Redox processes and electrode potentials

Q: What is the "OIL RIG" mnemonic used for in electrochemistry?

Oxidation Is Loss (of electrons) and Reduction Is Gain (of electrons).

Q: How is Standard Electrode Potential $ (E^{\ominus}) $ defined?

It is the voltage produced when a half-cell is connected to a Standard Hydrogen Electrode (SHE) under standard conditions.

Q: What are the specific "Standard Conditions" required for measuring $ E^{\ominus} $?

A concentration of 1.0 mol dm$^{-3}$ for ions, a pressure of 1 bar for gases, and a temperature of 298 K.

Q: What is the Standard Hydrogen Electrode (SHE) and why is it important?

The SHE is a reference electrode assigned a potential of exactly 0.00 V, consisting of H$_2$ gas at 1 bar and H$^{+}$ ions at 1.0 mol dm$^{-3}$ over a Platinum electrode.

Q: When is a Platinum electrode required in a half-cell setup?

A Platinum electrode is used if the half-cell does not contain a conductive solid metal, such as in gas/ion or ion/ion systems.

Q: How do you calculate the Standard Cell Potential $ (E^{\ominus}_{cell}) $ of a complete circuit?

$ E^{\ominus}_{cell} = E^{\ominus}_{reduction} - E^{\ominus}_{oxidation} $ (or $ E^{\ominus}_{cathode} - E^{\ominus}_{anode} $).

Q: Does the stoichiometry of a half-equation affect the value of $ E^{\ominus} $?

No, $ E^{\ominus} $ is an intensive property and its value remains the same regardless of the coefficients used to balance the overall equation.

Q: What is the criteria for a redox reaction to be considered spontaneous based on cell potential?

A reaction is spontaneous if the calculated Standard Cell Potential $ (E^{\ominus}_{cell}) $ is positive (greater than zero).

Q: How is Gibbs Free Energy change $ (\Delta G^{\ominus}) $ related to cell potential?

$ \Delta G^{\ominus} = -nFE^{\ominus}_{cell} $, where n is the moles of electrons and F is the Faraday constant.

Q: Why might a spontaneous reaction (positive $ E^{\ominus}_{cell} $) fail to occur in practice?

It may be limited by kinetics, meaning the activation energy is too high for the reaction to proceed at a detectable rate.

Chemistry (9476) · GCE A-Level - Higher 2 (H2) · 11 min read · Updated: 22 Oct 2026

Welcome to the World of Electrode Potentials!

Hello there! Today, we are diving into Electrode Potentials, a fascinating part of Electrochemistry. Think of this chapter as a "tug-of-war" for electrons. Some elements are "greedy" and want to pull electrons toward themselves, while others are happy to give them away. By the end of these notes, you’ll be able to predict exactly which way electrons will flow and whether a reaction will happen at all. Don’t worry if it seems like a lot of symbols at first—we’ll break it down step-by-step!

1. Prerequisite: The Basics of Redox

Before we start, let's remember a simple rule from O-Levels. Redox stands for Reduction and Oxidation. You can use the mnemonic OIL RIG:

Oxidation Is Loss (of electrons)
Reduction Is Gain (of electrons)

In this chapter, we focus on Standard Electrode Potential $ (E^{\ominus}) $, which measures how much a species "wants" to be reduced (gain electrons).

2. Defining Key Terms

To speak the language of electrochemistry, you need to know these three definitions by heart:

Standard Electrode (Redox) Potential $ (E^{\ominus}) $: The voltage produced when a half-cell is connected to a Standard Hydrogen Electrode (SHE) under standard conditions.

Standard Cell Potential $ (E^{\ominus}_{cell}) $: The potential difference between two half-cells. It tells us the "pushing power" of the entire battery.

Standard Conditions: To keep things fair (like a controlled experiment), we always use:
1. A concentration of 1.0 mol dm$^{-3}$ for all ions in solution.
2. A pressure of 1 bar for all gases.
3. A temperature of 298 K (25°C).

