Define the terms: (a) Electrochemical series (b) Corrosion. Write two applications of electrochemical series.

Question 94 of 111

Q. Define the terms:

(a) Electrochemical series

(b) Corrosion. Write two applications of electrochemical series.

Maharashtra MsbshseMaharashtra HSC (MSBSHSE) Board 2020Subjective· 4mImportance★★★★★

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Concept understanding — Electrochemical Series

The Intuition: A "Tug-of-War" for Electrons

Imagine a chemical reaction where two substances are fighting over electrons. One substance wants to give away electrons (get oxidised), the other wants to take electrons (get reduced). Who wins? That depends on how strongly each substance holds onto its electrons.

Some metals, like sodium, are extremely generous — they practically throw their electrons at anyone. Others, like gold, are miserly — they cling to their electrons and rarely let go.

The Electrochemical Series is simply a ranking of substances based on how badly they want to keep their electrons (or how eagerly they want to give them up). It's a leaderboard of electron greed.

The Precise Statement

The Electrochemical Series (also called the Activity Series or Reactivity Series) is a list of elements — mostly metals — arranged in order of their standard electrode potentials ( $E^\ominus$ ). These potentials are measured in volts, relative to the Standard Hydrogen Electrode (SHE), which is arbitrarily assigned a value of $0.00\ \text{V}$ .

E^\ominus_{\text{SHE}} = 0.00\ \text{V}

A more negative $E^\ominus$ means the element is a stronger reducing agent — it readily loses electrons and gets oxidised. A more positive $E^\ominus$ means the element is a stronger oxidising agent — it readily gains electrons and gets reduced.

Here is the series for some common metals and hydrogen:

Element	Half-Reaction (Reduction)	$E^\ominus$ (V)
Lithium	$\text{Li}^+ + e^- \rightarrow \text{Li}$	$-3.04$
Potassium	$\text{K}^+ + e^- \rightarrow \text{K}$	$-2.93$
Calcium	$\text{Ca}^{2+} + 2e^- \rightarrow \text{Ca}$	$-2.87$
Sodium	$\text{Na}^+ + e^- \rightarrow \text{Na}$	$-2.71$
Magnesium	$\text{Mg}^{2+} + 2e^- \rightarrow \text{Mg}$	$-2.37$
Aluminium	$\text{Al}^{3+} + 3e^- \rightarrow \text{Al}$	$-1.66$
Zinc	$\text{Zn}^{2+} + 2e^- \rightarrow \text{Zn}$	$-0.76$
Iron	$\text{Fe}^{2+} + 2e^- \rightarrow \text{Fe}$	$-0.44$
Tin	$\text{Sn}^{2+} + 2e^- \rightarrow \text{Sn}$	$-0.14$
Lead	$\text{Pb}^{2+} + 2e^- \rightarrow \text{Pb}$	$-0.13$
Hydrogen	$\text{2H}^+ + 2e^- \rightarrow \text{H}_2$	$0.00$
Copper	$\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$	$+0.34$
Silver	$\text{Ag}^+ + e^- \rightarrow \text{Ag}$	$+0.80$
Gold	$\text{Au}^{3+} + 3e^- \rightarrow \text{Au}$	$+1.50$

Watch out

The half-reactions are written as reductions (gaining electrons). A more negative $E^\ominus$ means the reverse reaction — oxidation (losing electrons) — is more favourable. So lithium, with $-3.04\ \text{V}$ , is the best at giving away electrons, not keeping them.

What the Series Tells You

1. Predicting Reactivity

The lower (more negative) a metal is in the series, the more reactive it is. Lithium, potassium, and sodium react violently with water. Gold and platinum sit at the bottom and do almost nothing.

2. Displacement Reactions

A metal higher in the series (more negative $E^\ominus$ ) can displace a metal lower in the series from its salt solution. For example, zinc can displace copper from copper sulphate:

\text{Zn}(s) + \text{CuSO}_4(aq) \rightarrow \text{ZnSO}_4(aq) + \text{Cu}(s)

Why? Zinc loses electrons more readily than copper. The zinc atoms donate electrons to the $\text{Cu}^{2+}$ ions, turning them into copper metal.

Tip

To check if a displacement happens: the metal doing the displacing must have a more negative $E^\ominus$ than the metal being displaced. If both are on the same side of hydrogen, the one with the more negative value wins.

3. Direction of Redox Reactions …

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The electrochemical series ranks electrode potentials; corrosion is a natural electrochemical process; the series predicts redox feasibility and metal reactivity order.

Electrochemical series: The arrangement of various electrodes (metal/metal-ion or non-metal/non-metal-ion couples) in the order of their increasing (or decreasing) standard reduction potentials ( $E^\circ$ ), measured relative to the standard hydrogen electrode (0.00 V), is called the electrochemical (or activity) series.
Corrosion: The spontaneous, gradual electrochemical degradation of a metal on exposure to environmental agents such as moisture, air (oxygen) and acidic gases, converting the metal into an oxide, hydroxide, carbonate or sulphide at its surface — e.g. rusting of iron (formation of hydrated $Fe_2O_3.xH_2O$ ). Applications of the electrochemical series:

(i) It helps predict the feasibility and direction of a redox reaction — a species with a higher (more positive) reduction potential can oxidise a species with a lower one. …

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Same / Similar Concept — real previous-year questions on the same or a closely similar concept, not this exact question.

