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  • Aaah, the controlled flow of electrons, making possible laptops and phones and cars and pacemakers.

  • Batteries, just like everything else in life, is just chemistry raised to the power of awesome.

  • The kind of chemistry that happens inside of a battery is called electrochemistry

  • because it involves reactions that produce or consume free electrons.

  • Specifically, they are oxidation or redox reactions, the ones where electrons are exchanged.

  • I've told you about redox reactions before and if you haven't seen that episode yet,

  • you should probably go watch that before you watch this.

  • Don't worry, I will still be here when you get back.

  • Now, when the flow of electrons in these kinds of reactions are sent through a conductor,

  • like a piece of metal, it can be used to do all sorts of work.

  • Like, for example, this kind of work.

  • The amount of work that can be done depends on how strong the push or pull on electrons is between the two reactants.

  • This is the reaction's electrical potential, but to its friends, it's known simply as voltage.

  • Basically, if the voltage is high, each electron can do a lot more work than if the voltage is low.

  • Many of the wonderful things in our modern lives are based on one simple premise:

  • putting a device between the two halves of just such a reaction;

  • the half that donates electrons, and the half that accepts them.

  • By harnessing that energy, a lot of the coolest things you've done today,

  • up to and including watching this episode of Crash Course Chemistry, has been made possible.

  • [Theme Music]

  • Part of what makes redox reactions so powerful, and powerfully excellent, is that they are complicated.

  • Because in each reaction, there's at least two things going on:

  • there's the part of the reaction where the electrons are being released

  • and another part where they're being eagerly demanded.

  • So when we deal with electrochemistry, we usually think of reactions in terms of half reactions.

  • Let's start with a typical redox reaction that happens in this alkaline battery as an example.

  • In here elemental zinc is going to react with manganese dioxide, also known as manganese

  • four oxide, to produce manganese three oxide and zinc oxide.

  • When you break this down in to half reactions,

  • first you have elemental zinc with an oxidation number of 0 being oxidized to zinc 2 ion.

  • At the same time manganese 4 is being reduced to manganese 3.

  • When we balance the half reactions,

  • we see that two electrons are released during the oxidation of each zinc atom,

  • and one electron is consumed by each manganese 4 atom.

  • The water and hydroxide ions, by the way, come from a solution of potassium hydroxide.

  • Which is a basic, or alkaline compound.

  • Which is why we call these things alkaline batteries.

  • Now if each of these half reactions occurred in contact with the other one,

  • they'd just spontaneously go to equilibrium releasing energy as a bunch of heat which wouldn't be very helpful.

  • So batteries are designed to harness that energy by isolating the half reactions from each other.

  • This allows excess electrons to build up in the negative terminal, called the cathode,

  • while an electron vacuum of sorts occurs in the positive terminal, the anode.

  • Electrons can then cross from one half reaction to the other,

  • only when we connect the cathode and the anode of the battery via conductors.

  • So the current can be used to do work. Which I can do by licking this 9 volt battery.

  • Ahhh! [laughs]

  • In these batteries, the zinc is in the center surrounded by a layer of cellulose that allows ions to pass through.

  • The manganese oxide is in the outer layer that surrounds the zinc core,

  • but the cellulose barrier doesn't allow the zinc and the manganese to mix.

  • Alkaline batteries are a type of galvanic cell.

  • Which is generally defined as an apparatus that generates electrical energy from a redox reaction.

  • Here's another example of a galvanic cell,

  • one where the interesting part is the flow of whole ions instead of the flow of electrons.

  • In this case, wires connect metal rods that are suspended in the solution.

  • They're are the anode and cathode here.

  • What's noteworthy, is that the metal atoms are actually consumed.

  • They're used up from the anode rod as they're oxidized, and the metal slowly wears away.

  • Meanwhile, the opposite happens at the cathode rod, where metal ions from the solution gain electrons

  • and precipitate on to the cathode as pure metal, gradually growing larger.

  • The circuit is completed by the wire, but also by a salt bridge,

  • which is also a U-shaped tube that contains a salt solution

  • that allows the metal ions to go from the anode to the cathode.

  • So, now you know how batteries work: developing and transferring a charge using electrochemistry.

  • But before any redox reaction can be used to text your boyfriend or girlfriend or whatever,

  • we need to know how much voltage it can generate.

  • Fortunately all of those amazing chemists who have come before us have done a lot of the work once again.

  • The voltage generated by many half reactions is already known and can be found in most textbooks or online.

  • And as I mentioned earlier, voltage is really just a way of expressing the electrical potential of each half reaction.

