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Bronsted Lowry and Lewis Acid-Base theory - Practice Questions & MCQ

Edited By admin | Updated on Sep 18, 2023 18:35 AM | #JEE Main

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Bronsted Lowry and Lewis Acid-Base theory is considered one the most difficult concept.
21 Questions around this concept.

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MEDIUM

The incorrect statement for the use of indicators in acid-base titration is :

Methyl orange may be used for a weak acid vs weak base titration.

Phenolphthalein is a suitable indicator for a weak acid vs strong base titration.

Methyl orange is a suitable indicator for a strong acid vs weak base titration.

Phenolphthalein may be used for a strong acid vs strong base titration.

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Concepts Covered - 2

Bronsted Lowry and Lewis Acid-Base theory

Bronsted-Lowry Acids and Bases

According to this concept, an acid and a base can defined as follows :
Acid: It is a substance that can donate a proton.
Base: It is a substance that can accept a proton.

Some examples include:

When HCl is dissolved in water, it donates a proton to H₂O which behaves as a base.
$\mathrm{HCl}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\ell) \rightarrow \mathrm{Cl}^{-}(\mathrm{aq})+\mathrm{H}_{3} \mathrm{O}^{-}(\mathrm{aq})$
When HCO₃^-is dissolved in water, it donates a proton to NH₃ which behaves as a base.
$\mathrm{HCO}_{3}^{-}(\mathrm{aq})+\mathrm{NH}_{3}(\mathrm{aq}) \rightleftharpoons \mathrm{CO}_{3}^{2-}(\mathrm{aq})+\mathrm{NH}_{4}^{+}(\mathrm{aq})$

The base formed from an acid is known as the conjugate base of the acid. Correspondingly, the acid formed from a base is called the conjugate acid of the base.
$\mathrm{HCl+N H_{3} \rightleftharpoons C l^{-}+N H_{4}^{+}}$
In the above reaction, Cl^- is the conjugate base of HCl and NH₄⁺is the conjugate acid NH₃.

Strength of Bronsted-Lowry Acid and Bases:
The strength of an acid or base is measured by its tendency to lose or gain proton. A strong acid is a substance which loses a proton easily to a base. Consequently, the conjugate base of a strong acid is a weak base.

The ability of an acid to lose a proton is experimentally measured by its equilibrium constant know as K_a. The larger the value of K_a, the more complete reaction or higher the concentration of H₃O⁺ and the stronger is the acid. Similarly, for bases, we have the equilibrium constant, K_b which determines the extent of the completion of the reaction.

Amphiprotic Compounds:
The compounds which can act either as acids or as bases, NaSH, NaHCO₃ etc are some of the examples.

Lewis Acid and Bases:
Acid: It is a substance that can form a covalent bond by accepting a shared pair of electrons.
Base: It is a substance that possess at least one unshared pair of electrons.

Substance that are bases in the Bronsted sysem are also bases according to the Lewis concept. However the Lewis definition of an acid considerably expands that number of substances that are classified as acid. A Lewis acid must have an empty orbital capable of receiving the electron pair of the base.
Lewis acids include molecules or atoms that have incomplete octets. For example molecules like BF₃, AlCl₃, etc. act as Lewis Acid.

Many simple cations can act as Lewis acids:
$\mathrm{Cu}^{2+}+4 \mathrm{NH}_{3} \rightarrow \mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}^{2+}$

Some metal atoms can function as acids in the formation of compounds such as:
$\mathrm{Ni}\: +\: 4\mathrm{C}\mathrm{O} \rightarrow \mathrm{Ni}(\mathrm{CO})_{4}$

Compounds that have central atoms capable of expanding their valence shells are Lewis acids in reactions in which this expansion occurs.
$\mathrm{SnCl}_{4}+2 \mathrm{Cl}^{-} \rightarrow \mathrm{SnCl}_{6}^{2-}$

Ionisation Constant of Acids and Bases and pH of strong Acids and Bases

Ionisation constant of acid is also known as dissociation constant or equilibrium constant at the dissociation of acid. Thus, the ionisation constant equation of acid is given as follows:

$\mathrm{K_{a}\: =\: K_{c}\: =\: \frac{[H+][A^{-}]}{[HA]}}$

Thus, K_a is very small for weak acids and it is very high for strong acids. For example:

$\\\mathrm{HCl\: \rightarrow \: H^{+}\:+\: Cl^{-}}\\\\\mathrm{Thus,\: K_{a}\: =\: K_{c}\: =\: \frac{[H+][Cl^{-}]}{[HCl]}}$
Since HCl concentration is low after the reaction due to 100% dissociation, K_a for Strong acids is very high.

