If you're seeing this message, it means we're having trouble loading external resources on our website.

If you're behind a web filter, please make sure that the domains ***.kastatic.org** and ***.kasandbox.org** are unblocked.

Main content

Current time:0:00Total duration:9:10

- [Voiceover] Here we have a table of Standard Reduction Potentials, and this is a shortened version, but you can see on the left side, we have different half-reactions. All of these half-reactions are written as reduction half-reactions. Remember gain of electrons is reduction. If we look at our first half-reaction, we have silver ion gaining an electron to form a solid silver. That's a reduction half-reaction. The standard reduction potential turns out to be +.80 volts. That's compared to this
half-reaction down here, which corresponds to the
standard hydrogen electrode, which is the reference value. This has a potential of zero volts. All of our other half-reactions
are compared to this one. The more positive the value is for the standard reduction potential, the more likely the
substance is to be reduced. Let's compare the
reduction of copper 2+ ions to the reduction of zinc 2+ ions. Let's compare these two half-reactions. If we are reducing copper
2+ to solid copper, the standard reduction
potential is +.34 volts. If we are reducing zinc 2+ to solid zinc, the standard reduction potential turns out to be -.76 volts. The more positive value, the more likely the
substance is to be reduced, so obviously +.34 is
more positive than -.76. We know that this must be
the reduction half-reaction. If we're talking about a redox reaction involving copper and zinc, this must be our reduction half-reaction. Let's go ahead and write that. This is our reduction half-reaction where we have copper 2+ ions gaining two electrons to turn into solid copper. The standard reduction
potential is +.34 volts. This is equal to +.34 volts. We know in a redox reaction something is reduced and
something is oxidized, and since we already have
our reduction half-reaction, we need an oxidation half-reaction. That must mean that this should be our oxidation half-reaction, but here we have it written as a reduction half-reaction, so we need to reverse this reaction to show it as an oxidation. We need to start with solid
zinc on the left side. We start with solid zinc on the left side, and zinc is oxidized into zinc 2+ ions, and we're losing two electrons. Remember loss of electrons is oxidation. Now this is an oxidation half-reaction. We need to find the
standard oxidation potential for this half-reaction. We can do that by looking
at our table here. So -.76 is the standard
reduction potential. Since we reversed our half-reaction, we just need to change the sign. The oxidation potential must be +.76. All we need to do is reverse the sign to get our standard oxidation potential, so we get +.76. To find our overall redox reaction, we just need to add together
our two half-reactions. To find our overall redox reaction here, we add the reduction half-reaction and the oxidation half-reaction. These electrons would cancel out, and on the left sides we
would get copper 2+ ions. This would be copper
2+ ions and solid zinc. On the right side for our products, we would get solid copper
and zinc 2+ ions in solution. We get solid copper and zinc 2+ ions. This overall redox reaction
should look very familiar to you because this is the
spontaneous redox reaction that we've talked about
in the last several videos as our example of a
voltaic cell, all right? We know this is a spontaneous reaction. How do we find the potential
for the cell, all right? How do we find the
standard cell potential? How do we find the potential
for the entire cell? To find the overall reaction, we add together our
reduction half-reaction and our oxidation half-reaction. That gave us our overall reaction. To find our standard cell potential, we just need to add together our reduction potential
for the half-reaction and the oxidation potential for
the oxidation half-reaction. To find the potential for the cell, we add the reduction potential and the oxidation potential. We get when we do that, we're gonna get +.34 volts is the potential for the
reduction half-reaction, and +.76 volts is the potential for the oxidation half-reaction. That gives us our standard cell potential. For our cell the potential
is equal to +1.10 volts, which we already know this
from previous videos, right? I talked about the fact
that you can use a voltmeter to measure the potential difference, to measure the voltage of a voltaic cell. You're gonna get a +1.10 volts under standard conditions. That's one of the nice things about the standard reduction potential table. We can calculate the voltage of our voltaic cells this way. Let's look in more detail
at our half-reactions. Let's start with the
oxidation half-reaction. We know that zinc is
being oxidized, right? Zinc is losing 2 electrons, and those two electrons that zinc loses are the same two electrons that caused the reduction of copper. Zinc is the agent for
the reduction of copper. We say that zinc is the reducing agent. Sometimes students find this confusing because zinc is being oxidized, so why is it the reducing agent? Zinc is the agent for the
reduction of something else, in this case, copper 2+ ions. So zinc is the reducing agent. Copper 2+ is gaining those two electrons, so copper 2+ is being reduced, but because copper 2+ is
gaining those two electrons, it allows zinc to be oxidized. Copper 2+ is the agent
for the oxidation of zinc. Copper 2+ is our oxidizing agent for our redox reaction. Let's look at our standard
reduction potential table and let's see if that
can help us understand oxidizing agents and reducing agents. We've been comparing these
two half-reactions, right? These two half-reactions right here. Let's compare copper 2+ ions to zinc 2+ ions, right? Copper 2+ we know is more
easily reduced, right? It has the higher, has the more positive value, I should say, for the standard reduction potential. Copper 2+ is more easily reduced, and therefore, copper 2+ is a stronger oxidizing
agent than zinc 2+. As you go up on your
standard reduction potential, you're increasing in the tendency for something to be reduced, and therefore, you're
increasing the strength as an oxidizing agent. As you move up on your
standard reduction potential, increased strength as an oxidizing agent. Copper 2+ is a stronger
oxidizing agent than zinc 2+. All right, let's think about the opposite. As you move down on your
reduction potentials, as you move down here, let's compare solid copper
and solid zinc, right? We know that solid zinc was our reducing agent in our reaction, and that's because the reduction potential was the more negative one. That means this is more likely to be the oxidation half-reaction. As you move down on your
reduction potential, you have an increasing
tendency to be oxidized. Therefore, you have an increasing strength as a reducing agent. Zinc, right, is a stronger
reducing agent than copper because, again, looking at
the reduction potentials, you know that it's more
likely to be oxidized. Going down on your reduction potentials, increased tendency to be oxidized, therefore, increased strength as a reducing agent. If you look at lithium, right? Lithium here is even more negative for the reduction potential. Therefore, this is even more likely to be an oxidation half-reaction. Lithium is a stronger
reducing agent than zinc.

AP® is a registered trademark of the College Board, which has not reviewed this resource.