Quick Review:

A positive $ E^{\ominus} $ means the species is easily reduced (a strong oxidising agent).
A negative $ E^{\ominus} $ means the species is hard to reduce but its reverse form is easily oxidised.

3. The Standard Hydrogen Electrode (SHE)

How do we measure the "greediness" of an element? We need a reference point, just like how we measure altitude relative to sea level. In Chemistry, "sea level" is the Standard Hydrogen Electrode.

The SHE is assigned a potential of exactly 0.00 V.

Components of the SHE:

Hydrogen gas at 1 bar.
H$^{+}$(aq) ions at 1.0 mol dm$^{-3}$ (usually HCl).
A Platinum (Pt) electrode. Why Platinum? It is inert (won't react) and provides a surface for the reaction to happen.

Analogy: The SHE is like the "zero" mark on a ruler. Everything else is measured as being "higher" or "lower" than it.

4. Measuring Electrode Potentials

Depending on what chemicals we are testing, we set up the half-cell differently before connecting it to the SHE.

A. Metal / Metal Ion Systems

Example: Measuring the $ E^{\ominus} $ of Copper.
You dip a Copper rod (the electrode) into a solution of Cu$^{2+}$ ions (1.0 mol dm$^{-3}$).

B. Non-metal / Non-metal Ion Systems

Example: Measuring the $ E^{\ominus} $ of Chlorine.
Since Chlorine is a gas, we bubble Cl$_2$ gas over a Platinum electrode dipped in a solution of Cl$^{-}$ ions.

C. Ion / Ion Systems (Different Oxidation States)

Example: Measuring Fe$^{3+}$ / Fe$^{2+}$.
Here, both species are in the same solution. We use a Platinum electrode dipped into a mixture of 1.0 mol dm$^{-3}$ Fe$^{3+}$ and 1.0 mol dm$^{-3}$ Fe$^{2+}$.

Key Takeaway: If there is no solid metal in the reaction to act as a wire, always use a Platinum electrode!

5. Calculating Standard Cell Potential $ (E^{\ominus}_{cell}) $

When you link two half-cells together, you create a circuit. To find the total voltage, use this simple formula:

\[ E^{\ominus}_{cell} = E^{\ominus}_{reduction} - E^{\ominus}_{oxidation} \]

Or, if you prefer the Red Cat / An Ox mnemonic:

Red Cat: Reduction happens at the Cathode (more positive $ E^{\ominus} $).
An Ox: Oxidation happens at the Anode (more negative $ E^{\ominus} $).

Step-by-Step Example:
Zn$^{2+}$ + 2e$^-$ $\rightleftharpoons$ Zn $ (E^{\ominus} = -0.76 V) $
Cu$^{2+}$ + 2e$^-$ $\rightleftharpoons$ Cu $ (E^{\ominus} = +0.34 V) $
1. The Cu half-cell is more positive, so it's the reduction side.
2. The Zn half-cell is more negative, so it's the oxidation side.
3. $ E^{\ominus}_{cell} = (+0.34) - (-0.76) = +1.10 V $

Common Mistake: Never multiply the $ E^{\ominus} $ value by any stoichiometric coefficients! Even if you multiply the whole equation by 2, the voltage stays the same.

6. Predicting Spontaneity

Will a reaction happen on its own? Check the $ E^{\ominus}_{cell} $!

If $ E^{\ominus}_{cell} > 0 $ (Positive): The reaction is spontaneous.
If $ E^{\ominus}_{cell} < 0 $ (Negative): The reaction is non-spontaneous.

The Gibbs Free Energy Connection

There is a mathematical link between cell potential and energy:
\[ \Delta G^{\ominus} = -nFE^{\ominus}_{cell} \]

$ \Delta G^{\ominus} $: Gibbs Free Energy change (must be negative for spontaneity).
n: Number of moles of electrons transferred.
F: Faraday constant (96500 C mol$^{-1}$).

