Chemistry · Redox Equilibrium · NEET
Standard electrode potential (E°) is a voltage that measures how strongly a species pulls electrons toward itself, so it gets reduced. Every E° in NEET tables is written as a reduction potential: the reaction is always written as oxidised form + electrons -> reduced form, for example Cu²⁺ + 2e⁻ -> Cu. A bigger (more positive) E° means the species grabs electrons easily and is easily reduced. 'Standard' means we measure it at fixed conditions: 298 K, 1 bar gas pressure, and 1 M concentration for ions.
We cannot measure the potential of a single electrode alone. So chemists picked one electrode as a reference and gave it the value 0 volts by agreement. That reference is the standard hydrogen electrode (SHE), based on the couple H⁺/H₂. Every other E° is really the voltage of a cell made with SHE on one side. So E° is a comparison, not an absolute number. The zero is a chosen starting line, like sea level for measuring height.
A positive E° means the reduced form is more stable than H₂, so the oxidised form easily takes electrons. That makes it a good oxidising agent (it oxidises others by taking their electrons). Example: F₂ and Cu²⁺ have high positive E°, so they are oxidising agents. A negative E° means the reduced form gives electrons easily, so it is a good reducing agent, like Na or K. Simple rule: more positive E° -> stronger oxidising agent; more negative E° -> stronger reducing agent.
A more negative E° means the metal loses electrons very easily and does not want them back. So that metal is a strong reducing agent and is more reactive (more electropositive). For example, E°(K⁺/K) = −2.93 V is very negative, so potassium is a very strong reducing agent. E°(Ag⁺/Ag) = +0.80 V is positive, so silver is a weak reducing agent. To order reducing power, list metals from most negative E° to most positive.
Electrode potential (E°) belongs to one half-cell (one couple like Zn²⁺/Zn). Cell potential (E°cell) belongs to the whole cell made of two half-cells. The formula is E°cell = E°cathode − E°anode. A positive E°cell means the reaction is spontaneous (it happens on its own). So one E° tells you the strength of a single couple; E°cell tells you if a full reaction between two couples will go forward.
The standard electrode potential (E°) values of Al³⁺/Al, Ag⁺/Ag, K⁺/K and Cr³⁺/Cr are −1.66 V, +0.80 V, −2.93 V and −0.74 V, respectively. The correct decreasing order of reducing power of the metals is:
Given the half-cell reactions: MnO₄⁻ + 8H⁺ + 5e⁻ -> Mn²⁺ + 4H₂O, E°(MnO₄⁻/Mn²⁺) = +1.510 V; ½O₂ + 2H⁺ + 2e⁻ -> H₂O, E°(O₂/H₂O) = +1.223 V. Will MnO₄⁻ liberate O₂ from water in the presence of acid?
The correct option for a redox couple is:
Try the real previous-year questions from this chapter — each with the answer and a full solution.
By convention, all standard electrode potentials in NEET and NCERT are written as standard reduction potentials. The half reaction is always written as oxidised form + electrons -> reduced form. If you ever need an oxidation potential, just flip the sign of the reduction potential.
Standard conditions are: temperature 298 K (25 °C), 1 bar pressure for any gas, and 1 M (unit activity) concentration for all ions in solution. The small circle in E° means these standard conditions are used.
Find E°cell = E°cathode − E°anode, where the cathode is the couple being reduced and the anode is the couple being oxidised. If E°cell is positive, the reaction is spontaneous (feasible). If it is negative, the reaction does not happen on its own.
It is a high-value shortcut. A single E° table lets you order reducing and oxidising power, predict displacement reactions, and check if a redox reaction is feasible, all without doing full experiments. NEET regularly asks direct questions on ordering reducing power and on E°cell feasibility.
