Redox Reactions and Electrode Processes

So far in this chapter, you have seen redox reactions where electrons jump directly from one substance to another, like when a zinc rod dipped into copper sulphate solution gets coated with copper while the blue colour fades. But what if you could force those electrons to travel through a wire instead of hopping directly between atoms? That simple idea is the bridge between redox chemistry and electricity, and it is exactly how batteries work.

From Direct Transfer to Indirect Transfer: Why It Matters

When zinc metal sits in a solution of $CuSO_4$ , the reaction is straightforward: zinc atoms hand their electrons directly to $Cu^{2+}$ ions. Zinc gets oxidised, copper gets reduced, and the solution heats up. The energy released is wasted as heat because the electron transfer happens right at the surface of the zinc rod.

Now imagine separating the two half-reactions physically. Put the zinc in one container and the copper sulphate in another. If you can find a way to connect them so that electrons still flow from zinc to copper ions, but this time through an external wire, you can capture that electron flow as an electric current. That is precisely what an electrochemical cell does.

Redox Couples: Pairing the Two Forms

Before setting up the cell, you need to understand a key concept. At each electrode, both the oxidised form and the reduced form of the same element are present together. For example, a zinc rod dipped in zinc sulphate solution has $Zn^{2+}$ ions (oxidised form) in the solution and $Zn$ metal (reduced form) as the rod. This pairing is called a redox couple (a pair consisting of the oxidised and reduced forms of a substance that participates in a half-reaction).

A redox couple is written with the oxidised form first, separated from the reduced form by a slash that represents the interface (typically solid/solution):

$Zn^{2+}/Zn \qquad \text{and} \qquad Cu^{2+}/Cu$

The first couple represents the oxidation half-reaction and the second represents the reduction half-reaction.

Building the Daniell Cell: A Step-by-Step Look

The Daniell cell is the classic electrochemical cell that demonstrates how a redox reaction can generate electricity. Here is how it is set up:

Two separate beakers: One contains $ZnSO_4$ solution with a zinc rod dipped in it. The other contains $CuSO_4$ solution with a copper rod dipped in it.
A salt bridge: A U-shaped tube filled with a solution of $KCl$ or $NH_4NO_3$ (solidified into a jelly using agar-agar) connects the two solutions. This bridge allows ions to migrate between the beakers, completing the internal electrical circuit, without letting the two solutions physically mix.
An external wire: The zinc rod and the copper rod are connected by a metallic wire. An ammeter and a switch are placed in this wire to detect and control the current.

Fig 7.3: The set-up for Daniell cell

What Happens When the Switch is Turned On

With the switch off, nothing happens. No reaction occurs in either beaker and no current flows. The moment the switch closes the circuit, the cell springs to life:

At the zinc electrode (anode): Zinc atoms lose electrons and enter the solution as $Zn^{2+}$ ions. This is the oxidation half-reaction:

$Zn \rightarrow Zn^{2+} + 2e^-$

At the copper electrode (cathode): $Cu^{2+}$ ions from the solution pick up the electrons arriving through the wire and deposit as metallic copper on the rod. This is the reduction half-reaction:

$Cu^{2+} + 2e^- \rightarrow Cu$

Through the wire: Electrons travel from the zinc rod (anode) to the copper rod (cathode) through the external metallic wire. This flow of electrons is the electric current the cell produces.
Through the salt bridge: Inside the cell, current is carried not by electrons but by the migration of ions through the salt bridge. Anions move toward the anode compartment and cations move toward the cathode compartment, maintaining electrical neutrality in both solutions.
Current vs. electron flow: By convention, the direction of current is opposite to the direction of electron flow. Electrons flow from zinc to copper, so conventional current flows from copper to zinc through the external circuit.

The key takeaway is that the same redox reaction ( $Zn + Cu^{2+} \rightarrow Zn^{2+} + Cu$ ) is happening, but the electron transfer is now indirect, routed through the wire, allowing you to harness it as electricity.

Electrodes and Electrode Potential

The zinc rod and the copper rod are called electrodes (conductors through which electricity enters or leaves a solution). Each electrode has a characteristic tendency to either lose electrons (get oxidised) or gain electrons (get reduced). The potential associated with each electrode, arising from this tendency, is called the electrode potential.

Think of it this way: zinc “wants” to lose electrons more strongly than copper does. This difference in tendency creates a voltage (a potential difference) between the two electrodes, and that voltage is what pushes electrons through the wire.

Standard Electrode Potential: A Common Reference

Electrode potential depends on the concentration of ions in solution, the pressure of any gas involved, and the temperature. To compare electrodes on a level playing field, chemists define the standard electrode potential ( $E^\circ$ ): the electrode potential measured under the following standard conditions:

Every dissolved species is at a concentration of 1 M (unit activity)
Every gas involved is at a pressure of 1 atm
The temperature is 298 K (25°C)

Since you cannot measure the absolute potential of a single electrode in isolation (you always need a complete circuit with two electrodes), a reference point is needed. By international convention, the standard hydrogen electrode (SHE) is assigned a potential of exactly 0.00 V:

$2H^+ + 2e^- \rightarrow H_2(g) \qquad E^\circ = 0.00 \text{ V}$

All other standard electrode potentials are measured relative to this reference.

Reading the Standard Electrode Potential Table

The table below lists $E^\circ$ values for selected electrode processes, all written as reduction reactions (oxidised form + electrons $\rightarrow$ reduced form). This is the standard way the electrochemical series is presented.

