Chapter 18: Electrochemistry: Galvanic Cells, Potentials, and Electrolysis

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Welcome back to The Deep Dive.

Today we're looking at something invisible but absolutely essential.

Think about your phone battery or maybe why that old bike in the shed gets rusty.

Right.

It seems simple but underneath it's this fascinating dance between chemistry and electricity.

Exactly.

It's happening constantly atom by atom and that dance has a name electrochemistry.

And it's everywhere.

I mean truly everywhere.

It's how chemical reactions can spit out electricity and how electricity can well make chemical reactions happen that wouldn't normally go.

Like flipping a switch chemically speaking.

Kind of.

From you know super complex things like making aluminum which used to be incredibly hard to.

Medical sensors detecting tiny traces of stuff even just starting your car.

It's all about swapping chemical and electrical energy.

So today our mission is to unpack chapter 18 from Zumdahl, Zumdahl and de Kass chemistry.

We want to guide you through the core ideas of electrochemistry.

Make them click even without any diagrams in front of you.

Yeah we'll break it down step by step.

We'll use analogies.

Try to paint a picture so you can visualize what's happening at that tiny scale.

The goal isn't just what happens but why it's important for health, the environment, industry,

all of it.

Okay let's dive in.

So you said it's a dance.

The fundamental steps are these things called redox reactions.

Exactly.

Redox.

That's short for reduction oxidation.

It's all about electrons moving from one chemical species to another.

And we split these into half reactions?

We do.

One half is oxidation.

That's loss of electrons.

Leo the lion says der loss of electrons is oxidation.

And the other half is reduction, the gain of electrons.

Gain of electrons is reduction.

They always happen together.

Can't have one without the other.

Okay electrons moving but how do we get power from that?

Ah well that was the big challenge.

See if you just mix the chemicals together in one beaker the electrons jump directly and usually the energy is just lost as heat.

Not very useful.

Right just gets warm.

Exactly.

So the brilliant idea was what if we physically separate the oxidizing agent from the reducing agent?

Force the electrons to take the long way around through a wire.

Uh -huh.

So you create a path for them.

Precisely.

And that setup is a galvanic cell.

Sometimes called the voltaic cell.

Think of it like a tiny chemical power plant.

You have the anode where oxidation happens electrons are released and the cathode where reduction happens electrons are consumed.

And the electrons flow anode cathode through the wire.

That's our current.

That's the current.

But wait if electrons leave the anode side and pile up at the cathode side wouldn't like one side get really positive and the other really negative wouldn't that just stop the flow?

Excellent point you've hit on a critical piece of the puzzle.

Yes that charge buildup would stop everything fast.

We need a way to keep things electrically neutral.

So we use the unsung hero the salt bridge or sometimes a porous disc does the same job.

It's basically a connection between the two halves often a tube filled with a salt solution that allows ions not the reacting ones but spectator ions to move between the compartments.

Ah okay so if negative charge builds up near the cathode positive ions from the salt bridge drift over and vice versa.

Exactly.

It completes the circuit internally by balancing the charge letting the electrons keep flowing externally through the wire.

Keeps the whole dance going.

Makes sense.

So we have flow what measures the um the push behind those electrons the driving force.

Good question.

That push or driving force is called the cell potential or sometimes electromotive force and we measure it in volts.

Volts okay like a battery voltage.

Precisely and one volt basically means one joule of work can be done for every coulomb of charge that moves.

It's a measure of potential energy difference.

And how do you measure the true potential without you know draining the battery while you measure it?

Right to get the maximum possible voltage the true thermodynamic potential you need to measure it under conditions of virtually zero current.

Modern digital voltmeters are great for this because they have very high resistance and draw almost no current so they don't disturb the system or waste energy as heat.

Okay so we can build a cell and measure its voltage but how do chemists compare different reactions?

There must be thousands of possible redox pairs.

Is there a standard?

There is.

It was a major step forward.

You can't really measure the potential of just one half reaction in isolation.

So chemists needed a universal reference point.

Like sea level for measuring altitude?

Exactly like sea level.

They decided by convention to define the potential for the reduction of hydrogen ions to hydrogen gas.

That's two H plus plus two electrons giving H2 gas as exactly zero volt under standard conditions.

Standard conditions being?

