Single replacement reactions | Chemistry | Khan Academy

If you put a copper wire in this silver nitrate solution, then you'll get this beautiful reaction. But instead of copper, if you were to put a wire of gold in the same silver nitrate solution, the same solution as before, this time nothing would happen—no reaction. Why? That's what we'll find out in this video. But by using that knowledge, we'll also be able to answer some curious questions, like why are acids stored in glass containers and not metallic ones. Or think about this: earlier when we used to build hydrogen balloons, where did we get that hydrogen from? If you're ready, let's find out the answers to these.

So let's start with what we saw earlier. We put a copper wire in the silver nitrate solution. So the reaction is solid copper reacting with silver nitrate—aqueous silver nitrate solution. What happens? We'll get copper(II) nitrate aqueous solution and silver in the elemental form. So let's think about what's going on over here. See, copper over here is swapping out silver. That's why we're getting the copper(II) nitrate, and silver is coming out in its elemental form. This is why the solution starts turning blue, because of the copper(II) nitrate solution, which is blue in color. And what's growing on top of this copper wire is the silver. If you were to zoom in over here, it's the silver crystals that keep growing on the copper wire, which makes this whole reaction beautiful.

So, because copper replaces silver in this reaction, we call such reactions a single replacement reaction. Of course, we can also call it the single displacement reaction because copper is displacing the silver over here. But anyways, if you're wondering why it is called single, well, because it turns out there's something called double replacement or double displacement reactions as well. We'll talk about that in future videos. But anyways, in general, we can denote the single replacement reaction as A reacts with BC, then A will swap out B. A will replace B, giving you AC plus B.

Okay, that now brings up the question: why does copper replace silver over here? Well, remember that all chemical reactions are mostly about electron transfers. So if you understand how the electrons are being transferred, we'll probably be able to get a clue as to what's going on over here.

Okay, so to think about electron transfers, we can just look at what's happening to the charges. Okay, so over here, the copper is neutral, so it has a zero charge. When it comes to silver nitrate, remember metals usually are the cations, so they have a positive charge. The nitrate would be having a negative charge—the whole nitrate, not just the oxygen; the whole nitrate has a negative charge. Okay? But on the other side, copper becomes the cation. Of course, a small detail is copper cation over here has a two plus charge—that's why this is called copper(II) nitrate.

Okay, but let's not worry about why that is the case. What matters is copper is the cation, nitrate is the anion, and of course, silver ends up becoming neutral over here; it's in the elemental form so we'll have zero charge. So now think about what's happening to copper. Copper went from being neutral to positive. How could it be? How could that be? Well, that's because it must have lost electrons.

So from this, we could say, "Ooh, copper has lost electrons," and what has happened to the silver? It went from being positive to neutral. How can that be? "Ooh, silver must have gained these electrons from copper." So silver must have gained electrons, and nothing has happened to the nitrate. Nitrate charge has not changed at all; it remains the same. So we can now get a clue as to what's going on: copper is losing the electrons, becoming cations, and these electrons are gained by the silver atoms, and that's how they are becoming neutral, coming out as the elemental form.

Then the copper cations will then combine with the nitrate ions to form copper(II) nitrate solution. That's what's going on over here. Now guess what? We've also learned about a formal way to talk about this stuff. Losing electrons is what we call oxidation, and gaining electrons is what we call reduction. So we can also say that copper is getting oxidized and the silver over here is getting reduced.

And by the way, a quick thing if you're getting confused between which is oxidation and which is called reduction, then the way I like to think about it is think in terms of what's happening to the charge. The positive charge is becoming zero, so the charge is reducing—that's why this is reduction, and the other one is oxidation. We've talked about that in our previous videos, but what's important is that this means single replacement reactions are all redox reactions.

Anyways, before we move on, one final thing over here is that this equation is not balanced. So because there are two nitrate ions over here, we can balance it by putting a two here and then by putting a two over here. Okay, so that's the story of what's going on over here. But now that brings up the question: what happens when you put gold in silver nitrate solution?

