# How does a Lewis acid differ from an oxidizing agent; Lewis bases from reducing agents?

A Lewis acid is defined as an electron acceptor. An oxidizing agent is also defined as an electron acceptor (in the electronic theory of oxidation).

So then what is the basic difference between an Lewis acid and an oxidizing agent?

Similarly;

A Lewis base is defined as an electron donor (Lewis theory). A reducing agent is also defined as an electron donor (electronic theory). Therefore, what is the difference between Lewis bases and reducing agents? Are they exactly same thing?

There is a similar question which, however, only considers bases and nucleophiles and does not go into detail about oxidation and reduction.

• For Lewis acid and base theory, the electrons must be in a pair. Also, we don't think of them as being transferred. The donation or acceptance of the pair creates a bond. Whereas for oxidation/reduction, you transfer some number of electrons, not necessarily 2. – Zhe Nov 23 '16 at 16:14
• @Zhe make it an answer. – Always Confused Aug 9 '18 at 13:11
• Could I write "all lewis acids are oxidising agents but all oxidising agents are not lewis acid"? similarly, "all lewis bases are reducing agents, but not all reducing agents are lewis bases"? If my understanding is too poor, please forgive me but I prefer to clear confusion. – Always Confused Aug 9 '18 at 13:15

An acid and an oxidising agent are very different concepts that apply to different phenomena — although a single compound can be both at the same time, e.g. $\ce{H2SO4}$. However, a compound can also be an acid and a reducing agent, e.g. $\ce{HI}$.

The Brønsted acid/base theory is clearly distinct from any type of redox reaction, so I suppose that the confusion lies with the extension of said theory that the Lewis theory provides. But first a short word on the Brønsted theory: therein, an acid is a compound that can transfer a $\ce{H+}$ ion onto another compound.

The Lewis theory realises, that the acidic transfer reagent need not be hydrogen. Instead, anything with an unoccupied low-energy orbital can function as a Lewis acid; most notably boron compounds and metal cations. Upon creation of a Lewis acid-base pair, a coordinate bond is formed. Basically, this means that the energy of the lone pair of the Lewis base is lowered because it interacts with the empty orbital of the Lewis acid whose energy is raised. This can be depicted as $\ce{LB\bond{->}LA}$. The key concept is that the interaction is fully reversible and that upon dissociation the Lewis base keeps its lone pair.

When discussing oxidising or reducing agents, however, the electrons are irreversibly transferred. They actually move from orbitals of the reducing agent (which will be oxidised) into orbitals of the oxidising agent (which will be reduced). This transfer is thermodynamically favourable and thus the electrons won’t just ‘go back’.

To exemplify this, let’s consider two reactions of $\ce{H2SO4}$; one where it is an oxidising agent and one where it is an acid (Brønsted type; the proton of $\ce{H2SO4}$ is actually the Lewis acid).

$$\text{Acid-base:}\\ \ce{H2SO4 (l) + NH3 (g) <=>> NH4+ HSO4- (s)}$$

Note that ammonia, our base, is a gaseous compound at room temperature. If you employ appropriate conditions (i.e. temperature and low pressure), the equilibrium will be drawn to the left, releasing ammonia as ammonia ($\ce{NH3}$ molecules) including its lone pair. $\ce{H2SO4}$ will stay behind.

$$\text{Oxidising agent:}\\ \ce{2H2SO4 + Cu -> CuSO4 + SO2 ^ + 2 H2O}$$

$\ce{SO2}$, being a gas, will diffuse away. However, even if you perform this in a pressurised vessel, there will be no significant back-reaction because the electrons have been fully transferred from copper to sulfur.

The argument for Lewis bases/reducing agents is the same in reverse.

• Thanks, but I am still confused. Could you add a brief statement about "if this, then Lewis base, and if this, then reducing agent" (similarly for acid and oxidising agents) – Always Confused Aug 9 '18 at 13:11