Did you know? This equation proves that a positive voltage always gives a negative Gibbs energy, which means the "battery" has enough energy to run!

7. Limitations of $ E^{\ominus} $

Don't be fooled! Just because $ E^{\ominus}_{cell} $ is positive doesn't mean you'll see a reaction immediately.
1. Kinetic Limitation: The reaction might have a very high activation energy, making it too slow to observe at room temperature.
2. Non-standard Conditions: If the concentration isn't 1.0 mol dm$^{-3}$, the actual potential will change.

8. Effects of Concentration (Qualitative)

If we change the concentration, we can use Le Chatelier’s Principle to predict what happens to the electrode potential.

Consider: $ M^{n+}(aq) + ne^- \rightleftharpoons M(s) $

If you increase the concentration of $ M^{n+} $, the equilibrium shifts to the right to use up the extra ions. This makes the reduction easier, so the potential becomes more positive.
If you decrease the concentration, the equilibrium shifts to the left, and the potential becomes more negative.

9. Real World: The $ H_2/O_2 $ Fuel Cell

A fuel cell is a special type of battery that doesn't "die" as long as you keep feeding it fuel.

The Reaction:
Anode: $ 2H_2(g) + 4OH^-(aq) \rightarrow 4H_2O(l) + 4e^- $
Cathode: $ O_2(g) + 2H_2O(l) + 4e^- \rightarrow 4OH^-(aq) $
Overall: $ 2H_2(g) + O_2(g) \rightarrow 2H_2O(l) $

Why is it great?

Eco-friendly: The only byproduct is water!
Efficient: High energy-to-mass ratio.
Lightweight: Great for space shuttles and modern electric vehicles.

Final Summary Checklist

✓ Standard conditions are 298K, 1 bar, 1.0 mol dm$^{-3}$.
✓ SHE is the reference point (0.00V).
✓ $ E^{\ominus}_{cell} = E^{\ominus}_{red} - E^{\ominus}_{ox} $.
✓ Positive $ E^{\ominus}_{cell} $ means the reaction is likely spontaneous.
✓ High activation energy can stop a "spontaneous" reaction from happening.

Keep practicing those calculations! You've got this!

Welcome to the World of Redox!

Welcome! Today, we are diving into one of the most important "behind-the-scenes" processes in Chemistry: Redox. Whether it is the battery powering your phone, the rust on a metal gate, or even the way your body turns food into energy, Redox processes are at work.

At first, the name "Redox" might sound a bit intimidating, but it is actually just a combination of two words: Reduction and Oxidation. In this guide, we will break down how electrons move between atoms and how we can use "Oxidation Numbers" to track them like a chemist detective!

1. The Basics: What is Redox?

A Redox reaction is a chemical reaction that involves the transfer of electrons between two substances. If one substance gives away electrons, another substance must take them. They always happen together—you can't have one without the other!

The "OIL RIG" Mnemonic

Don’t worry if you forget which is which! Use this classic trick to keep it straight:

OIL RIG

Oxidation Is Loss (of electrons)
Reduction Is Gain (of electrons)

Analogy: Think of electrons like money. If you give \$5 to a friend, you are "oxidised" (you lost money) and your friend is "reduced" (they gained money). The total amount of money in the room stays the same; it just changed hands!

Quick Review:
• Oxidation: Losing electrons $ (e^-) $.
• Reduction: Gaining electrons $ (e^-) $.

2. Oxidation Number (Oxidation State)

Sometimes, it is hard to see exactly where electrons are moving, especially in molecules that don't have clear ionic charges. This is where Oxidation Number (O.N.) comes in.

An Oxidation Number is a value assigned to an atom to represent the number of electrons it has "lost" or "gained" relative to its neutral state. Think of it as a "virtual charge."

