Table of Contents
Redox Reactions as Electron and Atom Donor–Acceptor Processes
In this chapter, redox reactions are viewed specifically as donor–acceptor processes. The general idea of redox, as introduced in the parent chapter, is now translated into precise language: Who gives what to whom? and how can this be recognized in concrete reactions?
Electron-Transfer View: Donor and Acceptor
The most compact description of a redox reaction is:
- Reductant (reducing agent) = electron donor
- Oxidant (oxidizing agent) = electron acceptor
In a redox process, electrons pass from the reductant to the oxidant:
$$
\text{Reductant} \rightarrow \text{Reductant}^{\text{oxidized}} + e^-
$$
$$
\text{Oxidant} + e^- \rightarrow \text{Oxidant}^{\text{reduced}}
$$
Adding these partial processes (later called half-reactions) gives the overall redox reaction.
Key donor–acceptor roles:
- The reducing agent is oxidized: it loses electrons (donor).
- The oxidizing agent is reduced: it gains electrons (acceptor).
This is a donor–acceptor concept fully analogous to the idea that in acid–base reactions, a Brønsted acid donates protons and a Brønsted base accepts protons. Here, however, electrons are the "particles" being donated and accepted.
Example: Reaction Between Zinc and Copper(II) Ions
Consider the reaction of zinc metal with a solution containing copper(II) ions:
$$
\text{Zn (s)} + \text{Cu}^{2+}(\text{aq}) \rightarrow \text{Zn}^{2+}(\text{aq}) + \text{Cu (s)}
$$
Written as electron donor–acceptor steps:
- Electron donation:
$$
\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-
$$ - Electron acceptance:
$$
\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}
$$
Roles:
- $\text{Zn}$: electron donor → reducing agent, is oxidized to $\text{Zn}^{2+}$.
- $\text{Cu}^{2+}$: electron acceptor → oxidizing agent, is reduced to Cu.
The reaction is nothing more than a coupling of these two donor–acceptor processes.
Oxidation as Electron Donation, Reduction as Electron Acceptance
Using the donor–acceptor language:
- Oxidation: process in which a species donates electrons.
- Reduction: process in which a species accepts electrons.
Thus, oxidation and reduction always occur together: every electron donated must be accepted by something else.
Some typical patterns:
- Metals often act as electron donors (reducing agents) and are oxidized to cations.
- Nonmetals in higher positive oxidation states often act as electron acceptors (oxidizing agents) and are reduced to lower oxidation states or to neutral species.
Atom-Transfer View: Oxygen, Hydrogen, and Electron Donors/Acceptors
Historically, oxidation and reduction were recognized in terms of atoms, not electrons. These atom transfers can also be interpreted as donor–acceptor processes.
Oxygen Transfer
In many classical reactions:
- Oxidation often corresponds to uptake of oxygen.
- Reduction often corresponds to loss of oxygen.
Example:
$$
\text{C} + \text{O}_2 \rightarrow \text{CO}_2
$$
Here, carbon behaves as an oxygen acceptor: it accepts oxygen atoms from $\text{O}_2$. On the electron level, carbon is the reducing agent (electron donor) and $\text{O}_2$ is the oxidizing agent (electron acceptor). Thus:
- Oxygen is donated by $\text{O}_2$ and accepted by C.
- Electrons are donated by C and accepted by $\text{O}_2$.
This shows that in many reactions, an oxygen atom donor (like $\text{O}_2$) is simultaneously an electron acceptor.
Hydrogen Transfer
In many inorganic and organic reactions:
- Oxidation can correspond to loss of hydrogen.
- Reduction can correspond to gain of hydrogen.
Example:
$$
\text{CuO} + \text{H}_2 \rightarrow \text{Cu} + \text{H}_2\text{O}
$$
Interpretation as donor–acceptor processes:
- $\text{H}_2$ acts as a hydrogen donor. It also donates electrons (is oxidized to $\text{H}^+$ in the intermediate steps).
- $\text{CuO}$ acts as a hydrogen acceptor and, at the same time, as an oxygen donor (giving oxygen to hydrogen).
- On the electron level:
- $\text{H}_2$ is the electron donor (reducing agent).
- $\text{Cu}^{2+}$ (in $\text{CuO}$) is the electron acceptor (oxidizing agent).
Hydrogen and oxygen transfers are therefore alternative descriptions of underlying electron donor–acceptor processes.
Internal (Disproportionation) and External (Comproportionation) Donor–Acceptor Processes
Donor–acceptor roles do not always involve two different elements or species; sometimes one and the same element participates simultaneously as donor and acceptor in different parts of the system.
Disproportionation: One Species as Both Donor and Acceptor
In a disproportionation reaction, a single species is both oxidized and reduced:
- It donates electrons in one pathway.
- It accepts electrons in another.
General scheme:
$$
2\,\text{X}^{n+} \rightarrow \text{X}^{(n+1)+} + \text{X}^{(n-1)+}
$$
Example: Disproportionation of hydrogen peroxide in basic solution:
$$
3\text{H}_2\text{O}_2 \rightarrow 2\text{H}_2\text{O} + \text{O}_2
$$
Part of the $\text{H}_2\text{O}_2$ acts as reductant (electron donor), and another part as oxidant (electron acceptor). The same chemical substance simultaneously provides both donor and acceptor roles.
Comproportionation: Two Oxidation States to One Intermediate
In a comproportionation reaction, species in two different oxidation states of the same element react to form an intermediate oxidation state:
General scheme:
$$
\text{X}^{(n-1)+} + \text{X}^{(n+1)+} \rightarrow 2\,\text{X}^{n+}
$$
One species donates electrons, the other accepts them; the product has an intermediate oxidation state. Donor–acceptor roles are thus divided between two different starting species containing the same element.
Electron-Transfer Chains and Redox Mediation
In many reactions, the ultimate electron donor and acceptor are not in direct contact. Instead, electrons are passed stepwise through mediators. Each step is itself a donor–acceptor process.
Example of a chain idea (schematic):
$$
\text{A} \rightarrow \text{B} \rightarrow \text{C}
$$
where:
- A donates electrons to B (A is oxidized, B is reduced).
- B then donates electrons to C (B is oxidized again, C is reduced).
Here, B plays a dual role:
- Electron acceptor in the first step.
- Electron donor in the second step.
Such mediator roles are important in many chemical and biological systems (details are treated elsewhere); conceptually, they are sequences of coupled donor–acceptor steps.
Recognizing Donor and Acceptor Roles in Practice
For any given redox reaction, you can systematically identify donor and acceptor roles by:
- Assigning oxidation numbers (covered in a later chapter).
- Determining which species increases its oxidation number:
- This species is oxidized → electron donor → reducing agent.
- Determining which species decreases its oxidation number:
- This species is reduced → electron acceptor → oxidizing agent.
Even without oxidation numbers, simple clues often exist:
- Neutral metals turning into cations: typically electron donors.
- Non-metal-containing ions with high positive oxidation numbers (e.g., $\text{MnO}_4^-, \text{Cr}_2\text{O}_7^{2-}$): typically electron acceptors.
- Molecular oxygen $\text{O}_2$: in most reactions, an electron acceptor (oxidizing agent).
- Molecular hydrogen $\text{H}_2$: in many reactions, an electron donor (reducing agent).
In every case, the essential donor–acceptor nature of a redox process remains:
- Electrons (and often atoms such as O or H) are transferred.
- One partner donates, the other accepts.
- Oxidation and reduction are intertwined halves of this donor–acceptor exchange.