Electrochemistry : Nernst Equation

Electrode Potential

(1) When a metal (M) is placed in a solution of its ions (M⁺⁺), either of the following three possibilities can occurs, according to the electrode potential solution pressure theory of Nernst.

(i) A metal ion Mⁿ⁺ collides with the electrode, and undergoes no change.

(ii) A metal ion Mⁿ⁺ collides with the electrode, gains n electrons and gets converted into a metal atom M, (i.e. the metal ion is reduced).

Mⁿ⁺(aq) + ne⁻ → M(s)

(iii) A metal atom on the electrode M may lose a electrons to the electrode, and enter to the solution as Mⁿ⁺, (i.e. the metal atom is oxidised).

M(s) → Mⁿ⁺(aq) + ne⁻

Thus, “the electrode potential is the tendency of an electrode to lose or gain electrons when it is in contact with solution of its own ions.”

(2) The magnitude of electrode potential depends on the following factors,

(i) Nature of the electrode,

(ii) Concentration of the ions in solution,

(iii) Temperature.

(3) Types of electrode potential : Depending on the nature of the metal electrode to lose or gain electrons, the electrode potential may be of two types,

(i) Oxidation potential : When electrode is negatively charged with respect to solution, i.e., it acts as anode. Oxidation occurs. M → Mⁿ⁺ + ne

(ii) Reduction potential : When electrode is positively charged with respect to solution, i.e. it acts as cathode. Reduction occurs. Mⁿ⁺ + ne⁻ → M

(4) Standard electrode potential : “If in the half cell, the metal rod (M) is suspended in a solution of one molar concentration, and the temperature is kept at 298 K, the electrode potential is called standard electrode potential, represented usually by E^o”. ‘or’

The standard electrode potential of a metal may be defined as “the potential difference in volts developed in a cell consisting of two electrodes, the pure metal in contact with a molar solution of one of its ions and the normal hydrogen electrode (NHE)”.

Standard oxidation potential for any half-cell = − (Standard reduction potential

Standard reduction potential for any half-cell = − (Standard reduction potential

(5) Reference electrode or reference half – cells

It is not possible to measure the absolute value of the single electrode potential directly. Only the difference in potential between two electrodes can be measured experimentally. It is, therefore, necessary to couple the electrode with another electrode whose potential is known. This electrode is termed as reference electrode or reference half – cells. Various types of half – cells have been used to make complete cell with spontaneous reaction in forward direction. These half – cells have been summarised in following table,

Various Types of Half – cells

Type	Example	Half – cell reaction	*Q =*	Reversible to	Electrode Potential (oxidn), E =
Gas ion half – cell	Pt(H₂)\|H⁺(aq) Pt(Cl₂)\|Cl⁻(aq)	$\[\frac{1}{2}{{H}_{2}}(g)\to {{H}^{+}}(aq)+{{e}^{-}}$ $\[C{{l}^{-}}(aq)\to \frac{1}{2}C{{l}_{2}}(g)+{{e}^{-}}$	[H⁺] $\[\frac{1}{[C{{l}^{-}}]}$	H+ Cl⁻	E^o – 0.0591 log[H⁺] E⁰ +0.0591 log[Cl⁻]
Metal – metal ion half – cell	Ag\|Ag⁺(aq)	Ag(s)→Ag⁺(aq) + e⁻	[Ag⁺]	Ag⁺	E⁰ – 0.0591 log[Ag⁺]
Metal insoluble salt anion half – cell	Ag, AgCl\|Cl⁻(aq)	Ag(s) + Cl⁻(aq)→AgCl(s) + e⁻	$\[\frac{1}{[C{{l}^{-}}]}$	Cl⁻	E^o + 0.0591 log[Cl⁻]
Calomel electrode	Hg, Hg₂Cl₂\|Cl⁻(aq)	2Hg(l) + 2Cl⁻→ Hg₂Cl₂(s) + 2e⁻	$\[\frac{1}{{{[C{{l}^{-}}]}^{2}}}$	Cl⁻	E^o + 0.0591 log[Cl⁻]
Metal – metal oxide hydroxide half – cell	Hg, HgO\|OH⁻(aq)	Hg(l) + 2OH⁻(aq) → HgO(s) + H₂O(l) + 2e⁻	$\[\frac{1}{{{[O{{H}^{-}}]}^{2}}}$	OH⁻	E^o + 0.0591 log[OH⁻]
Oxidation – reduction half – dell	$\[Pt\|\underset{(aq)}{\mathop{F{{e}^{2+}}}}\,,\,\,\,\underset{(aq)}{\mathop{F{{e}^{3+}}}}\,$	Fe²⁺(aq) → Fe³⁺(aq) + e⁻	$\[\frac{[F{{e}^{3+}}]}{[F{{e}^{2+}}]}$	Fe²⁺, Fe³⁺	$\[{{E}^{0}}-0.0591\,\,\,\log \frac{[F{{e}^{3+}}]}{[F{{e}^{2+}}]}$
Mercury –mercury sulphate half – cell	$\[Hg,\,\,\,HgS{{O}_{4}}\|SO_{4}^{2-}(aq)$	$\[SO_{4}^{2-}(aq)+Hg(l)\to HgS{{O}_{4}}(s)+2{{e}^{-}}$	$\[\frac{1}{[SO_{4}^{2-}]}$	$\[SO_{4}^{2-}$	$\[{{E}^{o}}+\frac{0.0591}{2}\log [SO_{4}^{2-}]$
Quinhydrone half – cell	Pt\|Quinhydrone\|H⁺(aq)		[H⁺]²	H⁺	E⁰ – 0.0591 log[H⁺]

