16th – 18th Century Discoveries#

• William Gilbert (16th Century): He spent lots of time studying magnets and also did some early work on electricity. He’s often called the “Father of Magnetism” for his efforts in producing and strengthening magnets.
• Otto von Guericke (1663): Built the first electric generator! It made static electricity using friction – basically rubbing a sulfur ball inside a glass globe as it spun.
• Charles François de Cisternay du Fay (Mid-18th Century): Found out there were two types of static electricity (he called them “vitreous” and “resinous”, we now call them positive and negative) and saw that same types repel, different types attract. This was the “two-fluid theory”.
• Benjamin Franklin (Later 18th Century): Proposed a simpler “one-fluid theory” of electricity, challenging Du Fay’s idea.
• Charles-Augustin de Coulomb (1785): Built on work by Joseph Priestley and developed the law describing how electric charges attract or repel each other – now known as Coulomb’s Law of electrostatic attraction.
• Luigi Galvani (Late 18th Century): This is a big one for electrochemistry! Galvani, a doctor, noticed that frog legs twitched when touched by different metals connected by a circuit. He thought there was a special “animal electricity” in the tissue.
• Alessandro Volta: Galvani’s idea sparked debate. Volta disagreed about “animal electricity,” believing the reaction came from the metals themselves interacting with the moisture. His experiments led him to invent the first practical battery, the voltaic pile, using zinc and copper. This device could produce a steady electric current for a much longer time than anything before it, showing that chemical reactions could indeed generate electricity.

19th Century Milestones#

The 1800s saw huge leaps, thanks in large part to Volta’s battery providing a steady current source.

• William Nicholson and Johann Wilhelm Ritter (1800): Used Volta’s battery to split water into hydrogen and oxygen, demonstrating electrolysis. Ritter also figured out electroplating and saw that the amount of material deposited depended on the electrode spacing.
• Johann Wilhelm Ritter (1801): Noticed thermoelectric currents, hinting at the connection between heat and electricity.
• Sir Humphry Davy (1810s): Used electrolysis extensively. His work suggested that electricity production in simple cells was due to chemical action between substances with opposite charges. He famously isolated metals like sodium and potassium by electrolyzing their molten salts.
• Hans Christian Ørsted (1820): Discovered that electric currents create magnetic fields, a foundational piece for electromagnetism.
• André-Marie Ampère: Quickly put Ørsted’s findings into mathematical laws.
• Thomas Johann Seebeck (1821): Showed that a temperature difference between junctions of different metals can create an electrical potential (the Seebeck effect, basis of thermocouples).
• Georg Ohm (1827): Published his famous law, connecting voltage, current, and resistance in an electric circuit (Ohm’s Law).
• Michael Faraday (1832): A giant in the field. His experiments led to his two laws of electrochemistry, which quantified the relationship between the amount of substance reacted/produced and the amount of electricity passed. He also coined many terms used today, like “electrode,” “electrolyte,” “anode,” and “cathode.”
• John Daniell (1836): Invented the Daniell cell, an improved battery that helped solve the issue of “polarization” (where the build-up of reaction products reduces the cell’s voltage over time) by using different electrolytes separated by a porous barrier or salt bridge.
• William Grove (1839): Created the first fuel cell, demonstrating that combining hydrogen and oxygen could directly produce electricity.
• Georges Leclanché (1868): Patented a cell that became the basis for the common zinc-carbon “dry cell,” making batteries more portable.
• Svante Arrhenius (1884): Proposed his theory of electrolytic dissociation, explaining that electrolytes dissolve in water into charged ions, which are responsible for conducting electricity in solutions.
• Paul Héroult and Charles M. Hall (1886): Independently developed the Hall–Héroult process, a major electrochemical method for producing aluminum from alumina using electrolysis.
• Walther Hermann Nernst (1888-1889): Developed the theory for the electromotive force (voltage) of voltaic cells. He showed how this voltage relates to the energy change in the chemical reaction and derived the Nernst equation, which describes how cell potential depends on the concentration of reactants and products.
• Fritz Haber (1898): Showed that specific products could be obtained in electrolysis by carefully controlling the voltage at the electrode (cathode).

