Electrochemical cell problem

[Lukraak_Sisser]

Maybe some other chemistry teachers here can help me out.

As part of the electrochemistry course my students are doing I want to show the ability to light a small (2.2V, .25A) lightbulb using an electrochemical cell.
This *should* be possible, after all, its just a battery.
Using copper in 0.1M Coppersulphate and zinc in 0.1M zinc chloride I can also clearly show that an electrode potential exists between the two electrodes, but only when a salt bridge is present to connect the two halfcells.
But for the life of me, I seem to be unable to get any current to actually flow trough an amp-meter let alone actually light a lightbulb.
So far its just been trials after school, but I would really like to actually get it to work before the students arrive at that point.
I'll definately be trolling the internet too, but maybe someone here has more experience with this.

 

[Myriad]

I can't address the electrochemistry directly (evaluating such things as the electrolyte concentrations and electrode geometry to determine what characteristics your cell should have), but from the point of view of straightforward electronics it's clear that your cell is simply not putting out enough power to light the bulb.

More specifically, while an electrical potential of 2.2V (or whatever your cell produces, you didn't say) might exist in an open-circuit conditions (no load), if the cell cannot produce enough current at that potential ("enough" in this case being the potential in volts divided by the resistance of the light bulb in ohms), the potential will drop. Try measuring the potential across the electrodes while the light bulb is connected, and you'll probably see it drop to close to zero. That indicates that your light bulb is way too large a load for your cell to handle.

The same problem would happen if you tried to start your car using eight AAA (or R03 or UM-4) cells wired in series. Even though the AAA battery pack would have the same 12V potential as a car battery, the AAA cells cannot source nearly enough current. Instead, the AAA's would be essentially shorted out through the coils of the starter motor, without getting the motor to crank, and would read close to zero volts while connected.

And in terms of power output (and total energy stored), a modern AAA battery is as superior to your average beakers-of-chemicals-and-bits-of-metal-foil lab demonstration battery as a car battery is to a AAA battery.

The only thing unexpected in your account is not measuring any current. In a short-circuit condition (just the cells, with the salt bridge in place, and the ammeter connected directly between the two electrodes), there should be a measurable current, even if only a few milliamps. Is the scale of the ammeter set correctly? Is it set to DC and sensitive enough to measure a few milliamps? Are you sure the polarity is correct? Try making the same measurements in the same way, but using a small store-bought battery, to make sure you're using the instrument correctly.

If you continue to see the cell generate a potential but no measurable current, it might be that your salt bridge and/or your electrodes just have very high resistances. Skinny wire electrodes without much surface area, a narrow salt bridge, or a low ion concentration in any portion could all contribute to high internal resistance in your cell, resulting in low measured current.

A lower power device such as a small single LED or a small LCD digital clock might make for an easier demonstration than a light bulb that requires more than half a Watt.

Respectfully,
Myriad

 

[technoextreme]

As part of the electrochemistry course my students are doing I want to show the ability to light a small (2.2V, .25A) lightbulb using an electrochemical cell.

That is a huge light bulb for what you are trying to do.

 

[Olowkow]

I agree with all of the above.

250 ma is a much too heavy current load for such a cell. The idea of using an LCD watch (cheap) is a good way to show the principle involved. One of my customers (prof) was using the "potato battery" and was trying to light an LED, NO WAY. Same problem essentially. I got her an LCD watch, and of course it worked, but I was unable to solder to the metal that was used, so it was kind of tricky to mechanically install wires. If you look on line, you can find the whole kit, very cheap.

 

[Lukraak_Sisser]

Thanks for the help.
I'll see if I can get such an LCD watch or something.
But I still need to fix the current problem.
Using Magnesium in water and platinum in dilute sulfuric acid I can get up to 2.1 V, but the amp meter doesnt show even a micro amp of current.
I've tested the meter, and it works for direct current on a AAA battery. It has a set of ranges, but as I said, on the electrocell even the most sensitive setting shows nothing.
THe direction of the current is not important to the machine, it can detect it going both ways and just adds a minus sign if its reversed.
The only thing I can currently think of is that there is a resistance somewhere in the wires or connections I use. They are very old and abused.

 

[Andrew Wiggin]

Not a chemistry teacher, but I enjoy tinkering with things like this.

Try zinc in zinc sulfate (or dilute sulfuric acid) and copper in copper sulfate for yout half cells. Daniel's cell is known to work, and with the exception of using zinc chloride, that's what you've got.

Make sure there's plenty of salt in your salt bridge, and plenty of copper sulfate and zinc sulfate in your solutions to reduce internal cell resistance. Daniel's cell is known already for its high internal resistance and you don't want to add to that. IIRC the antique cells operated at close to saturation; I know there was solid 'bluestone' at the bottom of the copper cell.

Make your electrodes large, as current will be proportional to surface area. The 'crowsfoot' style of electrode was used in the originals, where multiple strips of copper would be riveted together and then fanned out, or the zinc cast into a shape with multiple fanned out rods. If you're using the standard one inch by four inch or so metal strips that I see in chemistry education sets, where maybe three square inches is in the cell, you won't get much current at all.

This can also be constructed as a gravity cell, with no salt bridge needed; just make it before the class and leave it short circuited and stationary to keep it from mixing. IIRC you can do this without zinc sulfate or sulfuric acid; put the copper electrode in the bottom of the jar (it needs an insulated lead), pour copper sulfate around it, hang the zinc electrode at the top of the jar, and then pour distilled water in until the top electrode is covered. Leave it short circuited overnight; it'll start very slowly until enough sulfate is freed up in solution.

You'll also want to make sure you've got a lamp that will light with less than one volt or run several cells in series, as this sort of cell was usually rated as providing one volt. IIRC, the calculated potential is around 1.1 volts, and as I mentioned there's a pretty high internal resistance.

Hope some of that helps.

 

[John Hewitt]

Maybe some other chemistry teachers here can help me out.

As part of the electrochemistry course my students are doing I want to show the ability to light a small (2.2V, .25A) lightbulb using an electrochemical cell.
This *should* be possible, after all, its just a battery.
Using copper in 0.1M Coppersulphate and zinc in 0.1M zinc chloride I can also clearly show that an electrode potential exists between the two electrodes, but only when a salt bridge is present to connect the two halfcells.
But for the life of me, I seem to be unable to get any current to actually flow trough an amp-meter let alone actually light a lightbulb.
So far its just been trials after school, but I would really like to actually get it to work before the students arrive at that point.
I'll definately be trolling the internet too, but maybe someone here has more experience with this.

The traditional way of lighting a bulb in a school chemistry class is with one beaker of dilute sulphuric acid with plate electrodes of copper and zinc dipped into the same beaker. The two electrodes must not touch one another and are connected via the light bulb.
This works because, although zinc is reactive enough to react with acid, the cathodic reaction that would produce the hydrogen is kinetically very slow. (I think the cathodic reaction is the one that produces hydrogen, but correct me if I'm wrong.) Copper is not a reactive enough metal to react with the acid but, on that electrode surface, the cathodic reaction is fast. Hence, for the zinc to react with the acid, the electrons it produces are obliged to flow through the metal wire to the copper where they react with H+ to form hydrogen gas. If a bulb is in the way, it lights.

Because there is no salt bridge involved, the internal resistance of this cell is fairly low so that enough current can flow to light a small bulb.

(The same properties of these metals is the basis for the well known trick whereby you see hydrogen gas coming from copper wire but not the zinc if it is brought into contact with a piece of zinc in sulphuric acid, even though it is the zinc that is reacting with the acid - try it.)