Arrhenius, Brønsted-Lowry and Lewis Concepts — Revision Notes

⚡ 30-Second Revision

Arrhenius Acid: — Produces $H^+$ ($H_3O^+$) in water.
Arrhenius Base: — Produces $OH^-$ in water.
Brønsted-Lowry Acid: — Proton ($H^+$) donor.
Brønsted-Lowry Base: — Proton ($H^+$) acceptor.
Conjugate Pair: — Differ by one $H^+$. Acid $\leftrightarrow$ Conjugate Base + $H^+$.
Amphiprotic: — Can both donate and accept $H^+$ (e.g., $H_2O$, $HCO_3^-$).
Lewis Acid: — Electron pair acceptor (e.g., $BF_3$, $AlCl_3$, $H^+$, metal cations).
Lewis Base: — Electron pair donor (e.g., $NH_3$, $H_2O$, $OH^-$, $Cl^-$).
Scope: — Lewis > Brønsted-Lowry > Arrhenius.

2-Minute Revision

Quickly recall the three main acid-base concepts. Arrhenius is the simplest, defining acids as $H^+$ producers and bases as $OH^-$ producers, strictly in water. Its limitation is the aqueous-only requirement.

Brønsted-Lowry expands this, focusing on proton ($H^+$) transfer: acids donate protons, bases accept them. This introduces conjugate acid-base pairs, where an acid forms its conjugate base after donating a proton, and a base forms its conjugate acid after accepting one.

Substances like water that can do both are amphiprotic. This concept works in non-aqueous solvents. The most general is the Lewis concept, which shifts focus to electron pairs: acids accept electron pairs, and bases donate them.

Lewis acids are often electron-deficient (like $BF_3$, $AlCl_3$, or cations), while Lewis bases have lone pairs (like $NH_3$, $H_2O$, or anions). Remember that all Arrhenius and Brønsted-Lowry acids/bases can also be classified under Lewis, but not vice-versa.

For NEET, be ready to identify species based on each definition and correctly form conjugate pairs.

5-Minute Revision

To thoroughly revise acid-base concepts, start with the historical progression. Arrhenius definitions are narrow: acids increase $H_3O^+$ in water, bases increase $OH^-$ in water. Think $HCl \rightarrow H^+ + Cl^-$ and $NaOH \rightarrow Na^+ + OH^-$. Its main drawback is the 'aqueous only' restriction.

Next, the Brønsted-Lowry concept is broader, based on proton ($H^+$) transfer. An acid is a proton donor, and a base is a proton acceptor. This is crucial for understanding conjugate acid-base pairs.

For any acid $HA$, its conjugate base is $A^-$ (formed by losing $H^+$). For any base $B$, its conjugate acid is $BH^+$ (formed by gaining $H^+$). For example, in $HCl + H_2O \rightleftharpoons Cl^- + H_3O^+$, $HCl$ is the acid, $Cl^-$ is its conjugate base; $H_2O$ is the base, $H_3O^+$ is its conjugate acid.

Remember that amphiprotic substances like $H_2O$ or $HCO_3^-$ can act as both acid and base by donating or accepting a proton.

Finally, the Lewis concept is the most general, focusing on electron pair transfer. A Lewis acid is an electron pair acceptor (often electron-deficient, like $BF_3$, $AlCl_3$, $H^+$, or metal cations).

A Lewis base is an electron pair donor (has lone pairs, like $NH_3$, $H_2O$, $OH^-$, $Cl^-$). This explains reactions without protons, like $BF_3 + NH_3 \rightarrow F_3B \leftarrow NH_3$. The order of generality is Lewis > Brønsted-Lowry > Arrhenius.

Practice identifying each type in various reactions and forming conjugate pairs. Pay attention to the limitations of each theory to understand why the next one was developed.

Prelims Revision Notes

Arrhenius Theory:

- Acid: Produces $H^+$ (or $H_3O^+$) in aqueous solution. E.g., $HCl$, $H_2SO_4$. - Base: Produces $OH^-$ in aqueous solution. E.g., $NaOH$, $Ca(OH)_2$. - Limitation: Only for aqueous solutions; cannot explain basicity of $NH_3$ or acidity of $CO_2$.

Brønsted-Lowry Theory:

- Acid: Proton ($H^+$) donor. E.g., $HCl$, $H_2O$ (in presence of stronger base). - Base: Proton ($H^+$) acceptor. E.g., $NH_3$, $OH^-$, $H_2O$ (in presence of stronger acid). - Conjugate Acid-Base Pair: Two species differing by one proton.

Acid $\leftrightarrow$ Conjugate Base + $H^+$. Base + $H^+ \leftrightarrow$ Conjugate Acid. - Example: $H_2SO_4$ (acid) $\rightarrow HSO_4^-$ (conjugate base). - Example: $NH_3$ (base) $\rightarrow NH_4^+$ (conjugate acid).

- Amphiprotic/Amphoteric: Can act as both Brønsted-Lowry acid and base. E.g., $H_2O$, $HCO_3^-$, $H_2PO_4^-$. - Advantage: Explains reactions in non-aqueous solvents; explains $NH_3$ basicity. - Limitation: Requires proton transfer.

Lewis Theory:

- Acid: Electron pair acceptor. (Electrophiles). E.g., $BF_3$, $AlCl_3$, $H^+$, metal cations ($Fe^{3+}$), $SO_3$. - Base: Electron pair donor. (Nucleophiles). E.g., $NH_3$, $H_2O$, $OH^-$, $Cl^-$, $CN^-$. - Reaction: Forms a coordinate covalent bond. - Advantage: Most general theory; explains reactions without protons (e.g., $BF_3 + NH_3$); crucial for organic chemistry and coordination chemistry.

Relationship: — Lewis > Brønsted-Lowry > Arrhenius in terms of scope. All Arrhenius/Brønsted-Lowry acids/bases are also Lewis acids/bases, but the reverse is not always true (e.g., $BF_3$ is a Lewis acid but not Brønsted-Lowry).

NEET Focus: — Identify acid/base types, correctly form conjugate pairs, recognize amphiprotic species, understand limitations and scope of each theory.

Vyyuha Quick Recall

To remember the order and focus of the theories: All Boys Love Protons Everywhere.

Arrhenius: Focus on Aqueous solutions, Hydrogen ($H^+$) and OH ($OH^-$) ions.
Brønsted-Lowry: Focus on Proton ($H^+$) transfer.
Lewis: Focus on Electron Pairs (transfer).