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Lesson 1Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series

├── The starting experimentzinc in copper sulphate

├── The idea behind the galvanic cell

├── Construction of the Daniell cell

├── What happens at each electrode

├── When does the current stop?
├── The salt bridge — what it is and why it is needed

├── Why we need a reference

├── The standard hydrogen electrode

├── Standard electrode potentials

├── Measuring an electrode potential (idea of the experiment)

├── Definition

├── What the series tells us (the conclusions to know)

├── Calculating the cell emf

└── Cell notation (cell diagram)

Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series · 1 / 16

The starting experiment: zinc in copper sulphate

Dip a clean zinc plate into blue copper(II) sulphate solution. After a short time:

a reddish-brown layer of copper metal appears on the zinc surface;
the zinc plate slowly dissolves (it gets thinner);
if you wait long enough, the blue colour of the solution fades until it becomes almost colourless, because Cu²⁺ ions (which are blue in water) have been used up and colourless Zn²⁺ ions have taken their place.

The overall change is

Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

This is a spontaneous redox reaction→it goes on by itself once the reactants meet. It can be split into two half-reactions:

OxidationZn(s) → Zn²⁺(aq) + 2e⁻. A zinc atom loses two electrons, becomes a zinc ion and goes into the solution.

ReductionCu²⁺(aq) + 2e⁻ → Cu(s). A copper(II) ion gains the two electrons that came from zinc and is deposited as a copper atom on the zinc surface.

Zinc is the reducing agent (it is oxidised); Cu²⁺ is the oxidising agent (it is reduced).
The number of electrons lost always equals the number gained.

Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series · 2 / 16

The idea behind the galvanic cell

In the beaker experiment the electrons jump directly from Zn atoms to Cu²⁺ ions on the surface, and the energy is released as heat.
Scientists realised that if the two half-reactions are carried out in separate containers connected in a controlled way, the electrons must travel through an external wire to get from the zinc to the copper ions.
That moving stream of electrons is an electric current.

So a galvanic (voltaic) cell is a device made of two half-cells, each in its own container, joined by a salt bridge, with the electrodes joined by a wire.
It converts the chemical energy of a spontaneous redox reaction into electrical energy.

There is a second kind of electric cell, in which electrical energy from an outside source (for instance a dry cell) is used to force a non-spontaneous redox reaction to occur.
It is called an electrolytic cell (or electrolysis cell).
We study both kinds in this chapter.

Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series · 3 / 16

Construction of the Daniell cell

The Daniell cell is the classic example of a galvanic cell. It has:

a zinc plate dipped in zinc sulphate solution (the zinc half-cell);
a copper plate dipped in copper(II) sulphate solution (the copper half-cell);

a salt bridge joining the two solutions;

a wire (with a voltmeter or a lamp) joining the two metal plates.
The solution around each electrode is called the electrolyte. When the plates are connected by a wire, current flows.

Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series · 4 / 16

What happens at each electrode

Anode(zinc plate, the negative pole): oxidation occurs.
Zn(s) → Zn²⁺(aq) + 2e⁻

Cathode(copper plate, the positive pole): reduction occurs.
Cu²⁺(aq) + 2e⁻ → Cu(s)

Overall cell reaction(add the two half-reactions, the electrons cancel):
Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Comparison aspect	Main idea	Keywords
Remember the rule that holds for every kind of cell	oxidation always takes place at the anode and reduction always takes place at the cathode. In a galvanic cell the anode is negative and the cathode is positive the electrons leave the anode run through the external wire and enter the cathode	oxidationalwaystakesplace
•	(Conventional current is drawn in the opposite direction from the copper plate to the zinc plate through the wire.)	Conventionalcurrentdrawnopposite
Notice what happens as time passes	the zinc plate loses mass (thinner) the copper plate gains mass (thicker) the concentration of Zn²⁺ rises and the concentration of Cu²⁺ falls so the blue colour of the copper half-cell fades	zincplatelosesmass

Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series · 5 / 16

When does the current stop?

The current stops when the cell reaction can no longer continue
that is when

all the zinc metal of the anode has dissolved, or
all the Cu²⁺ ions of the cathode half-cell have been changed into copper atoms.
At that point the cell is "dead" (exhausted).

Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series · 6 / 16

The salt bridge — what it is and why it is needed

A salt bridge is usually a U-shaped glass tube filled with a concentrated solution of a strong electrolyte whose ions do not react with the ions in either half-cell or with the electrode materials (for example sodium sulphate, Na₂SO₄, or potassium nitrate; in the school course Na₂SO₄ is the example).
Its jobs:

Comparison aspect	Main idea	Keywords
It completes the electric circuit	by connecting the two solutions so ions can move between the half-cells and the circuit is closed	connectingsolutionsionsmove
It keeps each half-cell electrically neutral.	In the zinc half-cell positive Zn²⁺ ions keep being produced so positive charge builds up. In the copper half-cell positive Cu²⁺ ions keep being removed so negative SO₄²⁻ ions are left in excess. Without the bridge these charge imbalances would quickly oppose the flow of electrons and stop the reaction. The anions of the bridge move toward the zinc half-cell and its cations move toward the copper half-cell cancelling the excess charge	zincpositiveionsbeing
It keeps the two solutions apart	so that Cu²⁺ ions do not reach the zinc plate directly (which would give the useless direct reaction of Section 1.1)	reachzincplatedirectly

If the salt bridge is removed
the oxidation and reduction stop
and the current in the external wire stops.

Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series · 7 / 16

Why we need a reference

When a metal is dipped in a solution of its own ions there is a tendency for the metal to give ions to the solution (or for ions to take electrons from the metal).
This creates a potential difference between the metal and the solution, called the electrode potential.
There is no direct way to measure the potential of a single electrode against its own solution.
What we can measure easily is the difference in potential between two electrodes, i.e.
the voltage (emf) of a cell built from them.

So we build a cell from the electrode we want to study and a reference electrode whose potential is fixed and known.
By measuring the cell’s emf, we calculate the unknown electrode potential.
By international agreement the reference is the standard hydrogen electrode (SHE), and its potential is defined as exactly zero volt.

Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series · 8 / 16

The standard hydrogen electrode

The SHE consists of:

a small sheet of platinum (about 1 cm²) covered with a layer of finely divided, spongy "platinum black" (this surface adsorbs hydrogen gas and also acts as an inert conductor);
hydrogen gas at a constant pressure of 1 atmosphere bubbling over the platinum;
the platinum dips into an acid solution in which the concentration of H⁺ is 1 mol/L.
Under these standard conditions the electrode is the standard hydrogen electrode, with E° = 0.00 V. The half-reaction is
2H⁺(aq) + 2e⁻ ⇌ H₂(g)

Because the electrode potential depends on the H⁺ concentration and on the pressure of hydrogen, the potential of a hydrogen electrode changes if you alter the acid concentration or the gas pressure (or both).
That is why both conditions (1 mol/L and 1 atm) must be fixed to call it "standard".

Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series · 9 / 16

Standard electrode potentials

The standard electrode potential of a half-cell is its potential measured against the SHE when all dissolved species are at 1 mol/L, gases are at 1 atm and (as usually stated) the temperature is 25 °C.

If the half-cell is written as a reduction (Mⁿ+ + ne⁻ → M), the value is the standard reduction potential.
If it is written as an oxidation (M → Mⁿ+ + ne⁻), the value is the standard oxidation potential.
The oxidation potential of an electrode equals its reduction potential with the sign reversed.
For example, for copper E°(reduction) = +0.34 V, so E°(oxidation) = −0.34 V. For zinc: E°(reduction) = −0.76 V, so E°(oxidation) = +0.76 V.

Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series · 10 / 16

Measuring an electrode potential (idea of the experiment)

To measure the potential of zinc, connect a zinc half-cell (Zn in 1 M Zn²⁺) to a standard hydrogen electrode with a salt bridge and a voltmeter.
The voltmeter reads 0.76 V; because zinc is the one that is oxidised (it is the negative electrode) its oxidation potential is +0.76 V (reduction potential −0.76 V).
Repeating with copper, the voltmeter reads 0.34 V but now the hydrogen electrode is the anode and copper is the cathode (positive), so copper’s reduction potential is +0.34 V.

Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series · 11 / 16

Definition

The electromotive (electrochemical) series is the arrangement of the elements in descending order of their standard oxidation potentials — or, equivalently, in ascending order of their standard reduction potentials — relative to the standard hydrogen electrode (zero).
An illustrative part of the series (reduction potentials, volts; values are those commonly used in this course and may differ by a few hundredths of a volt from other textbooks):

Half-reaction (reduction)	E° reduction (V)	E° oxidation (V)
Li⁺ + e⁻ → Li	−3.05	+3.05
K⁺ + e⁻ → K	−2.92	+2.92
Na⁺ + e⁻ → Na	−2.71	+2.71
Mg²⁺ + 2e⁻ → Mg	−2.37	+2.37
Al³⁺ + 3e⁻ → Al	−1.67	+1.67
Zn²⁺ + 2e⁻ → Zn	−0.76	+0.76
Cr³⁺ + 3e⁻ → Cr	−0.74	+0.74
Fe²⁺ + 2e⁻ → Fe	−0.41	+0.41
Cd²⁺ + 2e⁻ → Cd	−0.40	+0.40
Ni²⁺ + 2e⁻ → Ni	−0.23	+0.23
Pb²⁺ + 2e⁻ → Pb	−0.13	+0.13
2H⁺ + 2e⁻ → H₂	0.00	0.00
Cu²⁺ + 2e⁻ → Cu	+0.34	−0.34
Ag⁺ + e⁻ → Ag	+0.80	−0.80
Au³⁺ + 3e⁻ → Au	+1.42	−1.42
F₂ + 2e⁻ → 2F⁻	+2.87	−2.87

Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series · 12 / 16

What the series tells us (the conclusions to know)

The oxidation potential of an electrode equals its reduction potential with the opposite sign.
A descending order of oxidation potential is an ascending order of reduction potential.
Elements at the top of the series (high oxidation potential, e.g. Li, K, Mg) are good reducing agents — they lose electrons easily.
Elements/ions at the bottom of the series (high reduction potential, e.g. F₂, Au³⁺) are good oxidising agents — they gain electrons easily.
The descending order of oxidation potential is also the descending order of chemical reactivity of the metals.
An element higher in the series acts as the anode when combined with any element lower in the series.
A metal higher in the series displaces any lower metal from the solution of its salt (zinc displaces copper from CuSO₄; copper does not displace zinc from ZnSO₄).
Every metal above hydrogen in the series displaces hydrogen from dilute acids (hydrogen gas is evolved), while metals below hydrogen (Cu, Ag, Au) do not. Example: Fe + 2HCl → FeCl₂ + H₂↑ (iron is oxidised, hydrogen ions are reduced).

Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series · 13 / 16

Calculating the cell emf

The electromotive force (emf), also called the cell potential E°cell, is the voltage the cell can deliver.
For a working galvanic cell it is always positive.
It can be found in three equivalent ways:

emf = E°(reduction of cathode) − E°(reduction of anode)
emf = E°(oxidation of anode) − E°(oxidation of cathode)
emf = E°(oxidation of anode) + E°(reduction of cathode)

Worked example 1 — the Daniell cell.→E°red(Cu²⁺/Cu) = +0.34 V; E°red(Zn²⁺/Zn) = −0.76 V.

Method 10.34 − (−0.76) = 1.10 V.

Method 2E°ox(Zn) = +0.76 V and E°ox(Cu) = −0.34 V, so 0.76 − (−0.34) = 1.10 V.

Method 30.76 + 0.34 = 1.10 V. All three agree: emf = 1.10 V.

Worked example 2 — new cell (nickel/silver).Given E°red(Ni²⁺/Ni) = −0.23 V and E°red(Ag⁺/Ag) = +0.80 V.

Comparison aspect	Main idea	Keywords
•	The electrode with the lower reduction potential (nickel) is the anode silver is the cathode	electrodelowerreductionpotential
Anode	Ni → Ni²⁺ + 2e⁻. Cathode: 2Ag⁺ + 2e⁻ → 2Ag (the silver half-reaction is doubled to balance electrons but the potential is not doubled — potential is an intensive property)	Cathodesilverhalfreaction
Overall	Ni(s) + 2Ag⁺(aq) → Ni²⁺(aq) + 2Ag(s)
•	emf = 0.80 − (−0.23) = 1.03 V
Worked example 3 — will the reaction occur?	Can iron reduce Pb²⁺? E°red(Fe²⁺/Fe) = −0.41 V E°red(Pb²⁺/Pb) = −0.13 V. Taking Fe as anode: emf = −0.13 − (−0.41) = +0.28 V. A positive emf means the reaction Fe + Pb²⁺ → Fe²⁺ + Pb is spontaneous. If you had written the reverse (lead as anode) the emf would be −0.28 V which signals a non-spontaneous reaction	ironreduceTakinganode

Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series · 14 / 16

Cell notation (cell diagram)

