log +
[H ]
[H+ ] = 10−𝑝𝐻
The measurement of pH
The pH of a given solution can be measured with the help of an
apparatus called pH meter. This consists of a voltameter connected to
two electrodes.
(a) a standard electrode
(b) a special electrode

Figure: A pH meter
Relation between 𝐻 , 𝑂𝐻 + −
𝑎𝑛𝑑 𝐾𝑊
−
In any aqueous solution, the product of [H+ ] and [OH ] is always equal
to K𝑾 . . At 25°C it is 1.0×10−14 . Thus,
−
[H+ ][OH ] = 1.0×10−14
The concentration of H+ and OH– ions can be calculated from the
expressions :
K𝑊
[H+ ] = −
[OH ]
− K𝑊
[OH ] = +
[H ]
BUFFER SOLUTION
A buffer solution is one which maintains its pH fairly constant even
upon the addition of small amounts of acid or base.
Two common types of buffer solutions are :
 Acid buffers. 𝑒. 𝑔., CH3 COOH + CH3 COONa
 Basic buffers. 𝑒. 𝑔., NH4 OH + NH4 Cl.
pH of Buffer solution: we use Henderson-Hasselbalch equation to pick
out pH of Buffer solution. The equation is,
[𝑠𝑎𝑙𝑡]
pH = p𝐾𝑎 + log
[𝑎𝑐𝑖𝑑]
[𝑠𝑎𝑙𝑡]
pOH = p𝐾𝑏 + log
[𝑏𝑎𝑠𝑒]
Significance of the Henderson-Hasselbalch
equation
 The pH of a buffer solution can be calculated from this equation.
 The dissociation constant of a weak acid (or weak base) can be
determined by this equation.
[𝑠𝑎𝑙𝑡]
pH = p𝐾𝑎 + log
[𝑎𝑐𝑖𝑑]
[𝑠𝑎𝑙𝑡]
if [salt] = [acid]. log = log 1 = 0
[𝑎𝑐𝑖𝑑]
pH = p𝐾𝑎

 A buffer solution of desired pH can be prepared by adjusting the

concentrations of the salt and the acid added for the buffer.
MECHANISM OF BUFFER SOLUTION
 Addition of HCl

 Addition of NaOH
ACID–BASE INDICATORS
An acid-base indicator is an organic dye that signals the end-point by
a visual change in colour.
Most indicators do not change colour at a particular pH. They do so over a
range of pH from two to three units. This is called the pH range which is
different for various indicators.
THEORIES OF ACID-BASE INDICATORS
Two theories have been put forward to explain the indicator action in acid-
base titrations :
 The Ostwald’s theory
 The Quinonoid theory
The Ostwald’s theory : According to this theory, a hydrogen ion indicator is a
weak organic acid or base. The undissociated molecule will have one colour
and the ion formed by its dissociation will have a different colour.
The Quinonoid theory : According to this theory, The acid-base indicators exist
in two tautomeric forms having different structures. Two forms are in
equilibrium. Phenolphthalein has benziod form in acidic medium and thus, it is
colourless while it has quinonoid form in alkaline medium which has pink
colour.
Thanks all

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