One molar concentration for the H plus ions and one atmosphere pressure for the H2 gas.

Usually 25 degrees C.

This setup is called the standard hydrogen electrode or SHE.

So everything else is measured relative to that zero point?

Precisely.

By connecting any other half cell to the SHE we can measure the voltage and assign a standard reduction potential to that half reaction.

These are usually listed in tables always for the reduction process.

Okay so we have these tables of standard reduction potentials.

How do we use them to find the voltage of a cell made from two different half reactions?

It's quite straightforward.

You look up the standard reduction potentials for both half reactions.

The one with the more positive or less negative E degree value will be the one that gets reduced.

That's your cathode.

The other one has to be flipped to run as oxidation.

That's your anode.

And when you flip the reaction for the anode you flip the sign of its E degree value right?

Exactly.

You reverse the sign.

Then you just add the reduction potential of the cathode and the oxidation potential, the flipped one of the anode together.

Or maybe simpler, you can just take E cathode, E degrees anode, using the reduction potentials directly from the table for both.

And if the final standard cell potential E degree is positive?

If E to cell is positive the reaction is spontaneous as written under standard conditions.

It can run as a galvanic cell and produce electricity.

If it's negative, it's not spontaneous, it would need energy input to go.

Got it.

And there's that Ah yes, line notation.

It's a quick way to represent a cell.

You put the anode components on the left, cathode on the right.

A single vertical line shows a phase boundary, like solid, electrode, and solution.

And a double vertical line represents the salt bridge.

So anode, anode solution, cathode solution, cathode.

Very concise.

Okay, positive E degree cell means spontaneous.

This feels like it connects to thermodynamics, doesn't it?

Spontaneity sounds like Gibbs free energy.

You're absolutely right.

This is where it all ties together beautifully.

That cell potential E is directly related to the maximum amount of useful work the reaction can do.

And the maximum useful work a process can do at constant temperature and pressure is the change in Gibbs free energy.

So there's an equation linking them.

There is, and it's fundamental.

G equals nfe.

Okay, break that down.

Bragey is Gibbs free energy change.

Right.

N is the number of moles of electrons transferred in the balanced redox reaction.

F is the Faraday constant.

It's a big number, about 96 ,485 coulombs per mole of electrons.

It connects charge to moles.

And E is the cell potential.

And the minus sign.

The minus sign is key.

It shows that a positive cell potential, E zero, which we just said means spontaneous, corresponds to a negative UL.

And a negative UL is the thermodynamic criterion for a spontaneous process.

They perfectly agree.

So measuring voltage gives us a direct window into the reaction's thermodynamic drive.

Exactly.

It's an experimental way to determine E joule.

And it confirms, for example, why trying to dissolve gold in regular nitric acid doesn't work.

The calculated E degrees is negative, which means E degrees is positive, not spontaneous.

You need something much stronger, like aqua regia, to get gold to react.

Okay, standard conditions are neat for comparing things, but, you know, life rarely happens at exactly one molar in one atmosphere.

What happens to the voltage when concentrations change?

Great question.

Standard conditions are a baseline, but cell potential is actually very sensitive to concentration.

It does change.

How do we predict that change?

Using the Nernst equation, this brilliant equation lets us calculate the actual cell potential, E, under non -standard conditions.

It basically adjusts the standard potential, E degrees,

based on the current concentrations of reactants and products,

using the reaction

Q.

Q, right.

That's the ratio of products to reactants at any given moment.

Exactly.

The Nernst equation shows that if the concentration of products is high relative to reactants, Q1, the voltage will be lower than standard.

If reactant concentrations are high, Q1, the voltage will be higher.

So you could even make a battery just by having different concentrations of the same stuff in the two half cells.

Absolutely.

That's called a concentration cell.

The electrodes and solutions are the same chemical species, just at different concentrations.

Nature wants to equalize concentrations, right?

So electrons flow from the more dilute side, which acts as the anode, to the more concentrated side, the cathode, trying to level things out.

Until the concentrations become equal.

At which point the driving force disappears, Q equals one, or K at equilibrium,

the cell potential E becomes zero, and you have what we call a dead battery.

No more potential difference.

Wow.

But the sensitivity to concentration,

it's not just a complication, is it?

Can we use it?

