Electrochemical Cells

What an Electrochemical Cell Is

Electrochemistry, as NCERT frames it, is the study of the production of electricity from the energy released during spontaneous chemical reactions and, conversely, the use of electrical energy to bring about non-spontaneous chemical transformations. The apparatus in which either of these conversions is carried out in a controlled way is an electrochemical cell. Every such cell contains two electrodes in contact with one or more electrolytes; at each electrode an oxidation or a reduction half-reaction occurs, and the two are coupled through an external wire and an internal ionic pathway.

The defining feature of a cell is that it physically separates the two halves of a redox reaction. In an ordinary beaker, when a zinc rod is dipped into copper sulphate, the electrons pass directly from zinc to the copper ions and the energy is simply dissipated as heat. A cell instead forces those electrons to travel through an external circuit, so that the same chemical change can be made to deliver — or be driven by — useful electrical work.

The single redox reaction NCERT keeps returning to is the displacement of copper by zinc:

$\ce{Zn(s) + Cu^{2+}(aq) -> Zn^{2+}(aq) + Cu(s)}$

Galvanic vs Electrolytic Cells

Electrochemical cells split into two families on the basis of the direction of energy conversion. A galvanic (voltaic) cell converts the chemical energy of a spontaneous redox reaction into electrical energy; the Gibbs energy released by the reaction appears as electrical work that can run a motor, a heater or any other gadget. An electrolytic cell does the opposite — it uses electrical energy supplied by an external source to drive a chemical reaction that would not occur on its own.

In both families the electrode where oxidation occurs is named the anode and the electrode where reduction occurs is the cathode — that naming never changes. What flips between the two families is the sign of each electrode, and NCERT stresses this because it is a frequent source of error. The most striking point in the NCERT treatment is that galvanic and electrolytic behaviour are not the property of two different devices but two regimes of the very same Daniell cell, separated by the applied voltage.

The Daniell Cell as the Model Galvanic Cell

The Daniell cell is the canonical galvanic cell of the NCERT syllabus. It is built from a zinc rod dipping in a zinc sulphate solution and a copper rod dipping in a copper sulphate solution, the two beakers joined by a salt bridge and the two metals connected externally through a voltmeter and a switch. When the concentrations of $\ce{Zn^{2+}}$ and $\ce{Cu^{2+}}$ are both unity (1 mol dm$^{-3}$), the cell develops a potential of 1.1 V.

The overall reaction is the sum of two half-reactions occurring in the two separate compartments — also called half-cells or redox couples:

Oxidation (anode): $\ce{Zn(s) -> Zn^{2+}(aq) + 2e^-}$
Reduction (cathode): $\ce{Cu^{2+}(aq) + 2e^- -> Cu(s)}$

Because the oxidation half-reaction occurs only on the zinc electrode and the reduction half-reaction only on the copper electrode, the electrons released by zinc are obliged to travel through the external wire to reach the copper ions. That externally routed electron flow is the electric current the cell delivers.

Half-Cells, Electrodes and Their Roles

Each half-cell consists of a metallic electrode dipped into an electrolyte. NCERT notes that one can build innumerable galvanic cells by combining different half-cells in the Daniell pattern. The two half-cells are connected externally by a metallic wire through a voltmeter and a switch, and internally through a salt bridge. In the special case where both electrodes dip into the same electrolyte solution, no salt bridge is required.

The behaviour at every electrode–electrolyte interface is a contest between two opposing tendencies: metal ions from the solution tend to deposit on the electrode, trying to make it positive, while metal atoms of the electrode tend to dissolve as ions and leave their electrons behind, trying to make it negative. At equilibrium a separation of charge results, and the electrode ends up positively or negatively charged relative to the solution depending on which tendency dominates.

Once the switch is closed, electrons flow from the negative electrode (anode) to the positive electrode (cathode) through the external circuit, and the conventional current flows in the opposite direction — from cathode to anode externally. This direction convention is worth committing to memory, because the terms "from" and "to" reverse depending on whether one is tracking electrons or conventional current.

The Salt Bridge

A salt bridge is an inverted U-tube filled with a concentrated solution of an inert electrolyte such as $\ce{KCl}$ or $\ce{NH4NO3}$, set in a gel of agar-agar so that it does not flow out during use. The electrolyte chosen is one that takes no part in the cell reaction. According to NIOS, the salt bridge has two functions, and both are essential to a working galvanic cell.