Reduction half-reaction	$E^\circ$ / V
$F_2(g) + 2e^- \rightarrow 2F^-$	+2.87
$Co^{3+} + e^- \rightarrow Co^{2+}$	+1.81
$H_2O_2 + 2H^+ + 2e^- \rightarrow 2H_2O$	+1.78
$MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O$	+1.51
$Au^{3+} + 3e^- \rightarrow Au(s)$	+1.40
$Cl_2(g) + 2e^- \rightarrow 2Cl^-$	+1.36
$Cr_2O_7^{2-} + 14H^+ + 6e^- \rightarrow 2Cr^{3+} + 7H_2O$	+1.33
$O_2(g) + 4H^+ + 4e^- \rightarrow 2H_2O$	+1.23
$MnO_2(s) + 4H^+ + 2e^- \rightarrow Mn^{2+} + 2H_2O$	+1.23
$Br_2 + 2e^- \rightarrow 2Br^-$	+1.09
$NO_3^- + 4H^+ + 3e^- \rightarrow NO(g) + 2H_2O$	+0.97
$2Hg^{2+} + 2e^- \rightarrow Hg_2^{2+}$	+0.92
$Ag^+ + e^- \rightarrow Ag(s)$	+0.80
$Fe^{3+} + e^- \rightarrow Fe^{2+}$	+0.77
$O_2(g) + 2H^+ + 2e^- \rightarrow H_2O_2$	+0.68
$I_2(s) + 2e^- \rightarrow 2I^-$	+0.54
$Cu^+ + e^- \rightarrow Cu(s)$	+0.52
$Cu^{2+} + 2e^- \rightarrow Cu(s)$	+0.34
$AgCl(s) + e^- \rightarrow Ag(s) + Cl^-$	+0.22
$AgBr(s) + e^- \rightarrow Ag(s) + Br^-$	+0.10
$2H^+ + 2e^- \rightarrow H_2(g)$	0.00
$Pb^{2+} + 2e^- \rightarrow Pb(s)$	-0.13
$Sn^{2+} + 2e^- \rightarrow Sn(s)$	-0.14
$Ni^{2+} + 2e^- \rightarrow Ni(s)$	-0.25
$Fe^{2+} + 2e^- \rightarrow Fe(s)$	-0.44
$Cr^{3+} + 3e^- \rightarrow Cr(s)$	-0.74
$Zn^{2+} + 2e^- \rightarrow Zn(s)$	-0.76
$2H_2O + 2e^- \rightarrow H_2(g) + 2OH^-$	-0.83
$Al^{3+} + 3e^- \rightarrow Al(s)$	-1.66
$Mg^{2+} + 2e^- \rightarrow Mg(s)$	-2.36
$Na^+ + e^- \rightarrow Na(s)$	-2.71
$Ca^{2+} + 2e^- \rightarrow Ca(s)$	-2.87
$K^+ + e^- \rightarrow K(s)$	-2.93
$Li^+ + e^- \rightarrow Li(s)$	-3.05

How to Interpret $E^\circ$ Values

The sign and magnitude of a standard electrode potential tell you two important things about the redox couple:

Negative $E^\circ$ : Strong Reducing Agents

A negative $E^\circ$ means the reduced form of that couple gives up electrons more readily than $H_2$ does. In other words, the couple is a stronger reducing agent than the $H^+/H_2$ reference couple.

For example, zinc has $E^\circ = -0.76$ V. This tells you that zinc metal has a greater tendency to lose electrons and get oxidised than hydrogen gas does. That is why zinc can displace hydrogen from acids.

The more negative the $E^\circ$ , the stronger the reducing agent. Lithium ( $E^\circ = -3.05$ V) sits at the very bottom of the table and is the strongest reducing agent listed.

Positive $E^\circ$ : Strong Oxidising Agents

A positive $E^\circ$ means the oxidised form of that couple pulls in electrons more strongly than $H^+$ does. The couple is a weaker reducing agent (or equivalently, a stronger oxidising agent) than the $H^+/H_2$ couple.

For example, copper has $E^\circ = +0.34$ V. $Cu^{2+}$ ions have a stronger tendency to gain electrons than $H^+$ ions do, so copper cannot displace hydrogen from acids, but $Cu^{2+}$ can oxidise hydrogen gas.

The more positive the $E^\circ$ , the stronger the oxidising agent. Fluorine ( $E^\circ = +2.87$ V) sits at the very top of the table and is the most powerful oxidising agent listed.

Two Trends in the Table

When you read the standard electrode potential table from top to bottom (from the most positive to the most negative $E^\circ$ values):

Oxidising strength decreases. Species at the top (like $F_2$ , $MnO_4^-$ ) are the strongest oxidising agents because they have the greatest tendency to gain electrons.
Reducing strength increases. Species at the bottom (like $Li$ , $K$ , $Na$ ) are the strongest reducing agents because they have the greatest tendency to lose electrons.

This arrangement is the electrochemical series, and it connects directly to the competitive electron transfer experiments you studied earlier. Metals lower in the series can displace metals higher in the series from their salt solutions, precisely because the lower metal is a stronger reducing agent.

Connecting Everything: From the Beaker to the Battery

The Daniell cell is a beautiful demonstration of how chemistry produces electricity. Every battery you have ever used, from the one in your phone to the one starting a car, works on this same fundamental principle: separate the two halves of a redox reaction, force the electrons through an external circuit, and harvest the current.

You will explore electrode reactions, cell EMF calculations, and applications of electrochemistry in much greater detail in Class XII. For now, the key insight is this: redox reactions are not just chemical bookkeeping. They are the engine behind every electrochemical device, and the standard electrode potential table is the roadmap that tells you which reactions will run and how strongly they will push electrons through a wire.