We absolutely can.

It's incredibly useful.

It's the fundamental principle behind ion -selective electrodes.

Think of the glass electrode in a pH meter.

It measures H plus concentration by developing a potential across a thin glass membrane that depends on the H plus difference inside and out.

So it's like a tiny concentration cell measuring just one ion.

Essentially, yes.

And we have electrodes selected for fluoride ions, silver ions, calcium ions, used in water quality testing,

medical diagnostics to measure electrolytes all sorts of things.

They translate ion concentration directly into a measurable voltage.

That's pretty clever.

Turning a variable into a measurement tool.

And there's another connection.

Remember how E is zero at equilibrium, and Q becomes the equilibrium constant, K.

The Nernst equation lets us relate the standard cell potential, E degrees, directly to the equilibrium constant, K, for a redox reaction.

So we can calculate K from electrochemical measurements.

Another powerful link.

Okay, let's talk practical applications.

Batteries.

They're everywhere.

Our phones, cars, laptops.

They're just galvanic cells, right?

Fundamentally, yes.

A battery is just one or more galvanic cells packaged conveniently, sometimes connected in series to get higher voltage, designed for portable power.

What about the classic car battery, the lead acid one?

The lead storage battery.

An absolute workhorse, been around since 1915.

It uses lead metal as the anode and solid lead dioxide as the cathode, both sitting in sulfuric acid solution.

As it discharges, both electrodes get coated with lead sulfate, and the sulfuric acid gets used up, becoming more dilute.

The beauty is, it's rechargeable.

Running current back through it reverses the reaction, turning the lead sulfate back into lead and lead dioxide, and regenerating the sulfuric acid.

Amazing it's lasted so long as the standard.

It really is.

Despite being heavy and having environmental concerns with the lead, its ability to deliver high currents reliably for starting engines is hard to beat, even after a century.

Though you do have to be careful jump -starting, it can produce hydrogen and oxygen gas, which is explosive.

Right.

And then there are the modern ones, like in my phone.

Yeah, we have many others.

The common alkaline batteries improve versions of the older dry cells.

Then things like silver oxide, mercury cells, less common now due to mercury issues, and the big players today.

Rechargeable nickel cadmium, NiCad,

and especially lithium ion batteries.

Lithium ion seems to be king for portable electronics and EVs.

What's special there?

Lithium ion batteries are remarkable.

They have high energy density, meaning lots of power in a small package.

The chemistry is clever.

It involves lithium ions actually moving into and out of the crystal structure of the electrode materials.

It's called intercalation.

They don't really dissolve and reform the electrode in the same way as lead acid.

The ions just sort of shuttle back and forth.

That helps them last for many charge cycles.

Okay, so batteries store energy.

What about something that generates power continuously as long as you feed it, like a fuel cell?

Exactly.

A fuel cell is different from a battery.

It's still a galvanic cell, converting chemical energy to electrical energy.

But the reactants aren't stored inside.

They're continuously supplied from an external source.

Like hydrogen and oxygen?

That's the most famous example, yes.

Hydrogen gas fed to the anode, oxygen gas to the cathode.

They react electrochemically to form water and generate electricity in the process.

Very clean.

The only byproduct is water.

Where are these used?

The space program used them extensively efficient power and drinking water.

They're being developed for cars, buses, and stationary power generation.

Still some challenges with cost and hydrogen storage, but they hold a lot of promise for clean energy.

Research is also looking into fuel cells that could directly use fuels like methane.

So electrochemistry gives us power, but it can also take things away, right?

Like rust,

corrosion.

Ah yes, the unwanted galvanic cell.

Corrosion is basically nature trying to return metals to their more stable lower energy state.

Usually oxides or sulfides like how they're found in ores.

And it costs a fortune.

A huge amount.

Estimates suggest maybe a fifth of all iron and steel produced is just to replace stuff that rusted away.

So why do metals corrode, especially iron?

Most metals want to oxidize.

Their standard reduction potentials are lower, less positive than oxygen's.

So thermodynamically oxygen wants to take their electrons.

Some metals like aluminum or chromium form a very thin, tough, invisible oxide layer that sticks tightly and protects the metal underneath from further attack.

But iron isn't so lucky.

Unfortunately not.

Iron oxide rust is flaky and porous.

