Are there any Chemistry buffs in here?

Factual Questions
1

Hello all!

I’m a student of the humanities who is way out of his element (pardon the pun) when it comes to science. I’m wondering if someone could give me the straight dope on electrolysis.

Specifically, I’m trying to understand the way in which batteries work. What makes one concoction good for the inside of a battery, and another one not? Is it simply its ability to store energy? Why, then, are their AAs made out nickel cadmium and not glucose, or propane? I get the feeling that I’m confusing an electrolyte with a fuel, and anyone that could steer me in the right direction would be my hero.

2

It’s not a simple subject.
Why not start with Wikipedia, and post back with questions.

3

Thanks for responding. Actually I looked at Wikipedia, How Stuff works, and “The complete idiot’s guide to chemistry” last night, all of which I have next to me. It seems I’m quite the idiot! :frowning:

Ok, this is my first road block, from the howstuffworks article:

Ok, what is said acidic paste? Is the paste itself a source of energy, or does it merely store an electric charge from somewhere else? Is it a fuel? Could I ostensibly (barring safety issues) burn it to give off heat?

4

OK, I’m not a chemist, so I might get some of this wrong, but:
The paste primarily serves as a conduit for ions moving from one electrode to the other. The energy in the battery is derived from the difference in electronegativity of the two electrodes. That said, there are chemical reactions going on between the electrodes and the electrolyte, so it is possible to “use up” the electrolyte, and to be able to renew a battery by just replacing the electrolyte. Here is some more in depth information on different batteries: Battery Chemistry Tutorial and FAQ from PowerStream describes the chemical properties of common and uncommon electrochemical cells

5

The acidic paste is just an electrolyte, a conductive medium that facilitates the transfer of charged ions. In a zinc-carbon battery, the electrolyte is typically ammonium chloride mixed with a bit of water.

Batteries function by exploiting redox reactions. Two reagents are separated by an electrolyte. One reagent in the reaction gives up electrons to the other via the electrolyte. This net transfer of charge results in an electromotive force that drives an electrical current.

6

Ok! I feel like I’m getting somewhere! :slight_smile:

So, if I understand so far, the paste itself is not a source of energy, but only its storage medium. What is the source of the energy then? Simply placing an electric current under through the paste to give it a few valent electrons?

I googled that redox reaction part. If I understand it right, this is why batteries start to rust when they lose their charge, yes?

7

Bah.

One reagent in the reaction gives up charge to the other via the electrolyte.

8

I am an organic chemist by training and profession. That said, this sort of thing is taught in a general manner in the undergraduate curriculum, so I know a little bit.

There are gigantic tables of eV (electron-volt) values for ions and changes in ionization. The back of my P. Chem. book has one that’s several pages long. The energy is quite literally chemical energy. You get electricity out of a battery because you are exploiting the energy given off in a redox reaction. The paste, as mentioned earlier, is just a conductor. For instance, in an alkaline battery (which is zinc and manganese dioxide) the half-reactions are:

Zn (s) + 2OH− (aq) → ZnO (s) + H2O (l) + 2e−
2MnO2 (s) + H2O (l) + 2e− →Mn2O3 (s) + 2OH− (aq)

In other words, zinc metal reacts with hydroxide to give zinc oxide, water, and two electrons. The manganese (IV) oxide then reacts with the water and the two electrons to give manganese (III) oxide and hydroxide. So the zinc is oxidized and the manganese reduced. This means that zinc is the anode and the manganese (III) oxide is the cathode. The electrons migrate from the anode to the cathode through the KOH paste. Based on those half-reactions, I could tell you what the expected voltage would be by looking up the values in a reference book.

9

The energy sources are the two chemical reagents themselves. Using the example of a typical alkaline battery, the two reagents are zinc and manganese(IV) oxide. Due to its electron configuration, the manganese in the manganese(IV) oxide attracts electrons more strongly than the zinc. The net redox reaction results in a transfer of negative charge from the zinc to the manganese.

Batteries rust because the casing is made of steel and the steel rusts in moist air. Rust forms as a result of a redox reaction between iron and oxygen, but it is not the redox reaction driving the battery.

10

Ok, so the two reagents (in this example, Zinc and Manganese) are the source of energy, in that when you cajigger them together, they form an exothermic reaction? And the electrolyte keeps these valent electrons from dissipating away?

Is this reaction continuously occurring? How does a battery (to anthropomorphize) “know” it’s being used?

11

Electrical engineer checking in. Once the voltage on the plates builds up to the potentials associated with the chemical reactions, it stops the reactions. In order for the reaction to start happening again, you have to reduce the voltage on the plates, and this happens when you connect the battery to a load, and transfer electrons from the negative plate to the positive. If the reaction rate is extremely sensitive to plate voltage, then the battery is said to have a low internal resistance.

