Free Printable Oxidation Reduction Worksheets - Free Printable

Let’s go step by step through each part of the worksheet.

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Part 1: Explain the meaning of each term

We need to define these four key redox terms simply and clearly.

- Oxidation: This is when a substance *loses electrons*. You can also think of it as an increase in oxidation number (like going from Fe²⁺ to Fe³⁺). A helpful memory trick: “OIL RIG” — Oxidation Is Loss, Reduction Is Gain (of electrons).

- Reduction: This is when a substance *gains electrons*. It’s the opposite of oxidation. Also, the oxidation number decreases (like Cu²⁺ becoming Cu metal). Remember: “RIG” — Reduction Is Gain.

- Oxidizing Agent: This is the substance that *causes oxidation* in another substance. To do that, it must *accept electrons itself*, so it gets reduced. So: oxidizing agent = gets reduced.

- Reducing Agent: This is the substance that *causes reduction* in another substance. To do that, it must *give away electrons*, so it gets oxidized. So: reducing agent = gets oxidized.

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Part 2: Explain what each phrase means

These are about predicting which substances can react with others based on their ability to gain or lose electrons. We’ll use standard reduction potentials implicitly (even if not listed) — stronger oxidizing agents have higher reduction potentials; stronger reducing agents have lower ones.

But since no table is given, we explain conceptually:

1. “Can be an oxidizing agent for Pb but not Sn²⁺”
→ Means this substance can take electrons from Pb (so Pb gets oxidized to Pb²⁺), but cannot take electrons from Sn²⁺ (so Sn²⁺ stays as is).
→ In other words, it’s strong enough to oxidize Pb, but not strong enough to oxidize Sn²⁺.
→ So its strength as an oxidizing agent is between Pb²⁺/Pb and Sn²⁺/Sn couples.

2. “Can be reduced by I⁻ but not by Fe²⁺”
→ Means I⁻ can give electrons to this substance (so I⁻ gets oxidized to I₂), but Fe²⁺ cannot give electrons to it (so Fe²⁺ stays as Fe²⁺).
→ So this substance is a weaker oxidizing agent than I₂ (because I⁻ can reduce it), but stronger than whatever would oxidize Fe²⁺ (since Fe²⁺ can’t reduce it).
→ Actually, more precisely: It can accept electrons from I⁻ (so it’s reduced), but not from Fe²⁺ → so it’s a stronger oxidizing agent than I₂? Wait — let’s flip it.

Actually: If I⁻ reduces it, then I⁻ is the reducing agent → so this substance is being reduced → so it’s an oxidizing agent. And since Fe²⁺ cannot reduce it, that means Fe²⁺ is not strong enough a reducing agent to donate electrons to it → so this oxidizing agent is too weak to pull electrons from Fe²⁺? No — wait.

Better way: Think in terms of spontaneity.

If X can be reduced by I⁻ → reaction: X + I⁻ → X⁻ + ½I₂ happens spontaneously → so X has higher reduction potential than I₂.

If X cannot be reduced by Fe²⁺ → reaction: X + Fe²⁺ → X⁻ + Fe³⁺ does NOT happen → so X has lower reduction potential than Fe³⁺/Fe²⁺ couple.

Wait — actually, if Fe²⁺ cannot reduce X, that means X is not a strong enough oxidizing agent to take electrons from Fe²⁺ → so X’s reduction potential is less than that of Fe³⁺/Fe²⁺.

But I⁻ CAN reduce X → so X’s reduction potential is greater than I₂/I⁻.

So overall: I₂/I⁻ < X < Fe³⁺/Fe²⁺ in reduction potential.

But again — without numbers, we just explain the behavior.

Simpler explanation for student:

→ This substance can take electrons from iodide ions (I⁻), turning them into iodine (I₂), but it cannot take electrons from iron(II) ions (Fe²⁺). So it’s a moderate oxidizing agent — strong enough to work with I⁻, but not strong enough to work with Fe²⁺.

3. “Can be oxidized by I₂ but not acidic SO₄²⁻”
→ Means I₂ can take electrons from this substance (so this substance gets oxidized, I₂ gets reduced to I⁻), but acidic sulfate (which usually means H⁺ and SO₄²⁻, acting as oxidizing agent via H⁺ or sometimes S in +6 state) cannot oxidize it.

→ So this substance is easy enough to oxidize that even I₂ (a mild oxidizing agent) can do it, but acidic sulfate (which is a stronger oxidizing agent under some conditions — like in hot concentrated form) still can’t? That seems contradictory.

Wait — actually, acidic SO₄²⁻ typically refers to dilute sulfuric acid, which isn’t a strong oxidizer — only hot concentrated H₂SO₄ is. But in many contexts, “acidic SO₄²⁻” might imply the presence of H⁺, which can act as oxidizing agent (to make H₂ gas).

