Chemistry: The Activity Series of the Elements

This lesson has not been reviewed.

Please purchase the lesson to review.

The beaker on your left originally contained a solution of silver nitrate. We put a copper wire in it and waited a little while. The solution turned this pretty blue color. At the same time the copper wire got coated with a fine metallic powder, corresponding to silver metal. We actually did this demo and discussed it a little earlier and you may want to review that reaction if this is unfamiliar to you. What happened in here is an example of a redox reaction. Copper metal reacted with silver as an oxidant to give us copper 2+, plus silver metal. Again, the copper is our reducing agent; the silver is our oxidizing agent in this reaction.
Let me ask you something. Suppose that I remove the copper wire and put this silver spoon in my solution, which I know now contains copper 2+. What do you think is going to happen? Will the following reaction occur? Will silver metal react with copper 2+, give electrons to the copper and generate copper 0 and silver +? Well, we can wait around all day if we want, but nothing is going to happen in this reaction. That is because the reverse reaction is not spontaneous. The reason it is not spontaneous is because the copper is a better reducing agent than the silver metal is a reducing agent. In other words, the copper 2+ is a poorer oxidant than the silver + is an oxidant. You'll note that there is his complimentary relationship between the left and the right side of our reaction. The stronger the reducing agent copper is, in fact, the weaker it is as an oxidant. Correspondingly so for the silver--the stronger silver is as an oxidant, the weaker its corresponding reduced product is.
Now, let me ask you another question. We'll put our copper wire back. In this beaker, I have a strip of zinc metal. I've put that zinc metal, once again, into a solution of silver nitrate. I have silver + in solution. This reaction also occurs, just like is the case with the copper. I end up with silver precipitated on my zinc here. This reaction is spontaneous in this direction. Once again, zinc is a good reducing agent, much better than silver metal is as a reducing agent. Silver is our good oxidant, much better than zinc 2+ is as an oxidant. So this reaction is spontaneous from left to right. Again, the reverse reaction would not be spontaneous.
Here's a question. Suppose that after generating copper 2+, after we start with copper metal, and we rip the electrons out of the copper with silver, can copper, in turn, come back and take electrons from zinc if we gave the zinc metal to the copper 2+? In other words, would the following reaction occur? Would copper 2+ react with zinc metal to give us copper metal and zinc 2+ or if not, would the reverse reaction, then, occur? Maybe it would be possible for copper metal to react with the zinc 2+ solution to give us copper and zinc metal.
Wouldn't it be great if we had some resource, some table that we could go to, some series perhaps, that would give us relative orders of reactivity based on the ability of a metal to act as a reducing agent or alternatively, based on the cation to act as the oxidant? Well, you guessed it. There is such a series. It's called the Redox Activity Series, because there are actually different forms of activity series, some dealing with acids and bases, and with this particular example dealing with the ability of species to act as redox partners. Let's focus on the redox series for a moment. What I'm showing you here is an abridged version of the full series. We'll eventually look at the entire series as we get into electro chemistry, but right now, let's just focus on this series alone. Each one of these half-reactions is associated, actually, with a number. Eventually we'll be able to quantify this relationship. Right now, I'd just like to introduce the series to you as a qualitative listing of different metals and their abilities to act as reducing agents and correspondingly of their cations to acts as oxidants.
How do we read this chart? These are listed according to oxidation power. That means that as I go from potassium to calcium to sodium, magnesium, down through, let's say to copper and to silver, I'm going from a weaker oxidant to a stronger oxidant. Remember that complimentary relationship that we spoke about earlier? Since it is going from weaker to stronger oxidant, then the reverse is going to be true for this side. In other words, as I go from silver to copper to iron, I'm going to a stronger reducing agent. Let me say that again because that's such an important idea. As I go down, I go from very, very bad oxidants, to much better oxidants for the cations. I also go from very good reducing agents for the metallic forms, to very poor reducing agents. Remember, we talked about silver and we talked about copper, just as an example. Copper 2+ is not as good of an oxidant as silver+ is as an oxidant and so we'd expect, then, for silver to react with copper in order to give us silver metal. Just the fact that the silver half-reaction is below that of copper, we expect for the silver to be able to get the electron. Once again, silver+ is a better oxidant than copper 2+ or correspondingly, copper 0 is a better reducing agent than silver. We would predict from that, then, that the initial reaction we saw, of copper reacting with silver, would be a spontaneous reaction.
Let's answer our question, then. We were worried about whether copper 2+ would be able to get electrons out of zinc. So here's copper 2+ and zinc is way up here. Notice on our scale, zinc is a much better reducing agent than copper metal is a reducing agent. That means that the reaction of copper 2+ and zinc should, indeed be spontaneous. This half-reaction is much lower on our series than this one is. Alternatively I could have said to you copper 2+ is a much better oxidant than zinc 2+. So copper 2+ would, indeed, react with zinc to give us, in this case, copper metal.
So who cares? What is this good for? It actually has a lot of good uses. Suppose that you had an iron pipe and your plan is to have this pipe outside for a number of years. You're worried about rust. You're worried about this reaction. You're worried about iron reacting with oxygen gas in the present amount of H+, which we could just get from water to give us iron 2+ and water. This goes on to actually give us rust. Let's just consider the first step of the rusting process. That reaction, if we go to our activity series, is certainly going to be spontaneous. There's iron. Way down here is oxygen. Once again, oxygen is a better oxidant than iron 2+ is so it is going to be a spontaneous reaction to have the oxygen to get what it wants. Essentially, it is going to take electrons from iron metal.
What if we could coat the pipe with something else, something that would react with iron 2+ to give us back our iron metal and not consume the pipe? We'd need to choose something to coat it with that was higher in the series, something which could give its electron to iron 2+. In other words, iron 2+ is a better oxidant than any of these guys. Likewise, these guys are better reducing agents. What if we coated it with zinc? Zinc is cheap. Zinc is more reducing as a metal than iron. That, indeed, is what a galvanized pipe is. It is an iron material. Galvanized nails, perhaps you've heard of, they're coated with zinc. As a result of being coated with zinc, if any of the iron is oxidized, before it can survive at all as iron 2+, that immediately would be reduced by the zinc, returning the iron to its metal form. In fact, that never really happens because the zinc is a better reducing agent than the iron is a reducing agent and the zinc, if you will, sacrifices itself to protect the iron. When this happens, it is coated with an oxide coating, zinc oxide. That turns out to be a material that blocks oxygen attacking. In a way, it's like a Band-Aid. It protects itself from further attack by air. If someone then scratches it, immediately this reaction happens a little bit creating more zinc oxide that blocks the surface again, or zinc hydroxide depending on if it reacts with water. That will block the surface again and protect it from further corrosion. So you see, by knowing the activity of zinc metal compared with iron metal we're able to design a product that, in this case, protects pipes from rusting.
Oxidation-Reduction Reactions
Looking In-Depth at Redox Reactions
The Activity Series of the Elements Page [1 of 2]