That said, the electrolyte typically acts as a bit of a load on it’s own…it leaks a little charge the “wrong” way. This is known as self discharge and battery manufacturers work to reduce it. Extremely low self discharge is a hallmark of many lithium based chemistries.

If instead, you connect a power supply, and raise the voltage above what the associated chemestry “wants” then you can usually reverse the reactions, and charge the battery. Some batteries don’t behave well when you do this though.

12

I’m bumping an old thread but now I have some new questions regarding this!

I remember seeing something like those equations you posted in my high school chemistry class once upon a day. Without actually going through a lengthy survey course in chemistry, is there any fast and messy reference tool I could use to deduce this? And by this I mean “Compound X and Compound Y release ‘z’ number of valent electrons via a ‘redox reaction’ to produce power.”

If it’s not too much to ask, can you recommend some follow up reading that won’t bog me down in jargon? I can buy a book or make a trip to the university library, that’s not a big problem!

Oh, and thanks for your thorough response back in the day. :slight_smile:

13

You want a table of values for an electrochemical series. Most first year chemistry textbooks will have an appendix with such a table. You can probably search online for more extensive tables.

14

Thanks!

15

Voltage created by various reactions (not the true voltages, but instead relative to hydrogen reaction)
http://en.wikipedia.org/wiki/Standard_electrode_potential_%28data_page%29

In the above table, the Hydrogen reaction actually produces around 4.4V, so in order to find the true absolute potential for each reaction, we must add 4.4V to all the listed potentials. For metal/electrolyte reactions, the metal plate is charged negative and the electrolyte becomes positive.

Really basic battery description: whenever metal touches electrolyte, for a moment the electrolyte dissolves the metal extremely rapidly. However, the dissolving metal atoms don’t bring any of the metal’s “electron sea” along with them as they’re torn free from the metal surface. Hence the dissolved metal atoms in solution will all be positively charged metal ions lacking one, two, or three electrons. This process of dissolving is creating a “charge pump,” a pumping action which quickly makes the metal negative and the electrolyte positive. The charge-pumping action is the source of the electrical energy sent out by a battery. To power the charge-pump, one of the battery’s plates must corrode (a chemical reaction, think of it as a slow form of burning.) But there’s more.

When metal first touches electrolyte, the charge pump action only runs briefly until a voltage appears between metal and electrolyte, and an intense electrostatic field appears in the atom-sized gap between the metal and electrolyte. (Gap is called the Helmholtz Double Layer.) This e-field halts the dissolution, since positive metal ions will be repelled by the electrolyte and attracted by the metal surface. Rather than dissolving like a lump of sugar, the metal charges up to a few volts negative, and afterwards sits quietly in contact with the electrolyte.

If we could somehow reduce the voltage between metal and electrolyte, we could turn the corrosion reaction back on again, and harvest electrical energy. The easy way to do this is to place two different metals in contact with the same blob of electrolyte. Each metal/electrolyte voltage (or “half-cell potential”) will be different. The voltage from electrolyte to metal1 might be lower, and the voltage from electrolyte to metal2 might be higher. If we electrically connect the two metals, their metal/electrolyte voltages are forced to the same value. The voltage between electrolyte and metal2 has been lowered, so metal2 begins dissolving rapidly, and a relatively huge electric current appears: the path of charge flow is through the metal-metal contact, but also through each of the two metal/electrolyte contacts. The charge-pumps on the two metal surfaces are connected back-to-back and in opposite directions, and the stronger pump wins. (Actually, their oppositely-directed voltages will subtract, and the difference determines the overall charge flow direction.) The weaker pump is forced to flow backwards: the voltage between metal1 and electrolyte has been forced to be higher than normal. Any reactions at the metal1 surface are forced to run backwards by the more energetic corrosion-reaction at the surface of metal2.

Note that a battery with terminals connected together is a “shorted” battery. The reactions run fast, and one plate dissolves rapidly with much heat output. In real world circuitry the currents are designed to be hundreds of times smaller than the short-circuit current. In other words, when a battery powers some electrical device, the “electrical friction” or resistance presented by that device will greatly slow down the charge flow and the chemical reactions at the electrode surfaces.

Typical misconceptions: there is no current in the electrolyte? (Wrong, in fact the path for current is always a circle, so if we have one amp at the battery terminals, there must be one amp going through the electrolyte.) Another misconception: the carbon rod in a Zinc flashlight battery is an electrode plate (no, actually the rod is a low-resistance current spreader. Instead, the cylinder of compressed carbon/manganese-oxide powder is one electrode, the zinc can is the other.)