Have you ever built a lemon battery? How do they work?
Have you ever built a lemon battery? How do they work? (0:32). We are distinguishing between two types of electrochemical cells: galvanic, also called voltaic cells, and electrolytic cells (1:05). Both cells have an anode, where the oxidation takes place, and a cathode, where the reduction occurs. (2:01). But where are the differences? Our episode describes the set-up of a galvanic cell (4:25) as well as electrolytic cell (5:09) and the function of all components.
Question: What metal (oxide) is used in common household batteries?
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Hi and welcome to the APsolute Recap: Chemistry Edition. Today’s episode will recap Galvanic and Electrolytic Cells.
Lets Zoom Out:
Unit 9 - Applications of Thermodynamics
Topic 9.7 - Galvanic (Voltaic) and Electrolytic Cells
Big idea - Energy
Introduction:
Have you ever built a lemon battery? It is one of those classic grade school or science fair experiments that we all should have done. No? Well, what are you waiting for! Get a lemon, a copper nail or penny, a galvanized nail, preferable with zinc, and some copper wires! Lemon batteries are an example of an electrochemical cell, more specifically a galvanic cell!
Let’s zoom in:
We need to distinguish between two types of electrochemical cells: galvanic, also called voltaic cells, and electrolytic cells.
Galvanic cells convert energy released by thermodynamically favored redox reactions to electrical energy, which we then can, for example, use to power a flashlight. Electrolytic cells, on the other hand, involve thermodynamically unfavored redox reactions. These types of cells convert electrical energy, which needs to be supplied externally, to chemical energy in order to initiate the chemical reaction. Both cells have an anode, where the oxidation takes place, and a cathode, where the reduction occurs. How can you remember this? Simple trick: Anode and oxidation both start with a vowel, cathode and reduction both with a consonant. Let’s take a closer look at the set-up.
A galvanic cell has two half-cells, one with an anode, the other one with a cathode. These two half-cells can, for example, be two beakers that are connected with a salt bridge and the flow of electrons is measured with a Voltmeter. Let’s make a specific example - sticking with zinc and copper from our lemon battery: We have Zinc metal as an anode in a solution of aqueous Zinc(II) sulfate. In the other beaker, we have copper as a cathode in an aqueous copper(II) sulfate solution. We now have two half-reactions: at the anode the zinc metal is oxidized and produces zinc(II) cations and two electrons which go into solution. Experimentally, we would be able to determine that by measuring the metal before and after - the mass will be less. These two electrons flow through the wires and the voltmeter to the cathode, our copper, where they are used in the reduction of copper(II) cations to elemental copper. If we measured the copper metal after the experiment, we would notice that it has gained mass. Our overall reaction therefore is: solid zinc, plus copper two plus, yields zinc two plus, plus solid copper.
You notice that while this reaction is occurring, the half cell with zinc loses two electrons and the half-cell with copper gains two electrons. To balance out the charges, we connect both cells with a salt bridge that is filled with an inert ionic solution, maybe with sodium chloride. The chloride anions flow to the anode, the zinc half-cell and the sodium cations flow to the cathode, the copper half-cell. Without the salt bridge, the cell wouldn’t work.
There can be some variations to this set up: We can replace the metal anode or cathode with an inert solid, like platinum, if we have gaseous or liquid reactant and/or products.
Electrolytic Cells also have anodes and cathodes, but these are in the same container, which means you don't need a salt bridge. What you will need, though, is a power source, because electrolytic cells contain reactions that are thermodynamically unfavored and therefore the chemical reaction needs electrical energy that can be converted to chemical energy. Equivalent to the galvanic cells, electrons flow from the anode to the cathode, but in this case via power source. The anode and cathode are submerged in an aqueous ionic solution or a molten ionic compound. Since we don’t have a salt bridge and the reaction is within the same container, the charges balance by a flow of cations to the cathode and anions to the anode.
When you look at images of galvanic cells and electrolytic cells, you often see the anode and cathode labelled with positive and negative. You will NOT have to label the electrodes with positive or negative! You will need to be able to identify where oxidation happens - at the anode, and where reduction happens - at the cathode, but not the sign of the electrode.
To recap:
Galvanic cells convert energy released by thermodynamically favored redox reactions to electrical energy. Electrolytic cells involve thermodynamically unfavored redox reactions and need an external power supply. Both electrochemical cells have an anode and a cathode and electrons flow from anode to cathode. Galvanic cells are usually set up with two half-cells and a salt bridge. Electrolytic cells are within one container and require an external power supply.
Coming up next on the APsolute RecAP Chemistry Edition: Electrolysis
Today’s Question of the day is about batteries.
Question: What metal (oxide) is used in common household batteries?