20th Century and Beyond#

• The Electrochemical Society (ECS) (1902) and International Society of Electrochemistry (ISE) (1949): Foundations of major professional societies dedicated to the field.
• Robert Andrews Millikan and Harvey Fletcher (1909-1911): Famous for the oil drop experiment, which accurately measured the fundamental charge of a single electron.
• Brønsted and Lowry (1923): Published their acid-base theory, which has electrochemical underpinnings.
• Arne Tiselius (1937): Developed electrophoresis for separating proteins, leading to a Nobel Prize.
• Revaz Dogonadze (1960s-1970s): Developed quantum electrochemistry, applying quantum mechanics to understand electron transfer in electrochemical systems.

This history shows a gradual building of knowledge, from early observations of static electricity and “animal sparks” to understanding fundamental laws and developing critical industrial processes and power sources.

Oxidation and Reduction (Redox Reactions)#

The term “redox” is just a shortcut for reduction-oxidation. These are the essential chemical changes that happen in electrochemical processes.

Redox Reaction: A chemical reaction involving the transfer of electrons between chemical species, resulting in changes in their oxidation states.

Think of it like a trade: one chemical species gives away electrons, and another species takes them.

Oxidation State: A number assigned to an atom in a molecule or ion that indicates the degree of oxidation (loss of electrons) or reduction (gain of electrons) compared to its neutral state. It’s like a hypothetical charge if all bonds were purely ionic.

When an atom or ion loses electrons, its oxidation state becomes more positive (or less negative). This is called oxidation. When an atom or ion gains electrons, its oxidation state becomes more negative (or less positive). This is called reduction.

Here are some handy ways to remember this:

• OIL RIG: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).
• LEO says GER: Lose Electrons Oxidation, Gain Electrons Reduction.

Oxidation and reduction always happen together. You can’t have one without the other. If something is oxidized, something else must be reduced.

• The substance that loses electrons (gets oxidized) is called the reducing agent or reductant. It causes reduction in another species.
• The substance that gains electrons (gets reduced) is called the oxidizing agent or oxidant. It causes oxidation in another species.

So, the oxidizing agent is the one that gets reduced in the reaction, and the reducing agent is the one that gets oxidized.

You might hear “oxidation” and think of oxygen, but it’s not always involved. Oxidation just means losing electrons. Oxygen is a common oxidizing agent because it tends to gain electrons very easily. Fluorine is even stronger!

Sometimes, especially in organic chemistry, you can spot oxidation and reduction by looking at oxygen and hydrogen:

• Gaining oxygen often means oxidation (the other species got reduced).
• Losing hydrogen often means oxidation (because hydrogen usually carries an electron away).
• Losing oxygen often means reduction.
• Gaining hydrogen often means reduction.

Balancing Redox Reactions#

Working with electrochemical reactions means dealing with these redox changes. To make sense of them, we often break them down into two “half-reactions”: one for oxidation and one for reduction. Then we combine and balance them, making sure the number of atoms and the total charge are equal on both sides of the equation. For reactions in water, we often need to add water molecules (H₂O) and hydrogen ions (H⁺) in acidic solutions or hydroxide ions (OH⁻) in basic solutions, plus the electrons (e⁻) being transferred, to get everything balanced.

Here’s a peek at how you might balance these half-reactions in different conditions:

Acidic Medium: Add H₂O and H⁺ ions.

• Example: Manganese(II) ions reacting with sodium bismuthate.
  • Unbalanced: Mn²⁺(aq) + NaBiO₃(s) → Bi³⁺(aq) + MnO₄⁻(aq)
  • Oxidation (Mn²⁺ loses electrons): Add water to balance oxygens, then H⁺ to balance hydrogens, then electrons to balance charge. 4 H₂O(l) + Mn²⁺(aq) → MnO₄⁻(aq) + 8 H⁺(aq) + 5 e⁻
  • Reduction (BiO₃⁻ gains electrons): Add H⁺ to balance charge and hydrogens, then water to balance oxygens, then electrons. (Note: Na⁺ is a “spectator ion” here, not directly involved in the redox). 2 e⁻ + 6 H⁺(aq) + BiO₃⁻(s) → Bi³⁺(aq) + 3 H₂O(l)
  • To combine, make the number of electrons equal (multiply the oxidation half by 2, reduction half by 5). 8 H₂O + 2 Mn²⁺ → 2 MnO₄⁻ + 16 H⁺ + 10 e⁻ 10 e⁻ + 30 H⁺ + 5 BiO₃⁻ → 5 Bi³⁺ + 15 H₂O
  • Add them up, cancelling out common terms (like electrons, some H⁺, some H₂O): 14 H⁺(aq) + 2 Mn²⁺(aq) + 5 BiO₃⁻(s) → 7 H₂O(l) + 2 MnO₄⁻(aq) + 5 Bi³⁺(aq) (If you bring back the Na⁺ spectator ions, you get the full equation from the original text).