Cell potential or EMF of the cell .

(1) “The difference in potentials of the two half – cells of a cell known as electromotive force (emf) of the cell or cell potential.”

The difference in potentials of the two half – cells of a cell arises due to the flow of electrons from anode to cathode and flow of current from cathode to anode.

Anode $\xrightleftharpoons[\text{Flow of current}]{\text{Flow of electrons}}$ Cathode

(2) The emf of the cell or cell potential can be calculated from the values of electrode potentials of two the half – cells constituting the cell. The following three methods are in use :

(i) When oxidation potential of anode and reduction potential of cathode are taken into account E^o_Cell= Oxidation potential of anode + Reduction potential of cathode $\[=E_{\text{ox}}^{0}\stackrel{\scriptscriptstyle\to}{\leftarrow}(\text{anode})+E_{\text{red}}^{\text{0}}(\text{cathode})$

(ii) When reduction potentials of both electrodes are taken into account E^o_Cell= Reduction Potential of cathode – Reduction potential of anode $\[=E_{\text{Cathode}}^{\text{0}}-E_{\text{Anode}}^{\text{0}}=E_{\text{right}}^{\text{o}}-E_{\text{left}}^{o}$

(iii) When oxidation potentials of both electrodes are taken into account E^o_Cell = Oxidation potential of anode – Oxidation potential of cathode $\[=E_{\text{ox}}^{0}(\text{anode})-E_{\text{ox}}^{0}(\text{cathode})$

(3) Difference between emf and potential difference : The potential difference is the difference between the electrode potentials of the two electrodes of the cell under any condition while emf is the potential generated by a cell when there is zero electron flow, i.e., it draws no current. The points of difference are given below

Emf	Potential difference
It is the potential difference between two electrodes when no current is flowing in the circuit.	It is the difference of the electrode potentials of the two electrodes when the cell is under operation.
It is the maximum voltage that the cell can deliver.	It is always less then the maximum value of voltage which the cell can deliver.
It is responsible for the steady flow of current in the cell.	It is not responsible for the steady flow of current in the cell.

(4) Cell EMF and the spontaneity of the reaction :

We know, ΔG = −nFE_Cell

(i) For a spontaneous process, DG is negative. Then, according to the equation for a spontaneous process, E_cell should be positive. Thus, the cell reaction will be spontaneous when the cell emf is positive.

(ii) For a non – spontaneous process, DG is positive. Then, according to equation for a non – spontaneous process, E_cell should be negative. Thus, the cell reaction will be non – spontaneous when the cell emf is negative.

(iii) For the process to be at equilibrium, ΔG = 0. Then, according to the equation E_cell should be zero. Thus, the cell reaction will be at equilibrium when the cell emf is zero. These results are summarized below

Nature of reaction	ΔG(or ΔG^o)	E_Cell (or E^o_Cell)
Spontaneous	–	+
Equilibrium	0	0
Non – spontaneous	+	–

Nernst’s equation.