How to Assign Oxidation Numbers (The Rules)

To be a Redox expert, you just need to follow these simple rules in order:

1. Free Elements: Any element in its uncombined state has an O.N. of 0.
Example: $ Na, H_2, O_2, P_4, S_8 $ all have an O.N. of 0.

2. Simple Ions: The O.N. is the same as the charge of the ion.
Example: $ Na^+ $ is +1; $ Mg^{2+} $ is +2; $ Cl^- $ is -1.

3. The Sum Rule:
• In a neutral compound, the sum of all O.N.s must be 0.
• In a polyatomic ion, the sum of all O.N.s must equal the charge of the ion.

4. Fixed Customers (The "Usually" Rules):
• Fluorine (F): Always -1.
• Oxygen (O): Usually -2 (Exception: -1 in peroxides like $ H_2O_2 $ and +2 in $ F_2O $).
• Hydrogen (H): Usually +1 (Exception: -1 in metal hydrides like $ NaH $).

Common Mistake to Avoid: Don't confuse charge with oxidation number. While they look the same (e.g., +2), we write the sign before the number for O.N. (e.g., +2) and after the number for ion charges (e.g., 2+). It’s a small detail that makes you look like a pro!

Key Takeaway: Oxidation numbers are just a bookkeeping system for electrons. If the O.N. increases, oxidation happened. If it decreases, reduction happened.

3. Identifying Redox via Oxidation Numbers

Now that we have our "bookkeeping" system, we can define Redox in a second way:

• Oxidation is an increase in oxidation number.
• Reduction is a decrease in oxidation number.

Memory Trick: If the number is being REDUCED (getting smaller/more negative), then it is REDUCTION!

Step-by-Step Example:

Let's look at the reaction: $ Zn (s) + Cu^{2+} (aq) \rightarrow Zn^{2+} (aq) + Cu (s) $

1. Assign O.N. to Zn: It’s a free element, so it starts at 0. It becomes $ Zn^{2+} $, so it ends at +2.
0 to +2 is an increase $\rightarrow$ Oxidation.

2. Assign O.N. to Cu: It starts as an ion $ Cu^{2+} $, so it's +2. It becomes a free element $ Cu $, so it ends at 0.
+2 to 0 is a decrease $\rightarrow$ Reduction.

Did you know? This exact reaction is what happens inside a standard zinc-carbon battery! Electrons flow from the Zinc to the Copper, creating the electricity we use.

4. Oxidising and Reducing Agents

This is the part that often trips students up, but here is a simple way to think about it. An "agent" is someone who causes something to happen.

• Oxidising Agent: The substance that oxidises something else. To do this, it must take electrons, so the agent itself gets reduced.
• Reducing Agent: The substance that reduces something else. To do this, it must give electrons, so the agent itself gets oxidised.

Analogy: A "Travel Agent" helps you travel, but the agent doesn't necessarily travel themselves. Similarly, an Oxidising Agent helps another atom get oxidised by taking its electrons!

Summary Table for Quick Reference:

5. Summary and Key Takeaways

You’ve made it through the core concepts of Redox! Here is what you need to remember for your exams:

• Redox stands for Reduction-Oxidation; they always happen at the same time.
• OIL RIG: Oxidation Is Loss, Reduction Is Gain of electrons.
• Oxidation Number: A tool to track electrons. Know your rules for F, O, and H!
• Identification: If the O.N. goes up, it's oxidation. If it goes down, it's reduction.
• Agents: They do the opposite of their name to themselves (e.g., a Reducing Agent is the one being oxidised).

Pro-tip for Struggling Students: If you are ever stuck on a complex equation, always start by assigning Oxidation Numbers to every single element. Once the numbers are written above the atoms, the "jumps" in values will tell you exactly what is being oxidised and what is being reduced!

Quick check

Can you answer these now?

Open each question to check the key ideas from this chapter.

What is the "OIL RIG" mnemonic used for in electrochemistry?