  • The difference between the chemical demand for the electrons in one half,

  • versus the tendency to lose them in the other.

  • These measurements are done at standard conditions, which we discussed in the enthalpy and entropy episodes.

  • And by convention they are written as if the substance is being reduced, not oxidized.

  • For this reason the value is known as the standard reduction potential of a substance.

  • To see how reduction potentials work in half reactions and combine in an overall reaction,

  • let's consider a galvanic cell where zinc is oxidized and copper ions are reduced.

  • Keep in mind that potentials are determined at standard state:

  • 25 degrees Celsius and 1 molar solutions of the copper and zinc ions.

  • Our cell needs to be set up under the same conditions or the voltage will be a bit different than expected.

  • The half reactions show more clearly what the electrons are doing.

  • We can see that the zinc is oxidized and the copper is reduced.

  • Now all standard reduction potentials are measured relative to the reduction of hydrogen ions to hydrogen gas,

  • which is set at zero, just as a baseline.

  • When copper is reduced, for example, it generates 0.34 volts more than hydrogen does,

  • so we say its standard reduction potential is +0.34 volts.

  • The standard reduction potential for zinc is -0.76 volts,

  • but because zinc is oxidized in this reaction, we can't use the reduction potential directly.

  • Instead, as a general rule, when you convert a reduction half reaction to oxidation,

  • the sign of the voltage is simply reversed.

  • So, the -0.76 volts for the reduction potential of zinc

  • becomes +0.76 volts for the oxidation potential of zinc ions.

  • And the electrical potential for the whole reaction, called the standard cell potential,

  • is just the sum of the standard potentials of both half reactions.

  • In this case, that would be 1.1 volts.

  • Now I gotta point out here that the electrical potential of a redox reaction is related to its equilibrium constant.

  • There's actually a way to determine the equilibrium constant from a measured voltage, and vice versa.

  • Both of these constants have a lot to do with the energy the reaction can release, or its Gibbs free energy.

  • But in brief, the higher the voltage, the more electrical energy can be produced.

  • So if the voltage is positive, it means that under standard state conditions, the reaction will spontaneously go forward.

  • If the sign is negative, the reaction will proceed backward.

  • And it totally makes sense, when you think about it, that reactions like this are used to make battery cells.

  • The reactions in batteries need to be spontaneous

  • because their whole purpose is to release energy, not consume it.

  • So what if we don't want to power a phone or a laptop or a toy shuttle.

  • What if instead we want to plate an iron car bumper with chrome?

  • This cannot be done with a spontaneous reaction.

  • Instead, a different electrochemical process is needed, one that you've probably heard of: electroplating.

  • This is done by immersing an object in a solution that contains an excess of ions of the coating metal.

  • A bar of the coating metal, in this case chromium, is used as the anode,

  • and the item to be plated, the iron bumper, acts as the cathode

  • When an electric current is applied, a redox reaction occurs in the solution

  • and atoms of the coating metal are deposited on the cathode.

  • This is essentially the opposite of a galvanic cell, known as an electrolytic cell,

  • and it performs electrolysis, which uses electricity, electro-, to do the breaking apart, -lysis.

  • In this case molecules in the solution are being broken down so that the metal atoms can be deposited on the surface.

  • Electrolysis is used for lots of other things too, like coating jewelry or flatware with gold or silver,

  • refining metals or separating mixtures of metal ions.

  • Also, converting water into hydrogen gas and oxygen.

  • So if you didn't already understand how much impact chemistry has in our daily lives,

  • you certainly should be able to see it now.

  • Not only are all the materials around you made of chemicals,

  • but even the electrical devices that power our lives depend on the reactions of electrochemistry.

  • Thank you for watching this episode of Crash Course Chemistry.

  • If you were listening, you learned that electrochemical reactions are redox reactions

  • that we describe in terms of half reactions.

  • You learned how an alkaline battery works and what's inside of it,

  • what a galvanic cell is and how it can be set up

  • and how to calculate the voltage that can be generated by a half reaction, the standard reduction potential,

  • and by the overall reaction, the standard cell potential.

  • You also learned how electrolysis and electroplating work.

  • This episode was written by Edi Gonzalez and edited by Blake de Pastino.

  • Our chemistry consultant was Dr. Heiko Langner. It was filmed, edited and directed by Nicolas Jenkins.

  • Michael Aranda is our sound designer and Thought Cafe is our graphics team.

Aaah, the controlled flow of electrons, making possible laptops and phones and cars and pacemakers.

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電化學。化學速成班#36 (Electrochemistry: Crash Course Chemistry #36)

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    songwen8778 發佈於 2021 年 01 月 14 日
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