(1) Nernst’s equation for electrode potential

The potential of the electrode at which the reaction,

Mⁿ⁺ + ne⁻ → M(s) takes place is described by the equation,

$\[{{E}_{{{M}^{n+}}/M}}=E_{{{M}^{n+}}/M}^{0}-\frac{RT}{nF}\ln \frac{[M(s)]}{[{{M}^{n+}}(aq.)]}$ or $\[{{E}_{{{M}^{n+}}/M}}=E_{{{M}^{n+}}/M}^{0}-\frac{2.303\,\,RT}{nF}\log \frac{[M(s)]}{[{{M}^{n+}}(aq)]}$ above eq. is called the Nernst equation.

Where, $\[{{E}_{{{M}^{n+}}/M}}$ = the potential of the electrode at a given concentration,

$\[{{E}_{{{M}^{n+}}/M}}$ = the standard electrode potential

R = the universal gas constant, 8.31 J K⁻¹ mol⁻¹,

T = the temperature on the absolute scale,

n = the number of electrons involved in the electrode reaction,

F = the Faraday constant : (96500 C),

[M(s)] = the concentration of the deposited metal,

[Mⁿ⁺(aq)] = the molar concentration of the metal ion in the solution,

The concentration of pure metal M(s) is taken as unity. So, the Nernst equation for the Mⁿ⁺/M electrode is written as,

$\[{{E}_{{{M}^{n+}}/M}}=E_{{{M}^{n+}}/M}^{0}-\frac{2.303\,\,RT}{nF}\log \frac{1}{[{{M}^{n+}}(aq)]}$

At 298 K, the Nernst equation for the electrode can be written as,

$\[{{E}_{{{M}^{n+}}/M}}=E_{{{M}^{n+}}/M}^{0}-\frac{0.0591}{n}\log \frac{1}{[{{M}^{n+}}(aq)]}$

For an electrode (half – cell) corresponding to the electrode reaction,

Oxidised form + ne⁻ → Reduced form

The Nernst equation for the electrode is written as,

$\[{{E}_{half-cell}}=E_{half-cell}^{0}-\frac{2.303\,\,RT}{nF}\log \frac{[\text{Reduced}\,\,\text{form }]}{\text{ }\!\![\!\!\text{ Oxidised}\,\,\text{form }\!\!]\!\!\text{ }}$

At 298 K, the Nernst equation can be written as,

$\[{{E}_{half-cell}}=E_{half-cell}^{0}-\frac{0.0591}{n}\log \frac{[\text{Reduced}\,\,\text{form }]}{\text{ }\!\![\!\!\text{ Oxidised}\,\,\text{form }\!\!]\!\!\text{ }}$

(2) Nernst’s equation for cell EMF

For a cell in which the net cell reaction involving n electrons is,

aA + bB → cC + dD

The Nernst equation is written as,

$\[{{E}_{cell}}=E_{cell}^{0}-\frac{RT}{nF}\text{ln}\frac{{{\text{ }\!\![\!\!\text{ C }\!\!]\!\!\text{ }}^{\text{c}}}{{[D]}^{d}}}{{{[A]}^{a}}{{[B]}^{b}}}$

Where, $\[E_{cell}^{0}=E_{cathode}^{0}-E_{anode}^{0}$ . The $\[E_{cell}^{o}$ is called the standard cell potential.

or $\[{{E}_{\text{cell}}}=E_{cell}^{o}-\frac{2.303\,\,RT}{nF}\log \frac{{{[C]}^{c}}{{[D]}^{d}}}{{{[A]}^{a}}{{[B]}^{b}}}$

At 298 K, above eq. can be written as,

or $\[{{E}_{\text{cell}}}=E_{cell}^{o}-\frac{0.0592}{n}\log \frac{{{[C]}^{c}}{{[D]}^{d}}}{{{[A]}^{a}}{{[B]}^{b}}}$

It may be noted here, that the concentrations of A, B, C and D referred in the eqs. are the concentrations at the time the cell emf is measured.

