Bronsted Lowry and Lewis Acid-Base theory

Introduction

The theories of Bronsted Lowry and Lewis acid-base theories were not developed by any single scientist at a particular time. These evolve over time . The Bronsted Lowry acid-base theory was developed by the scientists Johannes Nicolaus Bronsted and Thomas Martin Lowry. and their theories were published in the year of 1923. As compared to the Arrhenius theory these theories provide the broader context of acids and bases. According to the Bronsted theories the acids are those Which can donate the proton and bases are those that are the good acceptors of protons. This help to occur S the acid-base reaction at a wider extent not only in aqueous solution but also in other solvents. After donating a protein and accepting a proton acid base formed there conjugate pairs

The other Lewis acid-base theory was developed by Gilbert N Lewis in the year of 1923, the same as that of Bronsted theory but both do their work independently. Lewis says in his theory that Acids are those that can accept a pair of electrons and Bases are those that can donate a pair of electrons. His theory can be useful for that reaction where the transfer of proton does not take place. As in the reaction or compound formation of coordination chemistry. In Lewis acid-base reactions, the base provides a pair of electrons to form a bond with the acid, resulting in a new molecule or complex.

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.
$\mathrm{HCl}+\mathrm{NH}_3 \rightleftharpoons \mathrm{Cl}^{-}+\mathrm{NH}_4^{+}$
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 a proton. A strong acid is a substance that loses a proton easily to a base. Consequently, the conjugate base of a strong acid is weak.

$\underset{\text { ong Acid }}{\mathrm{HCl}}+\mathrm{H}_2 \mathrm{O} \longrightarrow \underset{\text { Weak base }}{\mathrm{Cl}^{-}}+\mathrm{H}_3 \mathrm{O}$

$\mathrm{HS}^{-}+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{H}_3 \mathrm{O}^{+}+\mathrm{S}^2$
The ability of an acid to lose a proton is experimentally measured by its equilibrium constant known as K_a. The larger the value of K_a, the more complete the reaction or the higher the concentration of H₃O⁺ and the stronger 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 that can act either as acids or as bases, NaSH, NaHCO_3, 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 possesses at least one unshared pair of electrons.

Substances that are based on the Bronsted system are also based on the Lewis concept. However, the Lewis definition of an acid considerably expands the 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:$\mathrm{Ni}+4 \mathrm{CO} \rightarrow \mathrm{Ni}(\mathrm{CO})_4$

Compounds that have central atoms capable of expanding their valence shells are Lewis acids in reactions in which this expansion occurs

$\mathrm{SnCl}_4+2 \mathrm{Cl}^{-} \rightarrow \mathrm{SnCl}_6^{2-}$

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Some Solved Examples

1. Which of the following salts is the most basic in aqueous solution?

1) $\mathrm{Pb}\left(\mathrm{CH}_3 \mathrm{COO}\right)_2$
2) $\mathrm{Al}(\mathrm{CN})_3$
3) (correct) $\mathrm{CH}_3 \mathrm{COOK}$
4) $\mathrm{FeCl}_3$

Solution

As we learned from the concept
$
\mathrm{CH}_3 \mathrm{COOK}+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{CH}_3 \mathrm{COOH}+\mathrm{KOH}
$Out of the given salts, this is the only one which is the salt of a strong base and hence, it is the most basic.

Hence, the answer is an option (3).

2. The following equilibrium is established when HClO₄ is dissolved in weak acid HF.

$\mathrm{HF}+\mathrm{HClO}_4 \rightleftharpoons \mathrm{ClO}_4^{-}+\mathrm{H}_2 \mathrm{~F}^{+}$

Which of the following is the correct set of conjugate acid-base pairs?

1) HF and $\mathrm{HClO}_4$
2) HF and $\mathrm{ClO}_4^{-}$
3) (correct) HF and $\mathrm{H}_2 \mathrm{~F}^{+}$
4) $\mathrm{HClO}_4$ and $\mathrm{H}_2 \mathrm{~F}^{+}$

Solution

Conjugate acid-base pair -

The acid-base pair that differ only by one proton is called a conjugate acid-base pair.