The neutrality role is the easy one to overlook. As the cell runs, $\ce{Zn^{2+}}$ ions build up in the anode compartment (excess positive charge) while $\ce{Cu^{2+}}$ ions are consumed at the cathode (excess negative charge from the spectator $\ce{SO4^{2-}}$). If this imbalance were left unchecked, charge separation would oppose the cell reaction and the current would die almost at once. The salt bridge's ions migrate to cancel exactly these excesses, keeping the current flowing. NCERT adds that in the Daniell cell a salt bridge can be replaced by a porous pot serving the same purpose, making the cell more compact to handle.

Electrode Potential and Cell Potential

The potential difference that develops between an electrode and its electrolyte is the electrode potential. When the concentrations of all species in a half-cell are unity, this becomes the standard electrode potential. By IUPAC convention, standard reduction potentials are now taken to be the standard electrode potentials. In a galvanic cell the half-cell undergoing oxidation (the anode) has a negative potential relative to its solution, while the half-cell undergoing reduction (the cathode) has a positive potential.

The potential difference between the two electrodes is the cell potential, measured in volts. When no current is being drawn, this is the cell's electromotive force (emf). NCERT defines it as the reduction potential of the cathode minus that of the anode:

$E_\text{cell} = E_\text{right} - E_\text{left} = E_\text{cathode} - E_\text{anode}$

A single electrode's potential cannot be measured in isolation — only the difference between two half-cell potentials, which is the cell emf, is accessible. To assign absolute numbers, one electrode is taken as a reference: the standard hydrogen electrode, written $\ce{Pt(s) | H2(g) | H+(aq)}$, is assigned a potential of exactly zero at all temperatures, and every other electrode potential is quoted against it.

Cell Representation and Notation

Drawing a beaker diagram for every cell is impractical, so cells are written in a compact symbolic form. The convention, stated identically in NCERT and NIOS, places the anode on the left and the cathode on the right. A single vertical line marks a phase boundary between a metal and its electrolyte; a double vertical line marks the salt bridge connecting the two electrolytes.

For the copper–silver cell, NCERT gives a second worked representation. The spontaneous reaction is:

$\ce{Cu(s) + 2Ag+(aq) -> Cu^{2+}(aq) + 2Ag(s)}$

with half-reactions

Anode (oxidation): $\ce{Cu(s) -> Cu^{2+}(aq) + 2e^-}$
Cathode (reduction): $\ce{2Ag+(aq) + 2e^- -> 2Ag(s)}$

so the cell is written

$\ce{Cu(s) | Cu^{2+}(aq) || Ag+(aq) | Ag(s)}$

and its emf is $E_\text{cell} = E_{\ce{Ag+|Ag}} - E_{\ce{Cu^{2+}|Cu}}$. Reading a notation back into a chemistry is a routine NEET skill: the left half always oxidises, the right half always reduces, and the emf is the right-hand potential minus the left-hand one.

Energy Conversion and the 1.1 V Crossover

The most instructive experiment in the NCERT chapter is what happens when an opposing external voltage $E_\text{ext}$ is applied to a Daniell cell and slowly increased. The cell's own potential is 1.1 V, and the behaviour falls into three regimes that map cleanly onto the galvanic/electrolytic distinction.

Below 1.1 V the spontaneous reaction proceeds and the cell delivers electrical energy — it is galvanic. At exactly 1.1 V the opposing voltage just balances the cell potential, no current flows and no chemical change occurs. Above 1.1 V the external source overpowers the cell, reverses every flow, and forces the reaction to run backwards: now electrical energy is being used to drive a non-spontaneous reaction, and the device functions as an electrolytic cell. This single experiment shows that "galvanic" and "electrolytic" describe the operating mode of a cell, not two unrelated objects.

The thermodynamic bookkeeping behind the galvanic mode is that the Gibbs energy released by the spontaneous redox reaction is converted into electrical work. The deeper quantitative relation — connecting $E_\text{cell}$ with $\Delta_r G$ and the equilibrium constant — is taken up in the sibling notes on cell potential and the Nernst equation, but the qualitative point belongs here: a galvanic cell is, at heart, a controlled conversion of chemical free energy into usable electrical energy.