(3) Nernst’s equation for Daniells cell : Daniell’s cell consists of zinc and copper electrodes. The electrode reactions in Daniell’s cell are,

At anode : Zn(s) → Zn²⁺(aq) + 2e⁻

At cathode : Cu²⁺(aq) + 2e⁻ → Cu(s)

Net cell reaction : Zn + Cu²⁺(aq) → Cu(s) + Zn²⁺(aq)

Therefore, the Nernst equation for the Daniell’s cell is,

$\[{{E}_{cdll}}=E_{cell}^{0}-\frac{2.303\,\,RT}{2F}\log \frac{[Cu(s)][Z{{n}^{2+}}(aq)]}{[Zn(s)][C{{u}^{2+}}(aq)]}$

Since, the activities of pure copper and zinc metals are taken as unity, hence the Nernst equation for the Daniell’s cell is,

$\[{{E}_{cell}}=E_{cell}^{0}-\frac{2.303\,\,RT}{2F}\log \frac{[Z{{n}^{2+}}(aq]}{[C{{u}^{2+}}(aq)]}$

The above eq. at 298 K is,

$\[{{E}_{cell}}=E_{cell}^{o}-\frac{0.0591}{2}\log \frac{[Z{{n}^{2+}}(aq]}{[C{{u}^{2+}}(aq)]}V$

For Daniells cell, $\[E_{cell}^{0}=1.1\,\,V$

(4) Nernst’s equation and equilibrium constant

For a cell, in which the net cell reaction involing n electrons is,

aA + bB → cC + dD

The Nernst equation is

$\[{{E}_{Cell}}=E_{cell}^{0}-\frac{RT}{nF}\,\,\ln \,\,\,\frac{{{[C]}^{c}}{{[D]}^{d}}}{{{[A]}^{a}}{{[B]}^{b}}}$ …..(i)

At equailibrium, the cell cannot perform any useful work. So at equilibrium, E_cell is zero. Also at equilibrium, the ratio

$\[\frac{{{[C]}^{c}}{{[D]}^{d}}}{{{[A]}^{a}}{{[B]}^{b}}}={{\left[ \frac{{{[C]}^{c}}{{[D]}^{d}}}{{{[A]}^{a}}{{[B]}^{b}}} \right]}_{equil}}={{K}_{c}}$

Where, K_c is the equilibrium constant for the cell reaction. Then, at equilibrium eq. ….(i) becomes,

$\[0=E_{cell}^{0}-\frac{RT}{nF}\ln \,\,\,{{K}_{c}}$ …..(ii)

or $\[{{K}_{c}}=\exp \,\,(nFE_{cell}^{0}/RT)$

or In $\[{{K}_{c}}=\frac{nFE_{cell}^{0}}{RT}$ …..(iii)

At 298 K, eq. (ii) can be written as,

$\[0=E_{cell}^{0}-\frac{0.0592}{n}\log \,\,{{K}_{c}}$

This gives, log $\[{{K}_{c}}=(nE_{cell}^{0}/0.0592)$

$\[{{K}_{c}}={{10}^{(nE_{cell}^{o}/0.0592)}}$ ….. (iv)

eq. (iii) and (iv) may be used for obtaining the value of the equilibrium constant from the $\[E_{cell}^{o}$ value.

Relationship between cell potential, Gibbs Energy and Equilibrium constant .

The electrical work (electrical energy) is equal to the product of the EMF of the cell and electrical charge that flows through the external circuit i.e.,

W_max = nFE_cell ……(i)

According to thermodynamics the free energy change (ΔG) is equal to the maximum work. In the cell work is done on the surroundings by which electrical energy flows through the external circuit, So

−W_max = Δ G ……(ii)

from eq. (i) and (ii) ΔG = − nFE_cell

In standard conditions ΔG^o = − nFE^o_cellWhereΔG^o= standard free energy change But $\[E_{cell}^{0}=\frac{2.303}{nF}RT\,\,\,\log {{K}_{c}}$

$\[\Delta {{G}^{0}}=-nF\times \frac{2.303}{nF}RT\,\,\,\log \,\,\,{{K}_{c}}$

ΔG^o = − 2.303 RT log K_c

ΔG^o = − RT In K_c

(2.303 log X = In X)

Electrochemical series .

(1) The standard reduction potentials of a large number of electrodes have been measured using standard hydrogen electrode as the reference electrode. These various electrodes can be arranged in increasing or decreasing order of their reduction potentials. The arrangement of elements in order of increasing reduction potential values is called electrochemical series.

The electrochemical series, also called activity series, of some typical electrodes is being given in Table.

