Predicting Reactions
Lajoy Tucker & Dr. Davinder Bhachu
Teachers
Contents
Introduction Predicting Reactions
Understanding the Electrochemical Series
What is the Electrochemical Series?
The electrochemical series is like a "league table" of half-reactions, ranked by their standard electrode potentials from most positive to most negative. Think of it as a competition where the most positive electrodes are the "strongest" at attracting electrons. Electrochemists, by convention, use reduction potentials.
The Electrochemical Series Table
Half-Reaction | E° (V) | Oxidizing Agent | Reducing Agent |
|---|---|---|---|
F₂ + 2e⁻ → 2F⁻ | +2.87 | F₂ (strongest) | F⁻ (weakest) |
MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O | +1.51 | MnO₄⁻ | Mn²⁺ |
Cl₂ + 2e⁻ → 2Cl⁻ | +1.36 | Cl₂ | Cl⁻ |
Br₂ + 2e⁻ → 2Br⁻ | +1.07 | Br₂ | Br⁻ |
Ag⁺ + e⁻ → Ag | +0.80 | Ag⁺ | Ag |
Fe³⁺ + e⁻ → Fe²⁺ | +0.77 | Fe³⁺ | Fe²⁺ |
I₂ + 2e⁻ → 2I⁻ | +0.54 | I₂ | I⁻ |
Cu²⁺ + 2e⁻ → Cu | +0.34 | Cu²⁺ | Cu |
2H⁺ + 2e⁻ → H₂ | 0.00 | H⁺ | H₂ |
Pb²⁺ + 2e⁻ → Pb | −0.13 | Pb²⁺ | Pb |
Ni²⁺ + 2e⁻ → Ni | −0.25 | Ni²⁺ | Ni |
Fe²⁺ + 2e⁻ → Fe | −0.44 | Fe²⁺ | Fe |
Zn²⁺ + 2e⁻ → Zn | −0.76 | Zn²⁺ | Zn |
Al³⁺ + 3e⁻ → Al | −1.66 | Al³⁺ | Al |
Mg²⁺ + 2e⁻ → Mg | −2.37 | Mg²⁺ | Mg |
Li⁺ + e⁻ → Li | −3.03 | Li⁺ (weakest) | Li (strongest) |
No answer provided.
Reading the Series - Key Insights
For Oxidizing Agents (species on the LEFT of equations):
More positive = stronger oxidizing agent
is the strongest - it desperately wants electrons
is the weakest - barely interested in electrons
For Reducing Agents (species on the RIGHT of equations):
More negative = stronger reducing agent
is the strongest - gives up electrons very easily
is the weakest - holds onto electrons tightly
No answer provided.
Predicting Reactions Question Explainer Video
Rule for Predicting Reactions
The Universal Principle
"The more positive half-reaction gains the electrons"
This makes perfect sense - positive charges attract negative electrons.
Step-by-Step Method for Predicting Reactions
Step 1: Identify the two relevant half-reactions
Step 2: Compare their values
Step 3: More positive → gains electrons (reduction, goes forward)
Step 4: More negative → loses electrons (oxidation, goes backward)
Step 5: Write the overall equation
Step 6: Calculate to confirm feasibility. A positive cell means that the reaction is feasible.
No answer provided.
Worked Example 1: Chlorine vs Bromine
Question: What happens when Cl₂ and Br⁻ are mixed?
Given half-reactions:
Step 1: Compare values
Step 2: Apply the golden rule
gets the electrons → (goes forward)
loses electrons → (reverse direction)
Step 3: Overall reaction
Step 4: Calculate
Conclusion: Positive confirms the reaction is feasible!
No answer provided.
Worked Example 2: Zinc and Copper
Question: Why do Zn and react spontaneously?
Given:
Analysis:
Half-reactions:
Reduction:
Oxidation:
Overall:
✓ Feasible!
No answer provided.
The Gibbs Free Energy Connection
Fundamental Relationship: ΔG = -nFE°cell
Where:
= Gibbs free energy change
= number of electrons transferred
= Faraday constant
= standard cell potential (V)
If is positive → is negative → Reaction is spontaneous
If is negative → is positive → Reaction is not spontaneous
For :
n = 2 electrons
Large negative ΔG confirms this reaction is highly favorable!
Predicting Whether Reactions Occur
Question 1:
Will oxidize to in acidic solution?
Relevant half-reactions:
Analysis:
Therefore: will oxidize
Balanced equation:
✓
No answer provided.
Question 2:
Rank MnO₄⁻, Cr₂O₇²⁻, and Fe³⁺ as oxidizing agents.
From the series:
Ranking:
Only and can oxidize to (both have )
No answer provided.
Practice Questions
Question 1:
a) Write an equation for the reaction that would take place if the half-cells were connected together.
b) Explain why this reaction takes place.
Given:
Answers:
a) Equation for the reaction:
Step 1: Compare E° values
(more positive)
(more negative)
Step 2: Apply the golden rule - more positive gains electrons
(reduction - goes forward)
(oxidation - reverse direction)
Overall equation:
b) Explanation: This reaction takes place because is more positive than has a greater tendency to gain electrons than . Therefore, will be reduced and Fe will be oxidized.
Question 2:
What reaction would take place if a piece of aluminum and zinc was placed in a solution containing a mixture of aluminum nitrate and zinc sulfate? Justify this using electrode potential data.
Given:
Answer:
Analysis:
(more positive)
(more negative)
Prediction: will be reduced, Al will be oxidized
Balanced equation:
Reduction:
Oxidation:
Overall:
Justification: will preferentially gain electrons from Al.
Question 3:
Which of the species and are able to liberate I₂ from an acidic solution of potassium iodide? Explain your reasoning using the electrode potentials.
Given:
Answer:
For a species to liberate from , it must have
Analysis:
✓ Can oxidize
✓ Can oxidize
✓ Can oxidize
All three species can liberate I₂ from acidic KI solution.
Sample equation for