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Redox Reactions and Electrochemistry: one-page cheat sheet

NEET Chemistry · Unit 7 · Physical Chemistry

ScoreChemby CSJ Sir
Weightage: 3–4 questions · 12–16 marks

1 · Oxidation numbers

Oxidationloss of electrons; O.N. rises (reductant is oxidised)

Fixed valuesfree element 0; alkali metals

+ 1

; alkaline earths

+ 2

; F always

- 1

Oxygen

- 2

; peroxide

- 1

; superoxide

- \frac{1}{2}

; OF

_{2}

+ 2

Hydrogen

+ 1

; metal hydrides (NaH, CaH

_{2}

)

- 1

FractionalC

_{3}

_{2}

+ \frac{4}{3}

; Br

_{3}

_{8}

+ \frac{16}{3}

; S

_{4}

_{6}^{2 -}

+ 2.5

(average)

Disproportionsame element is oxidised and reduced: H

_{2}

_{2}

, or Cl

_{2}

in OH

^{-}

Other typescombination, decomposition, displacement

2 · Balancing & titrations

Acidic mediumsplit into half-reactions; balance O with H

_{2}

O, H with H

^{+}

, charge with

e^{-}

; equalise electrons

Basic mediumbalance as acidic, then add OH

^{-}

to both sides to neutralise the H

^{+}

Cr_{2} O_{7}^{2 -} + 14 H^{+} + 6 Fe^{2 +} \to 2 Cr^{3 +} + 6 Fe^{3 +} + 7 H_{2} O

2 MnO_{4}^{-} + 5 C_{2} O_{4}^{2 -} + 16 H^{+} \to 2 Mn^{2 +} + 10 CO_{2} + 8 H_{2} O

n

-factorMnO

_{4}^{-}

(acid) 5; Cr

_{2}

_{7}^{2 -}

6; C

_{2}

_{4}^{2 -}

2; Fe

^{2 +}

KMnO

_{4}

self-indicator; first permanent light pink at the end point

Iodometry

I_{2} + 2 S_{2} O_{3}^{2 -} \to 2 I^{-} + S_{4} O_{6}^{2 -}

; starch, blue to colourless

3 · Galvanic cells & Nernst

Daniell cellZn

∣

^{2 +} ∥

^{2 +} ∣

Cu; anode Zn (

-

), cathode Cu (

+

)

E_{cell}^{\circ} = E_{cathode}^{\circ} - E_{anode}^{\circ} = 0.34 - (- 0.76) = 1.10 V

E_{cell} = E_{cell}^{\circ} - \frac{0.0591}{n} lo g Q (298 K)

SHE

0.00

V at all temperatures: Pt

∣

_{2}

(1 bar)

∣

^{+}

(1 M)

E^{\circ}

(V)Ag

^{+}

/Ag

+ 0.80

; Cu

^{2 +}

/Cu

+ 0.34

; Fe

^{2 +}

/Fe

- 0.44

; Zn

^{2 +}

/Zn

- 0.76

Higher

E^{\circ}

stronger oxidising agent

4 · Cell thermodynamics

Δ_{r} G = - n F E_{cell}, Δ_{r} G^{\circ} = - n F E_{cell}^{\circ}

lo g K_{c} = \frac{n E _{cell}^{\circ}}{0.0591}, F = 96500 C mol^{- 1}

Spontaneous

E_{cell} > 0 \Rightarrow Δ_{r} G < 0

; at equilibrium

E_{cell} = 0

5 · Galvanic vs electrolytic

	Galvanic	Electrolytic
Energy	chemical to electrical	electrical to chemical
Reaction	spontaneous	non-spontaneous
Anode	negative	positive
Cathode	positive	negative

Oxidation is always at the anode, reduction at the cathode.

6 · Conductance

G = \frac{1}{R} = κ \frac{A}{l}, G^{*} = \frac{l}{A} = R κ (cell constant)

Λ_{m} = \frac{1000 κ}{C} (S cm^{2} mol^{- 1}), α = \frac{Λ _{m}}{Λ _{m}^{\circ}}

Strong

Λ_{m} = Λ_{m}^{\circ} - A c

; extrapolate to get

Λ_{m}^{\circ}

Weakrises steeply on dilution;

Λ_{m}^{\circ}

cannot be extrapolated

Kohlrausch

Λ_{m}^{\circ} = ν_{+} λ_{+}^{\circ} + ν_{-} λ_{-}^{\circ}

K_{a}

\frac{c α ^{2}}{1 - α} = \frac{c Λ _{m}^{2}}{Λ _{m}^{\circ} ( Λ _{m}^{\circ} - Λ _{m} )}

7 · Electrolysis

w = Z I t = \frac{M I t}{n \times 96500}, \frac{w _{1}}{w _{2}} = \frac{E _{1}}{E _{2}} (series)

Cathodespecies with the higher reduction potential discharges first

Anodespecies with the lower reduction potential oxidises first

Dilute H

_{2}

_{4}

with Pt electrodes: O

_{2}

at the anode, H

_{2}

at the cathode.

8 · Batteries & corrosion

Dry cellZn anode; graphite + MnO

_{2}

cathode; NH

_{4}

Cl + ZnCl

_{2}

paste; 1.5 V

Mercury cellZn–Hg anode; HgO + C cathode; KOH + ZnO; 1.35 V

Lead storagePb anode, PbO

_{2}

cathode, 38% H

_{2}

_{4}

(

d = 1.30

)

Pb + PbO_{2} + 2 H_{2} SO_{4} discharge 2 PbSO_{4} + 2 H_{2} O

Fuel cellH

_{2}

–O

_{2}

in alkaline medium; about 70% efficient; product water

Rustinganode: Fe oxidised to Fe

^{2 +}

(

- 0.44

V); cathode: O

_{2}

reduced (

+ 1.23

V). Rust is Fe

_{2}

_{3} \cdot x

_{2}

Protectiongalvanising (Zn) or Mg sacrificial blocks

Seen in NEET 2026

Mass of Cu from CuSO $_{4}$ with 1.5 A for 10 min (Cu = 63).

Q = 1.5 \times 600 = 900 C, w = \frac{63}{2 \times 96487} \times 900 = 0.2938 g

Seen in NEET 2021

$Λ^{\circ}$ of CH $_{3}$ COOH from NaCl (126.45), HCl (426.16), CH $_{3}$ COONa (91.0).

91.0 + 426.16 - 126.45 = 390.71 S cm^{2} mol^{- 1}

Six traps to remember

$E_{cell}^{\circ} = E_{cathode}^{\circ} - E_{anode}^{\circ}$ ; no sign change for stoichiometry.
$Δ G = - n F E$ : positive $E$ means spontaneous.
Cell constant $= R \times κ$ , in cm $^{- 1}$ .
Weak electrolytes need Kohlrausch's law for $Λ_{m}^{\circ}$ , not extrapolation.
Faraday: compare equivalent masses for cells in series.
$KMnO_{4}$ is a self-indicator; $K_{2} Cr_{2} O_{7}$ is not.

Redox Reactions and Electrochemistry: free one-page cheat sheet