CBSE 2026Set A1 markMCQ
$Q. The standard reduction potentials of metals A, B, C and D are -3.05, -1.66, -0.40 and +0.80 volt respectively. Which metal of the following would have the highest reducing power ?$
(a) A
(b) B
(c) C
(d) D
›Reveal solutionSolution

The more negative the standard reduction potential, the stronger the tendency to be oxidised, i.e. the greater the reducing power. A has -3.05 V, the most negative.

A species with a very negative reduction potential is easily oxidised (gives up electrons readily), so it is a strong reducing agent. …

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CBSE 2026Set ANNUAL1 markMCQ
$Q. Copper sulphate cannot be stored in zinc vessel. This is because$
(a) EMF of the reaction is negative
(b) EMF of the reaction is positive
(c) reduction potential value of Zn is more than Cu
(d) reduction potential value of Cu is negative
›Reveal solutionSolution

Zinc's much lower (more negative) standard reduction potential than copper's makes the displacement reaction's EMF positive, i.e. spontaneous — which is exactly why $Zn$ metal corrodes when it touches $CuSO_4$ solution.

Standard reduction potentials: $E^{\circ}(Cu^{2+}/Cu) = +0.34\ V$ ; $E^{\circ}(Zn^{2+}/Zn) = -0.76\ V$ . Since $Zn^{2+}/Zn$ has the far more negative (lower) reduction potential, $Zn$ metal is the stronger reducing agent and is oxidised preferentially:

$Zn(s) + Cu^{2+}(aq) \rightarrow Zn^{2+}(aq) + Cu(s)$

$E^{\circ}_{cell} = E^{\circ}_{cathode} - E^{\circ}_{anode} = 0.34-(-0.76) = +1.10\ V$

Since $\Delta G^{\circ} = -nFE^{\circ}_{cell}$ , a positive $E^{\circ}_{cell}$ gives a negative $\Delta G^{\circ}$ — the reaction is thermodynamically spontaneous. So $CuSO_4$ solution spontaneously reacts with a zinc container, dissolving the zinc (corroding it) and depositing copper metal on it.

Why the other options are wrong: …

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CBSE 2023Set F1 markMCQ
$Q. The standard electrode potentials of A, B, C and D metals are -3.05 V, -1.66 V, -0.40 V and +0.8 V respectively. Which of the following would have the highest reducing power?$
(a) A
(b) B
(c) C
(d) D
›Reveal solutionSolution

The more negative the standard reduction potential, the stronger the reducing agent — so A (-3.05 V) has the highest reducing power.

Reducing power measures the tendency of a metal to lose electrons (get oxidised). A more negative standard electrode (reduction) potential means the metal is more easily oxidised and is therefore a stronger reducing agent.

…

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CBSE 2021Set A1 markMCQ
$Q. The standard reduction potential values of elements A, B and C are + 0.68 V, -2.50 V and -0.50 V respectively. The order of their reducing power is$
(a) A > B > C
(b) A > C > B
(c) C > B > A
(d) B > C > A
›Reveal solutionSolution
Lower (more negative) standard reduction potential means a stronger tendency to be oxidised, i.e. a stronger reducing agent.

Reducing power measures how readily a species gives up electrons (undergoes oxidation). The more negative the standard reduction potential (E°), the greater the tendency to lose electrons, so the stronger the reducing agent.

Given:

A: E° = +0.68 V (least negative → weakest reducing agent) …
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CBSE 2020Set WA1 markMCQ
$+1.5\,V$
(a) $A > C > D > B$
(b) $B > D > C > A$
(c) $A > B > D > C$
(d) $D > A > B > C$
›Reveal solutionSolution

Lower (more negative) $E^\circ$ ⇒ stronger reducing agent ⇒ more reactive metal; so $B>D>C>A$ .

Concept: A metal's chemical activity (its tendency to lose electrons and get oxidised) increases as its standard reduction potential becomes more negative. Metals high in the activity series (like the alkali/alkaline-earth metals) have strongly negative $E^\circ$ .

…

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CBSE 2019Set ANNUAL1 markMCQ
$A^+/A$
(a) $A^+ + B^+ \longrightarrow A + B$
(b) $A^+ + B \longrightarrow A + B^+$
(c) $A + B^+ \longrightarrow A^+ + B$
(d) $A + B \longrightarrow A^+ + B^+$
›Reveal solutionSolution

$A$ has the more negative (lower) reduction potential, so it is the better reducing agent and is oxidised, while $B^+$ (higher reduction potential) is reduced.

…

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CBSE 2018Set ANNUAL1 markMCQ
$Q. The standard reduction potential of four alkali metals A, B, C and D are -3.06, -1.63, -0.40 and -0.80 volt respectively. Strongest reducing agent amongst them is$
(a) A
(b) B
(c) C
(d) D
›Reveal solutionSolution

Most negative reduction potential = strongest tendency to be oxidised = strongest reducing agent, i.e. A (-3.06 V).

A reducing agent works by giving up electrons (being oxidised). The more negative the standard reduction potential, the greater the tendency of that species to undergo oxidation rather than reduction, and hence the stronger it is as a reducing agent.

…

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