Basic Medium: Add H₂O and OH⁻ ions.

• Example: Potassium permanganate and sodium sulfite.
  • Unbalanced: KMnO₄ + Na₂SO₃ + H₂O → MnO₂ + Na₂SO₄ + KOH
  • Consider just the ions/molecules changing: MnO₄⁻, SO₃²⁻, MnO₂, SO₄²⁻. (K⁺, Na⁺, OH⁻ are spectators or part of the environment).
  • Reduction (MnO₄⁻ gains electrons): Add water/OH⁻ to balance. 3 e⁻ + 2 H₂O + MnO₄⁻ → MnO₂ + 4 OH⁻
  • Oxidation (SO₃²⁻ loses electrons): Add water/OH⁻ to balance. 2 OH⁻ + SO₃²⁻ → SO₄²⁻ + H₂O + 2 e⁻
  • Combine by making electrons equal (multiply reduction by 2, oxidation by 3). 6 e⁻ + 4 H₂O + 2 MnO₄⁻ → 2 MnO₂ + 8 OH⁻ 6 OH⁻ + 3 SO₃²⁻ → 3 SO₄²⁻ + 3 H₂O + 6 e⁻
  • Add them up, cancelling common terms: H₂O + 2 MnO₄⁻ + 3 SO₃²⁻ → 2 MnO₂ + 3 SO₄²⁻ + 2 OH⁻ (Adding back spectator ions like K⁺ and Na⁺ gives the equation from the original text, often written with K and Na compounds).

Neutral Medium: Often, you can balance it using H⁺ and H₂O like the acidic method, and then if needed, convert H⁺ to OH⁻ by adding OH⁻ to both sides (H⁺ + OH⁻ → H₂O). But sometimes, the reaction naturally involves H₂O directly.

• Example: Combustion of propane (C₃H₈) – not typically done electrochemically in a simple way, but shows the balancing method.
  • Unbalanced: C₃H₈ + O₂ → CO₂ + H₂O
  • Reduction (O₂ gains electrons): Add H⁺ and H₂O. 4 H⁺ + O₂ + 4 e⁻ → 2 H₂O
  • Oxidation (C₃H₈ loses electrons): Add H₂O, H⁺, and electrons. 6 H₂O + C₃H₈ → 3 CO₂ + 20 e⁻ + 20 H⁺
  • Combine by making electrons equal (multiply reduction by 5). 20 H⁺ + 5 O₂ + 20 e⁻ → 10 H₂O 6 H₂O + C₃H₈ → 3 CO₂ + 20 e⁻ + 20 H⁺
  • Add them up, cancelling common terms: C₃H₈ + 5 O₂ → 3 CO₂ + 4 H₂O

Balancing redox reactions is a fundamental skill for understanding what’s happening chemically in an electrochemical system.

How a Galvanic Cell Works (e.g., Daniell Cell)#

Imagine two different metal electrodes, each in a solution containing ions of that metal. For example, a zinc electrode in zinc sulfate solution and a copper electrode in copper sulfate solution.