$\mathrm{HCl}+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{H}_3 \mathrm{O}^{+}+\mathrm{Cl}^{-}$

Cl^- is the conjugate base of HCl & H₃O⁺ is the conjugate acid of H₂O.

In the given case,

$
\mathrm{HF}^{+} \xrightarrow{+\mathrm{H}^{+}} \mathrm{H}_2 \mathrm{~F}^{+}
$
(Acid) (base)

Hence, the answer is the option (3).

3. Which of the following is a Lewis acid?

1) $\mathrm{PH}_3$
2) (correct) $\mathrm{B}\left(\mathrm{CH}_3\right)_3$
3) NaH
(4) $\mathrm{NF}_3$

Solution

Lewis acids and bases -

Lewis defined an acid as a species that accepts an electron pair and a base that donates an electron pair.

e.g BF₃

$B F_3+N H_3 \rightarrow B F_3: N H_3$

In the given question,

It is electron deficient, so it can accept a lone pair of electrons and behave as Lewis acid.

4. Which of the following are Lewis acids?

1) (correct)BCl₃ and AlCl₃

2)PH₃ and BCl₃

3) AlCl₃ and SiCl₄

4) PH₃ and SiCl₄

Solution

Lewis acids and bases -

Lewis defined an acid as a species that accepts an electron pair and a base that donates an electron pair.

In Lewis acid, many acids do not have protons.

e.g. $B F_3$

$\mathrm{BF}_3+\mathrm{NH}_3 \rightarrow \mathrm{BF}_3: \mathrm{NH}_3$

(i) BCl₃ and AlCl₃ have vacant p-orbitals

(ii) AlCl₃ has a vacant p-orbital, while SiCl₄ has a vacant d-orbital.

Hence, the answer is an option (1).

5. Species acting as both Bronsted acid and base is

1) (correct)(HSO₄)^-1

2)Na₂CO₃

3)NH₃

4)OH^-1

Solution

Bronsted-Lowry Acids and Bases-

According to Bronsted Lowry's theory, acid is a substance that is capable of donating a hydrogen ion H⁺ and bases are a substance that is capable of accepting H⁺ ions.

$\begin{aligned} & \mathrm{NH}_3+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{NH}_4^{+}+\overline{\mathrm{O}} \mathrm{H} \\ & \mathrm{NH}_4^{+} \rightarrow \text { add proton } \\ & \overline{\mathrm{O}} \mathrm{H} \rightarrow \text { lose proton }\end{aligned}$

The substance that can donate as well as accept ions can act as bronsted acid as well as bronsted base

Example HSO₄^-$\begin{aligned} & \mathrm{HSO}_4^{-}+\mathrm{H}^{+} \rightarrow \mathrm{H}_2 \mathrm{SO}_4 \\ & \mathrm{HSO}_4^{-} \rightarrow \mathrm{H}^{+}+\mathrm{SO}_4^{2-}\end{aligned}$

Hence, the answer is the option (1).

SUMMARY

Lewis's acid-base theory extends the concept of acids and bases more than their definitions cover. It includes reactions that are not covered by the Brønsted-Lowry or Arrhenius definitions, such as those involving compounds that do not donate or accept protons. It applies to a wide range of chemical reactions, including those in non-aqueous solvents and complex coordination chemistry. It helps in understanding reactions involving metal ions, ligands, and catalysts. Lewis acid-base reactions include the formation of complex ions, coordination compounds, and other processes where electron pair transfer occurs. It is widely used in organic chemistry, catalysis, and materials science.

Bronsted Lowry Acid Base Theory emphasizes the transfer of protons, which is a central aspect of many chemical reactions, particularly in aqueous solutions. It applies to reactions in both aqueous and non-aqueous solvents, providing a clearer understanding of acid-base behavior in various environments. It helps to explain the acid-base reactions that involve the formation and breaking of bonds involving protons, which is common in biological and chemical processes. In this theory, each acid has a corresponding base (its conjugate base) and vice versa. For example, in the reaction of hydrochloric acid (HCl) with water, HCl donates a proton to water, forming the conjugate base Cl⁻ and the conjugate acid H₃O⁺.