Hence, as K_a increases, the strength of acid increases.

Similarly, for bases we have:

$\mathrm{MOH \rightleftharpoons M^++O H^{-}}$

Again, the equilibrium constant K_b or K_c is given as follows:

$\mathrm{K_{b}\: =\: K_{c}\: =\: \frac{[M]^{+}[OH]^{-}}{[MOH]}}$

Again, K_b is very small for weak bases and it is very high for strong bases like NaOH.

Hence, as K_b increases, the strength of base increases.

pH and Strong Acids/Bases

pH is also referred to as potential or power of hydrogen. Mathematically, it can be represented as follows:

$\mathrm{pH\: =\: -log_{10}[H^{+}]}$

If solution is neutral, then:
$\mathrm{k_w = [H^+][OH^-]=10^{-14}}$
On Solving we have,

$\mathrm{[H^+]=[OH^-]=10^{-7}}$
$\mathrm{Thus, \ pH = 7}$

For Acidic solutions: For Basic solutions:
For acidic solutions, we must have $\mathrm{[H^+] > [OH^-] }$ For basic solutions, we must have $\mathrm{[H^+] < [OH^-] }$
Thus, $\mathrm{[H^+] > 10^{-7} }$ Thus, $\mathrm{[H^+] < 10^{-7} }$
Thus, $\mathrm{[H^+] }$ for acids can be 10^-6, 10^-5, 10^-4, etc. Thus, $\mathrm{[H^+] }$ for bases can be 10^-8, 10^-9, 10^-10, etc.
Thus, pH of acids can be 6, 5, 4, etc. Thus, pH of basics can be 8, 9, 10, 11, etc.
Hence, pH of acidic solutions is less than 7 Hence, pH of basic solutions is greater than 7

pH of Strong Acids
Strong acids are those acids which dissociate completely in solutions. For example:

2 x 10^-3M HNO₃
Since HNO₃ is a strong acid, thus it will dissociate completely into H⁺ and OH^- ions as follows:

$\\\mathrm{HNO_{3}\: \rightarrow \: H^{+}\: +\: NO_{3}^{-}}\\\\\mathrm{Thus,\: [H^{+}]\: =\: 2\: x\: 10^{-3}M\: \: \: \: \: \: \: \: \: (given)}\\\\\mathrm{\Rightarrow\: pH\: =\: -log_{10}(2\: x\: 10^{-3})}\\\\\mathrm{\Rightarrow\: pH\: =\: -log_{10}(2)\: -\: log_{10}(10^{-3})}\\\\\mathrm{\Rightarrow\: pH\: =\: -0.30\: +\: 3\: =\: 2.7}$
Thus, pH of HNO₃ is 2.7

10^-4 M H₂SO₄
Since H₂SO₄ is a strong acid, thus it will dissociate completely into H⁺ and OH^- ions as follows:

$\\\mathrm{H_{2}SO_{4}\: \rightarrow \: 2H^{+}\: +\: SO_{4}^{2-}}\\\\\mathrm{Thus,\: [H^{+}]\: =\: 2\: x\: 10^{-4}M\: \: \: \: \: \: \: \: \: (given)}\\\\\mathrm{\Rightarrow\: pH\: =\: -log_{10}(2\: x\: 10^{-4})}\\\\\mathrm{\Rightarrow\: pH\: =\: -log_{10}(2)\: -\: log_{10}(10^{-4})}\\\\\mathrm{\Rightarrow\: pH\: =\: -0.30\: +\: 4\: =\: 3.7}$
Thus, pH of H₂SO₄ is 3.7