• The anode is where oxidation happens. In the Daniell cell, zinc is more easily oxidized than copper. So, zinc metal atoms lose electrons and become zinc ions (Zn²⁺) dissolving into the solution. The electrons are left behind on the zinc electrode.
  • Anode half-reaction (oxidation): Zn(s) → Zn²⁺(aq) + 2 e⁻
• These electrons travel through an external circuit (like a wire) from the anode to the cathode. This flow of electrons is the electric current we can use!
• The cathode is where reduction happens. In the Daniell cell, copper ions (Cu²⁺) from the solution gain electrons from the copper electrode surface and turn back into solid copper metal, which plates onto the electrode.
  • Cathode half-reaction (reduction): Cu²⁺(aq) + 2 e⁻ → Cu(s)
• The overall reaction combines the two: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s). This reaction happens spontaneously, releasing energy.

For the circuit to be complete, ions need to flow in the electrolyte between the two solutions to balance the charge changes happening at the electrodes. As positive Zn²⁺ ions form at the anode and positive Cu²⁺ ions are used up at the cathode, positive charge builds up in the anode compartment and negative charge builds up in the cathode compartment (from the sulfate ions left behind). To counteract this, a salt bridge or porous barrier connects the two solutions.

Salt Bridge: A laboratory device used to connect the oxidation and reduction half-cells of a galvanic cell, allowing the flow of ions to maintain electrical neutrality in each compartment and complete the circuit. It typically contains an electrolyte solution (like KCl or KNO₃) in a gel.

In the salt bridge, negative ions flow towards the anode compartment to neutralize the excess positive charge, and positive ions flow towards the cathode compartment to neutralize the excess negative charge. This ionic flow in the electrolyte, combined with the electron flow in the external circuit, creates a complete loop for charge.

The voltage difference between the anode and cathode is the electromotive force (emf) or cell potential, which drives the current in the external circuit.

We can represent an electrochemical cell using a cell diagram:

• It shows the components in order from anode to cathode.
• A single vertical line (|) shows a phase boundary (e.g., solid electrode touching liquid electrolyte).
• A double vertical line (||) represents the salt bridge or porous barrier connecting the two half-cells.

For the Daniell cell (Zinc anode, Copper cathode): Zn(s) | Zn²⁺ (concentration) || Cu²⁺ (concentration) | Cu(s)

This diagram tells you:

• On the left (anode): Solid zinc is in contact with a zinc ion solution. Zinc is getting oxidized.
• In the middle: There’s a salt bridge.
• On the right (cathode): A copper ion solution is in contact with solid copper. Copper ions are getting reduced.

The concentrations of the electrolyte solutions are important because, as we’ll see, they affect the cell voltage.

Concentration Cells#

A cool example of the Nernst equation in action is a concentration cell.

Concentration Cell: A type of electrochemical cell where both half-cells are made of the same materials, but the electrolyte concentrations in the two half-cells are different. A potential difference is generated solely due to the difference in concentration.

Think back to the Daniell cell example, but imagine having two copper electrodes. One is in a dilute copper sulfate solution (e.g., 0.05 M Cu²⁺), and the other is in a concentrated copper sulfate solution (e.g., 2.0 M Cu²⁺).

The reactions are the same in both half-cells: Cu²⁺ + 2e⁻ ⇌ Cu(s). However, according to Le Chatelier’s principle, the reduction reaction (Cu²⁺ + 2e⁻ → Cu) is more favored where the reactant (Cu²⁺) concentration is higher. So, reduction will happen in the concentrated side, making it the cathode. Oxidation (Cu → Cu²⁺ + 2e⁻) is the reverse, favored where the product (Cu²⁺) concentration is lower. So, oxidation will happen in the dilute side, making it the anode.

The cell diagram for this copper concentration cell is: Cu(s) | Cu²⁺ (0.05 M) || Cu²⁺ (2.0 M) | Cu(s)

The overall reaction is just the movement of Cu²⁺ ions from the concentrated side to the dilute side until the concentrations are equal: Cu²⁺ (2.0 M) → Cu²⁺ (0.05 M)

For a concentration cell, the standard cell potential (E°cell) is always zero because the electrodes and ion types are the same in both half-cells (they have the same standard reduction potential). The potential difference comes only from the concentration difference.

Using the Nernst equation (at 298 K): Ecell = E°cell - (0.05916 V / n) log Q Ecell = 0 - (0.05916 V / 2) log ([Cu²⁺]diluted / [Cu²⁺]concentrated) Ecell = - (0.05916 V / 2) log (0.05 / 2.0) Ecell = - (0.02958 V) log (0.025) Ecell = - (0.02958 V) * (-1.602) Ecell ≈ +0.0474 V

A small but measurable voltage is generated as the system tries to reach equilibrium (equal concentrations).