A galvanic cell is shown by a short-hand diagram→anode on the left, cathode on the right, a single vertical line for each phase boundary and a double vertical line for the salt bridge:

Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

For a reaction such as Ni²⁺ + Fe → Ni + Fe²⁺ the iron is oxidised, so the diagram is Fe(s) | Fe²⁺(aq) || Ni²⁺(aq) | Ni(s): iron is the anode (negative), nickel is the cathode (positive), electrons flow in the wire from iron to nickel and conventional current flows from nickel to iron.
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Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series · 15 / 16

Part 16

📖

🔍 In-Depth Explanations, Deductions & Multi-Book Insights:

Chemistry Thanaweya Amma (2026/2027)

Lesson 1 – Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series
Detailed Academic Lesson Overview:
Electrochemistry explores interconversions between chemical energy and electrical energy via spontaneous or forced redox reactions.
Galvanic cells produce electricity spontaneously.
The Standard Hydrogen Electrode (SHE) establishes the zero baseline for standard oxidation and reduction potentials in the Electrochemical Series.

🌳 Hierarchy Classification & Conceptual Roadmap

Galvanic Cells and Electrode Potentials

├── 1. Galvanic Cell Fundamentals (Daniell Cell)

│ ├── Anode (Negative electrode) Oxidation occurs (Zn → Zn²⁺ + 2e⁻)
│ ├── Cathode (Positive electrode) Reduction occurs (Cu²⁺ + 2e⁻ → Cu)
│ ├── External wire Electrons flow spontaneously from Anode to Cathode
│ └── Salt Bridge (U-tube with inert electrolyte Na₂SO₄ or KNO₃):
│ ├── Anions (SO₄²⁻, NO₃⁻) migrate toward Anode half-cell
│ └── Cations (Na⁺, K⁺) migrate toward Cathode half-cell
├── 2. Standard Hydrogen Electrode (SHE: E° = 0.00 V)
│ ├── Platinum sheet covered with black spongy platinum, dipped in 1 M H⁺, with H₂ gas bubbled at 1 atm
│ └── E° changes from 0.00 V if [H⁺] ≠ 1 M or P_H2 ≠ 1 atm
└── 3. Electrochemical Series (Electromotive Series)
├── Top elements Highest standard oxidation potentials, most active, strong reducing agents
├── Bottom elements Lowest oxidation potentials (highest reduction), weak reducing agents
└── Cell Electromotive Force E°cell = E°ox(anode) + E°red(cathode) > 0 for spontaneous cell
📋 Master Teacher's Blackboard Rules & Exam Secrets
RULE 1 In ALL electrochemical cells (galvanic or electrolytic), OXIDATION always occurs at the ANODE, and REDUCTION always occurs at the CATHODE.
RULE 2 In galvanic cells: Anode is NEGATIVE (source of e⁻); Cathode is POSITIVE.
RULE 3 If calculated E°cell is POSITIVE, the redox reaction is spontaneous (Galvanic cell). If E°cell is NEGATIVE, the reaction is non-spontaneous and requires external electrical power (Electrolytic cell).
RULE 4 Any element higher in the electrochemical series displaces hydrogen from acids and displaces any metal below it from its aqueous salt solution.
📊 Detailed Comparison Matrix Daniell Galvanic Cell Component and Operational Architecture

Component / Phenomenon	Anode Half-Cell (Zn)	Cathode Half-Cell (Cu)
Electrode Material	Zinc plate (Zn)	Copper plate (Cu)
Electrode Sign / Polarity	Negative pole (-)	Positive pole (+)
Electrochemical Process	Oxidation (Zn → Zn²⁺ + 2e⁻)	Reduction (Cu²⁺ + 2e⁻ → Cu)
Mass Change of Electrode	Decreases (Zn dissolves into ions)	Increases (Cu metal deposits)
Solution Concentration	[Zn²⁺] increases	[Cu²⁺] decreases
Salt Bridge Ion Migration	Anions (SO₄²⁻, NO₃⁻) migrate here	Cations (Na⁺, K⁺) migrate here

Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series · 16 / 16

Curriculum

Chemistry

Chapter 1 - Transition Elements and Iron Metallurgy

Chapter 2 - Chemical Analysis

Chapter 3 - Chemical Equilibrium

Chapter 4 - Electrochemistry

Chapter 5 - Organic Chemistry

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Lesson 1: Galvanic Cells, Standard Hydrogen Electrode and Electrochemical Series