Oxidation Is Loss (of electrons) and Reduction Is Gain (of electrons).

How is Standard Electrode Potential $ (E^{\ominus}) $ defined?

It is the voltage produced when a half-cell is connected to a Standard Hydrogen Electrode (SHE) under standard conditions.

What are the specific "Standard Conditions" required for measuring $ E^{\ominus} $?

A concentration of 1.0 mol dm$^{-3}$ for ions, a pressure of 1 bar for gases, and a temperature of 298 K.

What is the Standard Hydrogen Electrode (SHE) and why is it important?

The SHE is a reference electrode assigned a potential of exactly 0.00 V, consisting of H$_2$ gas at 1 bar and H$^{+}$ ions at 1.0 mol dm$^{-3}$ over a Platinum electrode.

When is a Platinum electrode required in a half-cell setup?

A Platinum electrode is used if the half-cell does not contain a conductive solid metal, such as in gas/ion or ion/ion systems.

How do you calculate the Standard Cell Potential $ (E^{\ominus}_{cell}) $ of a complete circuit?

$ E^{\ominus}_{cell} = E^{\ominus}_{reduction} - E^{\ominus}_{oxidation} $ (or $ E^{\ominus}_{cathode} - E^{\ominus}_{anode} $).

Does the stoichiometry of a half-equation affect the value of $ E^{\ominus} $?

No, $ E^{\ominus} $ is an intensive property and its value remains the same regardless of the coefficients used to balance the overall equation.

What is the criteria for a redox reaction to be considered spontaneous based on cell potential?

A reaction is spontaneous if the calculated Standard Cell Potential $ (E^{\ominus}_{cell}) $ is positive (greater than zero).

How is Gibbs Free Energy change $ (\Delta G^{\ominus}) $ related to cell potential?

$ \Delta G^{\ominus} = -nFE^{\ominus}_{cell} $, where n is the moles of electrons and F is the Faraday constant.

Why might a spontaneous reaction (positive $ E^{\ominus}_{cell} $) fail to occur in practice?

It may be limited by kinetics, meaning the activation energy is too high for the reaction to proceed at a detectable rate.

How does increasing the concentration of ions in a half-cell affect its electrode potential?

Increasing the concentration of the reactant ions shifts the equilibrium to the right, making the electrode potential more positive.

What is the overall reaction and a major benefit of the Hydrogen-Oxygen fuel cell?

The overall reaction is $ 2H_2(g) + O_2(g) \rightarrow 2H_2O(l) $, and its main benefit is that water is the only byproduct.

What is a redox reaction?

A chemical reaction involving the transfer of electrons between two substances, where oxidation and reduction occur simultaneously.

How does the "OIL RIG" mnemonic help define oxidation and reduction?

OIL RIG stands for Oxidation Is Loss and Reduction Is Gain of electrons.

What is an Oxidation Number (O.N.)?

A value assigned to an atom in a compound that represents the number of electrons lost or gained relative to its neutral state.

What is the oxidation number of an element in its uncombined (free) state?

The oxidation number of any free element is always 0.

How is the oxidation number of a simple (monatomic) ion determined?

The oxidation number is equal to the charge of the ion.

What is the rule for the sum of oxidation numbers in compounds and polyatomic ions?

In a neutral compound, the sum of all O.N.s is 0; in a polyatomic ion, the sum equals the charge of the ion.

What are the standard oxidation numbers for Fluorine, Oxygen, and Hydrogen?

Fluorine is always -1; Oxygen is usually -2; Hydrogen is usually +1.

How are oxidation and reduction identified using oxidation numbers?

Oxidation is an increase in O.N., while reduction is a decrease in O.N.

What is an oxidising agent?

A substance that oxidises another substance by accepting electrons, causing its own oxidation number to decrease (it is reduced).

What is a reducing agent?

A substance that reduces another substance by donating electrons, causing its own oxidation number to increase (it is oxidised).

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