It doesn't seal the surface, so the corrosion continues.

And rusting is an electrochemical process.

You get tiny anodic and cathodic regions on the surface of the steel.

Like many galvanic cells.

Exactly.

At the anode spots, iron metal loses electrons.

8E2 plus 2E.

Those electrons travel through the metal itself to the cathode spots.

There they react with oxygen from the air and water to form hydroxide ions.

O2 plus 2H2O plus 4E plus 4OH.

And the iron ions meet the hydroxide ions.

The F2 plus ions migrate through any moisture on the surface, meet the hydroxide, often get further oxidized by oxygen to F3 plus rust,

and eventually precipitate as hydrated iron oxide F2O3 NH2O which is rust.

Often the rust forms at a spot different from where the iron initially dissolved.

And water is key.

And salt.

Absolutely essential.

Water provides the medium for ions to move a salt bridge.

And salt,

like which dramatically speeds up the electrochemical process.

So how do we fight it?

Besides just painting things?

Painting is one way of creating a barrier.

Plating with a more resistant metal like chromium or tin is another.

Then there's galvanizing coating steel with zinc.

Why zinc?

Zinc is more easily oxidized than iron.

It has more negative reduction potential.

Or more positive oxidation potential.

So if the coating gets scratched, the zinc acts as a sacrificial anode.

It corrodes instead of the iron.

Lever.

Sacrificing itself.

Similar idea is cathodic protection.

You electrically connect the steel structure, like a pipeline or a ship hole, to an even more active metal, like magnesium or aluminum blocks.

These act as the anode and corrode away, protecting the steel cathode.

You just have to replace the sacrificial anodes periodically.

Alloying also works stainless steel has chromium and nickel mixed in, which form that protective oxide layer.

Okay, so galvanic cells give us electricity from spontaneous reactions.

What about the opposite?

Using electricity to force a non -spontaneous reaction.

That's electrolysis.

And the setup is called an electrolytic cell.

You're essentially using an external power source, like a battery or power supply, to pump electrons in the wrong direction, thermodynamically speaking.

Assuring the reaction uphill.

Exactly.

You apply an external voltage that's greater than the cell's natural potential, if it had one.

This forces reduction to happen at the electrode connected to the negative terminal, not called the cathode, and oxidation of the electrode connected to the positive terminal, now the anode.

It reverses the spontaneous process, or drives one that wouldn't happen at all.

Like recharging a battery.

Recharging a battery is a perfect example of electrolysis.

You're forcing the discharge reaction to run backward.

Can we predict how much chemical change happens, like how much metal gets plated?

Yes, absolutely.

It's quantitative.

If you know the current in amperes, which are coulombs per second, and the time the current flows, you can calculate the total charge passed.

Charge equals current x time.

Okay, total charge.

Then you use the Faraday constant, f96 ,485 semole, to convert that charge into moles of electrons.

Got it.

Moles of electrons.

Finally, you look at the balanced half reaction for the process you're interested in, say, plating silver, Ag++Ea.

The stoichiometry tells you how many moles of electrons are needed per product.

So you can calculate the moles, and thus the mass of product formed.

Very useful for electroplating.

So we use this for things like splitting water into hydrogen and oxygen?

Yes.

Electrolysis of water is a classic example.

Or plating a thin layer of silver or gold onto jewelry or cutlery.

You mentioned something earlier, though, about surprises.

Sometimes the reaction you expect doesn't happen.

Ah, yes.

Overvoltage.

It's a bit of chemical reality intruding on simple predictions.

Sometimes, especially for reactions involving gases, you need a significantly higher voltage than the standard potential calculation suggests to actually make the reaction happen at a reasonable rate.

Why is that?

It's complex, often related to the kinetics, the energy barriers involved in forming gas bubbles on the electrode surface.

A key example is electrolyzing aqueous sodium chloride, brine.

Okay, NaCl in water.

You've got Na +, AshESi, and water molecules.

Right.

Based purely on standard potentials, water is easier to reduce to H2 and OH than Na +, ions are.

And water is slightly easier to oxidize to O2 and H +, than Cl ions are.

So you'd expect hydrogen at the cathode and oxygen at the anode?

You would.