Concentration cells are important not just in theory but also in biological systems, like nerve cells, where ion concentration differences across membranes create electrical potentials.

Batteries#

Batteries are perhaps the most common application. They are basically self-contained galvanic cells (or collections of cells) that provide portable electrical power from spontaneous chemical reactions.

• Wet Cells: Early batteries used liquid electrolytes (like the first telegraph and telephone systems).
• Dry Cells: Like the old zinc-carbon cells, they use a paste electrolyte, making them portable and non-spillable (flashlights, early radios).
• Lead-Acid Batteries: The classic car battery. These are secondary (rechargeable) batteries. The reaction producing electricity is reversible. You can “recharge” them by applying an external voltage to drive the non-spontaneous reverse reaction, converting electrical energy back into stored chemical energy. They typically use lead plates in a sulfuric acid/water mix. While heavy and sensitive to deep discharge, they are reliable and cost-effective for high power needed for starting engines.
• Mercury Batteries: Used in early small electronic devices, but phased out due to environmental concerns from mercury pollution.
• Lithium-ion Batteries: The workhorse of modern portable electronics (phones, laptops) and increasingly electric vehicles. They use non-aqueous electrolytes, allowing for higher voltages per cell, lighter weight, and better low-temperature performance compared to water-based batteries. They are rechargeable (secondary).
• Flow Batteries: An experimental type for large-scale energy storage (like grid backup). Reactants are stored in external tanks and pumped through the cell. This design allows for scaling energy capacity independently of power output.
• Fuel Cells: Convert chemical energy from fuels (like hydrogen, natural gas) and an oxidant (like oxygen) directly into electrical energy with high efficiency. They don’t need to be “recharged” in the traditional sense; they run as long as fuel and oxidant are supplied. Used in spacecraft and potential future power sources.

Corrosion#

Corrosion, like rust on iron, is a natural electrochemical process where metals are oxidized (lose electrons) by substances in their environment, often oxygen and water.

Corrosion: The destructive attack of a metal by chemical or electrochemical reaction with its environment. Rust (iron oxide) and tarnish are common examples.

How Iron Rusts (Iron Corrosion): Iron rusting is an electrochemical phenomenon. It needs iron metal, oxygen, and water. Different spots on the metal surface act as tiny anodes and cathodes.

1. Anode: Iron is oxidized, losing electrons: Fe(s) → Fe²⁺(aq) + 2 e⁻
2. Cathode: Oxygen from the air is reduced. In acidic conditions (like where CO₂ from the air dissolves in water to make carbonic acid), this is: O₂(g) + 4 H⁺(aq) + 4 e⁻ → 2 H₂O(l) (In neutral or basic water, the reaction involves water and forms OH⁻ ions: O₂(g) + 2 H₂O(l) + 4 e⁻ → 4 OH⁻(aq)).
3. Overall Reaction: The electrons flow through the metal, and ions flow through the water film (electrolyte). 2 Fe(s) + O₂(g) + 4 H⁺(aq) → 2 Fe²⁺(aq) + 2 H₂O(l) (or in neutral water: 2 Fe(s) + O₂(g) + 2 H₂O(l) → 2 Fe(OH)₂(s), followed by further oxidation)
4. Further Oxidation: The Fe²⁺ ions formed are then further oxidized by oxygen in the presence of water to form hydrated iron(III) oxide (Fe₂O₃·xH₂O), which is rust. 4 Fe²⁺(aq) + O₂(g) + (4 + 2x) H₂O(l) → 2 Fe₂O₃·xH₂O(s) + 8 H⁺(aq)

Rust doesn’t stick tightly to the iron and flakes off, exposing fresh metal to corrode. This is why iron can rust completely through. Electrolytes like salt dissolved in water make corrosion happen faster because they improve the ionic conductivity, allowing the internal circuit to function more efficiently.