Perhaps better interpretation:

“Can be oxidized by I₂” → I₂ takes electrons from it → so this substance is a reducing agent strong enough to reduce I₂ to I⁻.

“Not by acidic SO₄²⁻” → acidic sulfate (probably meaning H⁺ in acid solution) cannot oxidize it → so this substance is not strong enough a reducing agent to reduce H⁺ to H₂.

That makes sense! Because many metals don’t react with acids (like Cu doesn’t react with HCl), but might react with I₂.

Example: Copper metal can be oxidized by I₂ to Cu²⁺? Actually, no — copper doesn’t react with I₂ easily. Let’s pick zinc: Zn reacts with acid (H⁺) AND with I₂. Not good.

What about silver? Ag doesn’t react with H⁺ (doesn’t dissolve in non-oxidizing acids), but can it be oxidized by I₂? Silver iodide forms, but that’s precipitation, not redox.

Actually, perhaps they mean something like sulfite ion? SO₃²⁻ can be oxidized by I₂ to SO₄²⁻, and acidic SO₄²⁻ won’t oxidize it further.

Yes! That fits.

So: Substance = e.g., SO₃²⁻ or similar.

Explanation: This substance loses electrons to iodine (I₂ becomes I⁻), so it gets oxidized. But when you put it in acid with sulfate ions, nothing happens — the acid/sulfate doesn’t take its electrons. So it’s a mild reducing agent — strong enough to reduce I₂, but not strong enough to reduce H⁺ (from acid).

4. “Can oxidize Co and reduce H⁺”
→ This substance can take electrons from cobalt metal (Co → Co²⁺ + 2e⁻), so it acts as oxidizing agent toward Co.

AND it can give electrons to H⁺ (so H⁺ → H₂), meaning it acts as reducing agent toward H⁺.

Wait — that would mean the same substance is both oxidizing and reducing? That’s possible only if it’s undergoing disproportionation, but here it says “can oxidize Co and reduce H⁺” — probably meaning it can do both reactions, but not necessarily at the same time.

More likely: It’s describing a substance that is capable of acting as an oxidizing agent (toward Co) and also capable of acting as a reducing agent (toward H⁺). But that would require it to have intermediate oxidation states.

Alternatively — maybe it’s poorly worded, and it means: There exists a substance that can oxidize Co (so it’s reduced) and separately, there exists a substance that can reduce H⁺ (so it’s oxidized). But the phrasing suggests one substance doing both.

Another interpretation: Perhaps it’s saying “this substance can cause Co to be oxidized (so it itself is reduced), and this same substance can cause H⁺ to be reduced (so it itself is oxidized)” — which implies it can act as both oxidizing and reducing agent — i.e., it’s amphoteric in redox sense.

Example: Hydrogen peroxide (H₂O₂) can act as oxidizing agent (reduced to H₂O) or reducing agent (oxidized to O₂).

So: H₂O₂ can oxidize Co? Cobalt metal to Co²⁺? Possibly. And H₂O₂ can reduce H⁺? No — reducing H⁺ means making H₂, which H₂O₂ doesn’t do.

Actually, to reduce H⁺, you need a reducing agent like Zn or Mg.

To oxidize Co, you need an oxidizing agent like H⁺ (if Co were reactive) or stronger.

Cobalt metal reacts with acids to produce H₂ — so H⁺ oxidizes Co. So if a substance can oxidize Co, it could be H⁺ itself. But then can H⁺ reduce H⁺? No.

This is confusing.

Let me re-read: “can oxidize Co and reduce H⁺”

Perhaps it’s two separate capabilities of the same species? Like, for example, Fe²⁺ can be oxidized to Fe³⁺ (so it reduces something else), and Fe³⁺ can be reduced to Fe²⁺ (so it oxidizes something else). But here it’s “oxidize Co” — so the substance must be reduced while oxidizing Co. And “reduce H⁺” — so the substance must be oxidized while reducing H⁺.

So for one substance to do both, it must be able to be both oxidized and reduced — i.e., have multiple oxidation states.

Example: MnO₄⁻ can oxidize many things, but can it reduce H⁺? No.

Better example: Nitrous acid, HNO₂. It can act as oxidizing agent (reduced to NO) or reducing agent (oxidized to NO₃).

Can HNO₂ oxidize Co? Cobalt metal to Co²⁺? Possibly, if acidic.

Can HNO₂ reduce H⁺? Reducing H⁺ means producing H₂ gas — HNO₂ doesn’t do that.

Actually, to reduce H⁺, you need a metal above hydrogen in activity series.

Perhaps the intended meaning is: The substance can serve as an oxidizing agent (for Co) and also serve as a reducing agent (for H⁺) — but not necessarily in the same reaction.

And for that, it needs to have an intermediate oxidation state.

Common example: Fe²⁺. Fe²⁺ can be oxidized to Fe³⁺ (so it reduces