Standard reduction electrode potentials at 298K

Element	Electrode Reaction (Reduction)	Standard Electrode Reduction potential E⁰, volt
Li	Li⁺+ e^– = Li	–3.05
K	K⁺+ e^– = K	–2.925
Ca	Ca²⁺ + 2e⁻ = Ca	–2.87
Na	Na⁺ + e⁻ = Na	–2.714
Mg	Mg²⁺ + 2e⁻ = Mg	–2.37
Al	Al³⁺ + 3e⁻ = Al	–1.66
Zn	Zn²⁺+ 2e^–=Zn	–0.7628
Cr	Cr³⁺+ 3 e^– = Cr	–0.74
Fe	Fe²⁺+ 2e^– = Fe	–0.44
Cd	Cd²⁺+ 2e^– = Cd	–0.403
Ni	Ni²⁺+ 2e^– = Ni	–0.25
Sn	Sn²⁺+ 2e^– = Sn	–0.14
H₂	2H⁺+ 2e^– = H₂	0.00
Cu	Cu²⁺+ 2e^– = Cu	+0.337
I₂	I₂+ 2e^– = 2I^–	+0.535
Ag	Ag⁺+ e^– = Ag	+0.799
Hg	Hg²⁺+ 2e^– = Hg	+0.885
Br₂	Br₂+ 2e^– = 2Br^–	+1.08
Cl₂	Cl₂+ 2e^– = 2Cl^–	+1.36
Au	Au³⁺+ 3e^– = Au	+1.50
F₂	F₂+ 2e ^–= 2F^–	+2.87

(2) Characteristics of Electrochemical series

(i) The negative sign of standard reduction potential indicates that an electrode when joined with SHE acts as anode and oxidation occurs on this electrode. For example, standard reduction potential of zinc is –0.76 volt, When zinc electrode is joined with SHE, it acts as anode (–ve electrode) i.e., oxidation occurs on this electrode. Similarly, the +ve sign of standard reduction potential indicates that the electrode when joined with SHE acts as cathode and reduction occurs on this electrode.

(ii) The substances, which are stronger reducing agents than hydrogen are placed above hydrogen in the series and have negative values of standard reduction potentials. All those substances which have positive values of reduction potentials and placed below hydrogen in the series are weaker reducing agents than hydrogen.

(iii) The substances, which are stronger oxidising agents than ion are placed below hydrogen in the series.

(iv) The metals on the top (having high negative value of standard reduction potentials) have the tendency to lose electrons readily. These are active metals. The activity of metals decreases form top to bottom. The non-metals on the bottom (having high positive values of standard reduction potentials) have the tendency to accept electrons readily. These are active non-metals. The activity of non-metals increases from top to bottom.

(3) Application of Electrochemical series

(i) Reactivity of metals: The activity of the metal depends on its tendency to lose electron or electrons, i.e., tendency to form cation (Mⁿ⁺). This tendency depends on the magnitude of standard reduction potential. The metal which has high negative value (or smaller positive value) of standard reduction potential readily loses the electron or electrons and is converted into cation. Such a metal is said to be chemically active. The chemical reactivity of metals decreases from top to bottom in the series. The metal higher in the series is more active than the metal lower in the series. For example,

(a) Alkali metals and alkaline earth metals having high negative values of standard reduction potentials are chemically active. These react with cold water and evolve hydrogen. These readily dissolve in acids forming corresponding salts and combine with those substances which accept electrons.

(b) Metals like Fe, Pb, Sn, Ni, Co, etc., which lie a little down in the series do not react with cold water but react with steam to evolve hydrogen.

(c) Metals like Cu, Ag and Au which lie below hydrogen are less reactive and do not evolve hydrogen from water.

(ii) Electropositive character of metals : The electropositive character also depends on the tendency to lose electron or electrons. Like reactivity, the electropositive character of metals decreases from top to bottom in the electrochemical series. On the basis of standard reduction potential values, metals are divided into three groups

(a) Strongly electropositive metals : Metals having standard reduction potential near about – 2.0 volt or more negative like alkali metals, alkaline earth metals are strongly electropositive in nature.

(b) Moderately electropositive metals : Metals having values of reduction potentials between 0.0 and about – 2.0 volt are moderately electropositive Al, Zn, Fe, Ni, Co, etc., belong to this group.