Corrosion of Other Metals:

• Copper: Forms a greenish patina (copper carbonate) over time.
• Silver: Tarnishes to black silver sulfide when exposed to sulfur compounds.
• Aluminum and Titanium: These metals corrode very easily, but they form a very thin, tough, invisible oxide layer instantly when exposed to air. This layer sticks tightly and protects the metal underneath from further corrosion. This is called passivation.

Preventing Corrosion: Since corrosion is electrochemical, preventing it involves stopping the electrochemical cell from forming or operating.

1. Coating: Covering the metal with paint, plastic, or another less reactive metal (passivation like chrome plating) physically blocks contact with the electrolyte (water, oxygen). If the coating is scratched, corrosion can start at the exposed spot, sometimes even faster because the exposed metal becomes the anode relative to the larger coated area (which acts like a cathode).
2. Sacrificial Anodes: Attach a more reactive metal (one that is more easily oxidized, i.e., has a lower standard reduction potential) to the metal you want to protect. This more reactive metal becomes the anode and corrodes away (“sacrifices” itself), forcing the protected metal to be the cathode, where reduction happens instead of oxidation.
   • Examples: Zinc bars on steel ship hulls, or blocks of magnesium buried near steel pipelines. The zinc or magnesium corrodes, protecting the steel. These sacrificial anodes must be replaced periodically. Zinc is often chosen because it’s effective and relatively inexpensive compared to magnesium (which would work even better but is more expensive).

Electrolysis#

Electrolysis is the opposite of a battery; it uses electricity to drive a chemical reaction that wouldn’t happen spontaneously. This happens in an electrolytic cell.

Electrolysis: The process of using electrical energy to drive a non-spontaneous chemical reaction.

You need an external power supply (like a battery or DC generator) to provide the voltage and current needed for electrolysis.

Electrolysis of Molten Sodium Chloride (Downs Cell): This is how we get metallic sodium and chlorine gas. Solid NaCl won’t conduct electricity well, but molten NaCl (melted salt, around 800°C) has mobile Na⁺ and Cl⁻ ions.

• The power supply pumps electrons into the cathode and pulls electrons out of the anode.
• At the cathode (where electrons are supplied), positive sodium ions (Na⁺) gain electrons and are reduced to liquid sodium metal: Cathode (Reduction): 2 Na⁺(l) + 2 e⁻ → 2 Na(l)
• At the anode (where electrons are removed), negative chloride ions (Cl⁻) lose electrons and are oxidized to chlorine gas: Anode (Oxidation): 2 Cl⁻(l) → Cl₂(g) + 2 e⁻
• Overall Reaction: 2 Na⁺(l) + 2 Cl⁻(l) → 2 Na(l) + Cl₂(g)

This process is non-spontaneous. The standard potential is around -4 V. You need to apply at least 4 V from the power supply to make it happen, and usually more to make it go at a practical speed (this extra voltage is called overvoltage).

Electrolysis of Water: Splitting water into hydrogen and oxygen gas also requires energy and is done by electrolysis. Water itself doesn’t conduct electricity well, so an electrolyte (like sulfuric acid or sodium chloride) is usually added to carry charge. Inert electrodes, often platinum, are used.

• At the anode (oxidation): Water is oxidized to oxygen gas, releasing H⁺ ions and electrons. Anode (Oxidation): 2 H₂O(l) → O₂(g) + 4 H⁺(aq) + 4 e⁻
• At the cathode (reduction): Water is reduced to hydrogen gas, producing OH⁻ ions. Cathode (Reduction): 2 H₂O(l) + 2 e⁻ → H₂(g) + 2 OH⁻(aq) (Alternatively, if there’s sufficient H⁺ from an acidic electrolyte: 2 H⁺(aq) + 2 e⁻ → H₂(g) - this is the reaction at the SHE, for example).
• To get the overall reaction, balance the electrons (multiply the cathode reaction by 2): 4 H₂O(l) + 4 e⁻ → 2 H₂(g) + 4 OH⁻(aq) Combine with the anode reaction: 2 H₂O(l) → 2 H₂(g) + O₂(g) (Notice that if you add the 4 H⁺ from the anode to the 4 OH⁻ from the cathode, you get 4 H₂O, which balances with the starting 4 H₂O and the 2 H₂O used in the reduction half-reaction, resulting in the net equation shown).

