Identify oxidation and reduction, assign oxidation numbers, balance simple redox reactions, and explain galvanic and electrolytic cells conceptually.
A lesson of Chemistry, in Equilibrium, Acids and Bases, and Redox, about 35 minutes.
By the end of this lesson, you will be able to assign oxidation numbers, identify what was oxidized and reduced, explain how a battery works, and say why corrosion and electroplating are the same chemistry pointed different ways.
Zinc lost two electrons, so zinc was oxidized. Copper gained them, so copper was
reduced.
Neither happens alone. Electrons have to go somewhere, so every oxidation is
paired with a reduction and the two half-reactions always occur together. That is
why the topic is one word rather than two.
"Reduction" sounds like an increase and is not. The name comes from
metallurgy: reducing an ore reduced its mass by stripping away oxygen, and the
electron definition came centuries later. The word is historical, and the mnemonic
is more reliable than the etymology.
The two agents are named for what they do to the other species, which is the
second confusing piece of naming:
The reducing agent is the one that gets OXIDIZED, because it hands its
electrons to something else
The oxidizing agent is the one that gets REDUCED
Zinc is the reducing agent here and is itself oxidized.
Oxidation Numbers
A bookkeeping device: the charge an atom would have if every bond were fully
ionic. It tracks electrons without requiring you to see them.
The rules, in order of priority:
A free element is $0$. $\mathrm{Zn}$, $\mathrm{O_2}$, $\mathrm{Cl_2}$
A monatomic ion is its charge. $\mathrm{Na^+}$ is $+1$
Oxygen is $-2$, except in peroxides where it is $-1$
Hydrogen is $+1$ with nonmetals, $-1$ with metals
Fluorine is always $-1$
The total must equal the overall charge, which is $0$ for a neutral compound
That last rule is the one that does the work, because it lets you solve for an
unknown.
$$\mathrm{H_2SO_4}: \quad 2(+1) + S + 4(-2) = 0 \quad \Rightarrow \quad S = +6$$
The change in oxidation number tells you what happened.
$$\text{oxidation number INCREASES} \Rightarrow \text{oxidized}$$
$$\text{oxidation number DECREASES} \Rightarrow \text{reduced}$$
That is consistent with losing electrons making you more positive, and it lets you
identify redox in reactions where no obvious ions appear at all.
Electrons appear on the right of an oxidation and the left of a reduction,
because they are being produced in one and consumed in the other.
The electron counts must match before adding, and if they do not, multiply one
half-reaction until they do. Adding the two halves then cancels the electrons and
recovers the whole equation.
This is more than presentation. In a battery the two halves happen in
physically separate places, connected by a wire, and the half-reaction is
literally what is occurring at each electrode.
Sort It: Oxidation or Reduction Half-Reaction
A Battery Separates the Halves
Put zinc metal into copper sulfate solution and the electrons transfer directly at
the surface, releasing energy as heat. Useless.
Separate the two halves and make the electrons travel through a wire, and the
same reaction becomes electricity.
The anode is where oxidation happens, and it is negative in a battery
The cathode is where reduction happens, and it is positive
Electrons flow through the external wire from anode to cathode
A salt bridge lets ions move between the two compartments to keep each side
electrically neutral
Without the salt bridge nothing flows for more than an instant. The anode
compartment accumulates positive charge as metal dissolves, the cathode side
accumulates negative charge, and that imbalance stops the electron flow at once.
An mnemonic that survives:AN OX and RED CAT, for ANode OXidation
and REDuction at the CAThode.
The voltage comes from the activity series. The further apart the two metals
sit, the more strongly one wants the electrons and the higher the cell voltage. A
zinc-copper cell gives about $1.1$ V, and that difference in electron-hunger is all
a battery is.
Worked Example 2: Reading a Cell
A zinc electrode in zinc sulfate, a copper electrode in copper sulfate, joined
by a wire and a salt bridge.
Iron gives up electrons to oxygen. Water is required, which is why iron does not
rust in dry air, and salt accelerates it badly by carrying charge between the
anodic and cathodic patches on the metal surface.
Rust flakes off, exposing fresh metal underneath, so it never protects. That is
the crucial difference from aluminum, which forms a thin adherent oxide that seals
the surface and stops further attack. Aluminum is the more reactive metal and the
more durable one.
Sacrificial protection exploits redox deliberately: bolt a block of a more
reactive metal to the thing you want to keep. Zinc or magnesium blocks on a ship's
hull corrode instead of the steel, because they give up electrons more readily, and
they are replaced periodically. Galvanized steel is the same idea as a coating.
Electroplating runs a redox reaction backwards using electricity. Make the
object the cathode, put the plating metal's ions in solution, and drive current
through; metal ions gain electrons and deposit on the surface. That is how chrome,
silver and gold plating are done, and it is the same equipment as electrolysis of
water or purification of copper.
A steel pipeline is protected by attached magnesium blocks.
Common Mistakes
Mistake 1: Thinking reduction means getting smaller
It means GAINING electrons. The name is historical.
Mistake 2: Swapping the agents
The reducing agent is the one that gets oxidized.
Mistake 3: Forgetting the total-charge rule
Oxidation numbers in a neutral compound must sum to zero.
Mistake 4: Adding half-reactions with unequal electrons
Multiply until the counts match.
Mistake 5: Putting oxidation at the cathode
AN OX and RED CAT.
Where This Goes Next
Redox is the largest category of chemistry there is, and it powers everything you
do.
Respiration is a redox reaction. Glucose is oxidized and oxygen is reduced,
and the electrons pass down a chain of carriers inside your mitochondria, releasing
their energy in controlled steps rather than all at once. Combustion does the same
chemistry in one violent step; your cells do it in about thirty, and capture the
energy as ATP.
Photosynthesis is that reaction run backwards, driven by sunlight, and it is
the only reason there is oxygen in the atmosphere at all.
Battery technology is applied redox, and its limits are chemical. A
lithium-ion cell works because lithium sits at the extreme reactive end of the
series, giving the largest possible voltage per unit mass. The search for better
batteries is a search for redox pairs that are far apart, light, cheap, safe and
reversible for thousands of cycles, and those requirements conflict.
Fuel cells combine hydrogen and oxygen electrochemically rather than by
burning, producing electricity and water with no combustion at all, at roughly
twice the efficiency of an engine.
Almost every energy transformation that matters, biological or industrial, is
electrons moving from something that holds them loosely to something that holds
them tightly. Learning to track that movement is what the oxidation-number
bookkeeping is for.
Check Your Understanding
Five of them, no instructions this time.
Summary
Oxidation is loss, reduction is gain, and the two always happen together.
The reducing agent is the one that gets oxidized, which is the naming that
catches everyone.
Oxidation numbers track electrons on paper, and the totals must equal the
overall charge.
A battery separates the two half-reactions so the electrons travel through a
wire: AN OX, RED CAT, with a salt bridge keeping both sides neutral.
Corrosion is unwanted oxidation, and electroplating is the same chemistry driven
backwards by electricity.