(c) Weakly electropositive : The metals which are below hydrogen and possess positive values of reduction potentials are weakly electropositive metals. Cu, Hg, Ag, etc., belong to this group.

(iii) Displacement reactions

(a) To predict whether a given metal will displace another, from its salt solution: A metal higher in the series will displace the metal from its solution which is lower in the series, i.e., The metal having low standard reduction potential will displace the metal from its salt’s solution which has higher value of standard reduction potential. A metal higher in the series has greater tendency to provide electrons to the cations of the metal to be precipitated.

(b) Displacement of one nonmetal from its salt solution by another nonmetal: A non-metal higher in the series (towards bottom side), i.e., having high value of reduction potential will displace another non-metal with lower reduction potential, i.e., occupying position above in the series. The non-metal’s which possess high positive reduction potentials have the tendency to accept electrons readily. These electrons are provided by the ions of the nonmetal having low value of reduction potential,. Thus, Cl₂ can displace bromine and iodine from bromides and iodides.

Cl₂ + 2KI → 2KCl + I₂

2I⁻ → I₂ + 2e⁻ …..(Oxidation)

Cl₂ + 2e⁻ → 2Cl⁻ …..(Reduction)

[The activity or electronegative character or oxidising nature of the nonmetal increases as the value of reduction potential increases.]

(c) Displacement of hydrogen from dilute acids by metals : The metal which can provide electrons to H⁺ ions present in dilute acids for reduction, evolve hydrogen from dilute acids.

Mn → Mnⁿ⁺ + ne⁻ …..(Oxidation)

2H⁺ + 2e⁻ → H₂ …..(Reduction)

The metal having negative values of reduction potential possess the property of losing electron or electrons.

Thus, the metals occupying top positions in the electrochemical series readily liberate hydrogen from dilute acids and on descending in the series tendency to liberate hydrogen gas from dilute acids decreases.

The metals which are below hydrogen in electrochemical series like Cu, Hg, Au, Pt, etc., do not evolve hydrogen from dilute acids.

(d) Displacement of hydrogen from water : Iron and the metals above iron are capable of liberting hydrogen from water. The tendency decreases from top to bottom in electrochemical series. Alkali and alkaline earth metals liberate hydrogen from cold water but Mg, Zn and Fe liberate hydrogen from hot water or steam.

(iv) Reducing power of metals: Reducing nature depends on the tendency of losing electron or electrons. More the negative reduction potential, more is the tendency to lose electron or electrons. Thus reducing nature decreases from top to bottom in the electrochemical series. The power of the reducing agent increases, as the standard reduction potential becomes more and more negative. Sodium is a stronger reducing agent than zinc and zinc is a stronger reducing agent than iron. (decreasing order of reducing netur)

Element : Na > Zn > Fe

Recution potential : − 2.71 > − 0.76 > − 0.4

Alkali and alkaline earth metals are strong reducing agents.

(v) Oxidising nature of non-metals : Oxidising nature depends on the tendency to accept electron or electrons. More the value of reduction potential, higher is the tendency to accept electron or electrons. Thus, oxidising nature increases form top to bottom in the electrochemical series. The strength of an oxidising agent increases as the value of reduction potential becomes more and more positive.

F₂ (Fluorine) is a stronger oxidant than Cl₂, Br₂ and I₂. Cl₂ (Chlorine) is a stronger oxidant than Br₂ and I₂

Element : I₂ Br₂ Cl₂ F₂

Reduction potential : + 0.53 + 1.06 + 1.36 + 2.85

Oxidising nature increases

Thus, in electrochemical series

(vi) Thermal stability of metallic oxides : The thermal stability of the metal oxide depends on its electropositive nature. As the electropositivity decreases form top to bottom, the thermal stability of the oxide also decreases from top to bottom. The oxides of metals having high positive reduction potentials are not stable towards heat. The metals which come below copper form unstable oxides, i.e., these are decomposed on heating.

(vii) Extraction of metals : A more electropositive metal can displace a less electropositive metal from its salt’s solution. This principle is applied for the extraction of Ag and Au by cyanide process. silver from the solution containing sodium argento cyanide, NaAg(CN)₂, can be obtained by the addition of zinc as it is more electro-positive than

2NaAg(CN)₂ + Zn → Na₂Zn(CN)₄ + 2Ag

Electrochemistry : Nernst Equation

Electrochemistry : Nernst Equation

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