But because of overvoltage, especially for oxygen formation on many common electrode materials, it turns out to be easier, requires less extra voltage,

to oxidize the chloride ions to chlorine gas, Cl2 instead.

So typically, electrolysis of brine gives hydrogen gas at the cathode and chlorine gas at the anode.

Little twist.

Interesting.

So theory gets you close, but experiment reveals these quirks.

Always.

And this quirk is hugely important industrially.

How so?

What are the big industrial uses of electrolysis?

Oh, massive.

Think about aluminum.

We use it everywhere, but it's incredibly reactive.

It's never found as a pure metal in nature, only as its oxide and bauxite ore.

And you can't just smelt it with carbon like iron.

Nope, aluminum oxide is too stable.

And electrolyzing in water doesn't work, as we said, water reduces first.

Melting pure aluminum oxide takes over 2000 degrees C, totally impractical.

So how do we get aluminum cans in foil?

Through the Hall -Herald process, discovered independently by Charles Hall in the U .S.

and Paul Herald in France in 1886.

A major breakthrough.

That was the secret.

They found that aluminum oxide dissolves in molten cryolite, N3LF6, another mineral.

This mixture melts at a much more manageable temperature, around 1000 degrees C.

You can then electrolyze this molten solution using carbon electrodes.

Aluminum metal is produced at the cathode, and oxygen reacts at the carbon anodes, producing CO2.

And that made aluminum cheap.

Dramatically.

It went from being a precious metal,

Napoleon III supposedly served his most honored guests using aluminum cutlery because it was rarer than gold, to the everyday lightweight material we know today.

A total game changer, all thanks to finding the right electrochemical conditions.

Amazing story.

What else relies on industrial electrolysis?

Lots of things.

Electrical finding is used to purify metals like copper.

You start with impure copper as the anode, pure copper as the cathode, in a copper sulfate solution.

Electrolysis dissolves copper from the impure anode and plates out ultra -pure copper onto the cathode, leaving impurities behind.

Essential for good electrical wiring.

And metal plating.

Like chrome bumpers.

Exactly.

Metal plating uses electrolysis to deposit a thin, protective, or decorative layer of one metal, like chromium, tin, nickel or zinc, onto another, usually steel.

Protects against corrosion.

Makes things look good.

And that brine electrolysis we talked about.

That's one of the largest scale electrolytic processes, often called the chloralkali process.

Electrolyzing molten NaCl in a down cell is how we produce sodium metal and chlorine gas.

But the aqueous version?

Electrolyzing aqueous NaCl brine, because of that overvoltage effect, gives us hydrogen gas, chlorine gas, and leaves behind sodium hydroxide, NaOH, in the solution.

NaOH, or caustic soda, is a hugely important industrial chemical, used in making soap, paper, textiles, and much more.

So controlling electricity lets us make fundamental chemicals.

Absolutely.

Historically, they used mercury cells for this, which caused environmental problems.

Now, more modern diaphragm or membrane cells are used, which separate the products more efficiently and avoid mercury contamination.

It's a continuous drive for efficiency and sustainability, driven by understanding electrochemistry.

Wow.

What a journey.

We went from electrons just swapping partners in redox reactions.

To figuring out how to harness that swapping galvanic cells to make electricity.

Developing standards like SGE to compare them, linking it all to thermodynamics with Gibbs free energy.

Accounting for real -world concentrations with the Nernst equation, leading to sensors.

Engineering practical power sources like batteries and fuel cells.

Understanding and fighting the destructive side in corrosion.

And then flipping it all around with electrolysis to force reactions and produce essential materials like aluminum and chemicals like chlorine and sodium hydroxide.

It really is the study of that fundamental back and forth between chemical energy and electrical energy.

It underpins so much from tiny sensors and medicine to giant industrial plants.

Understanding these principles lets you see how much intricate chemistry is quietly running our world.

It's incredible.

So the next time you plug something in, start your car or even just see a rusty nail, maybe pause for a second.

You're tapping into or witnessing this silent, powerful conversation between electrons and ions that chemists learn to understand and direct.

It makes you wonder what other hidden chemical conversations are out there, just waiting for us to tune in and take a deep dive.

And from the Last Minute Lecture Team, thank you for joining us today.

We really hope this deep dive helps solidify your understanding of electrochemistry.