This reaction requires a voltage (around 1.23 V thermodynamically, but practically closer to 2 V or more due to overvoltage, especially for oxygen formation). Platinum helps by lowering the overvoltage needed.

Electrolysis of Aqueous Solutions (e.g., NaCl solution): This gets a bit trickier because you have water and ions from the dissolved salt present. At each electrode, there’s a competition between the ions and water to be oxidized or reduced.

• Consider a solution of NaCl in water. You have Na⁺, Cl⁻, and H₂O.
• At the Cathode (Reduction): Possible reactions are Na⁺ + e⁻ → Na (E°red = -2.71 V) and 2 H₂O + 2 e⁻ → H₂ + 2 OH⁻ (E°red = -0.83 V). Water is much easier to reduce than Na⁺ (less negative E°red), so hydrogen gas is produced at the cathode.
• At the Anode (Oxidation): Possible reactions are 2 Cl⁻ → Cl₂ + 2 e⁻ (E°oxi = -1.36 V, so E°red = +1.36 V) and 2 H₂O → O₂ + 4 H⁺ + 4 e⁻ (E°oxi = -1.23 V, so E°red = +1.23 V). Based purely on standard potentials, water seems easier to oxidize (less positive E°red). However, overvoltage for oxygen formation at many electrode surfaces is significantly higher than for chlorine formation. This means it takes more extra voltage than expected to make oxygen. So, often, the oxidation of Cl⁻ is kinetically favored and happens even though water’s standard potential suggests it should be easier. Chlorine gas is usually produced at the anode in concentrated NaCl solutions.

So, the overall reaction for electrolyzing concentrated aqueous NaCl is: 2 NaCl(aq) + 2 H₂O(l) → 2 NaOH(aq) + H₂(g) + Cl₂(g)

This process produces hydrogen gas, chlorine gas, and sodium hydroxide (caustic soda), all important industrial chemicals. The concentration of OH⁻ ions increases in the solution.

Quantitative Electrolysis: Faraday’s Laws#

Michael Faraday’s work in the 1830s put numbers to electrolysis. His laws relate the amount of chemical change directly to the amount of electricity passed through the cell.

Faraday’s Laws of Electrolysis: Quantitative relationships describing the amount of chemical change that occurs during electrolysis based on the amount of electrical charge passed.

• First Law: The mass of a substance produced or consumed at an electrode during electrolysis is directly proportional to the total amount of electrical charge passed through the cell. Think of it this way: more electrons you push through, the more stuff reacts. Mathematically: m = (Q * M) / (n * F) Where:

  • m = mass of substance produced/consumed (grams)
  • Q = total electric charge passed (Coulombs). Q = current (Amps) * time (seconds)
  • M = molar mass of the substance (grams/mole)
  • n = number of electrons transferred per ion/molecule in the half-reaction (valence number)
  • F = Faraday’s constant (approx. 96485 C/mol e⁻)

  This equation can be rearranged to find the amount of charge needed to produce a certain mass, or the mass produced by a certain current over time.

• Second Law: When the same amount of electrical charge is passed through different electrolytic cells containing different substances, the masses of the substances produced or consumed at the electrodes are in the ratio of their chemical equivalent weights.

  Chemical Equivalent Weight: The molar mass (M) of a substance divided by the number of electrons transferred (n) in its redox reaction (M/n). It represents the mass of the substance that reacts with or produces one mole of electrons. Essentially, one Faraday (one mole of electrons) will react exactly one equivalent weight of any substance. If you pass 1 Faraday of charge through cells containing molten NaCl and molten CaCl₂, you’ll produce 1 mole of Na (Na⁺ + e⁻ → Na) but only 0.5 moles of Ca (Ca²⁺ + 2e⁻ → Ca), because Ca requires 2 electrons per atom. The masses produced (1 mol Na * 23 g/mol = 23 g, vs 0.5 mol Ca * 40 g/mol = 20 g) are proportional to their equivalent weights (23/1 = 23 for Na, vs 40/2 = 20 for Ca).

Faraday’s laws are fundamental for quantitative applications of electrolysis, like electroplating and calculating yields in industrial electrochemical processes.