Well, turns out nothing; you get no reaction over here at all. Gold is unable to replace silver, and that's why nothing happens. But why not? Why is copper able to replace silver over here, but gold is unable to? Well, here's what's going on: see, all metals have tendencies to lose electrons, but that tendency is not the same. Copper, for example, is more active, meaning it has a higher tendency to lose electrons compared to silver— that's why copper ends up getting oxidized, and silver is forced to gain electrons and get reduced.

But when it comes to gold, it turns out it's less active compared to silver, and therefore its tendency to lose electrons is less compared to that of silver. Silver has a higher tendency to lose electrons, and that's why gold cannot oxidize; it's less active. So look, if you want to replace someone, you have to be more active. You need to have a higher tendency to lose electrons; otherwise, the replacement reaction does not happen.

So guess what this means? If you want to be able to predict whether a single replacement reaction happens or not, we need to know whether this metal is more active than this metal or not. But how do we know that in general? Well, that's where the activity series comes into the picture. Here, the more reactive, or the more active metals are on the top, and the lesser ones are on the bottom. So the ones on the top can replace the ones on the bottom, but not the other way around.

So, for example, sodium can replace zinc from some solution, but iron, for example, cannot replace, say, potassium from some solution because iron is less reactive; it's below. So it's less reactive compared to potassium. And now, if we come back to our original example, you can see copper is above silver, which means copper is more active compared to silver, so it can replace it.

But look at gold; it's less active—it's below the silver, so it's less active compared to silver, and that's why it cannot replace silver in this particular reaction. So look, the activity series is pretty cool; it helps us predict whether the replacement reactions are going to happen or not.

So let's take a few more examples. For the first one, let's put zinc in copper sulfate solution. It would be a great idea to pause the video and make a prediction of what's going to happen. Use the activity series and think about whether the replacement reaction happens, and if it does, what does that reaction look like? Who's getting oxidized? Who's getting reduced? So pause and try.

Okay, here it goes. So we're adding solid zinc—zinc pellets—to copper sulfate aqueous solution of copper sulfate. So the question is: can zinc replace copper? For that, I go back to my activity series and I see zinc is on the top; copper is on the bottom, which means zinc is more active compared to copper. Yes, zinc can replace copper, and the way it does that is zinc is going to lose electrons and become cations.

So zinc will usually lose two electrons to give you zinc(II) plus, and if you were to write the charge over here, here it was neutral; it had zero charge here. Copper was the cation; metals are usually the cations, sulfate was the anion. Okay? So zinc becomes two plus; it now combines with sulfate to form zinc sulfate—so you get two minus over here. This will be the aqueous solution, and the copper over here will gain the electrons and become solid copper—so it gets reduced.

So zinc gets oxidized, copper gets reduced, which means the reaction should proceed. And I tried it at home; I put some zinc pets into the copper sulfate solution and waited for a long time. Here's a time lapse: if you zoom in, you can actually see the black stuff that is being formed is the copper being deposited on the zinc pellets. And because now copper is moving out of the solution, eventually the solution turns colorless—most of the copper sulfate is gone, and what's left now in the solution is zinc sulfate, and the black stuff is the copper. Beautiful, right?

Alright, here's a second example: suppose we were to take some elemental silver and put it in zinc sulfate solution. What will happen? Why don't you pause and try?

Alright, so the question is: can silver replace zinc? And for that, I go back to my activity series. Zinc is over here; silver is here. Silver is below zinc, so that means silver is less active than zinc. That means it can't— it cannot replace zinc, and therefore nothing happens over here.

Now, before we move on, do you see something weird in the activity series? It's supposed to be all metals; this is all about metals, but guess what? We also see hydrogen over here. What's hydrogen doing over here? This is not a metal, right? Well, guess what? Hydrogen can also lose electrons and form cations and can now combine with anions like chloride ions to form HCl or it can combine with sulfate ions to form H₂SO₄ and other compounds we call these acids.

And so, since we can form such solutions, it is also part of the activity series. And so let's do an example of that. Suppose I take magnesium and put it in hydrochloric acid. What do you think will happen? Again, pause and try.

Here goes. We're adding magnesium solid to HCl, which is the aqueous solution. The question I ask myself is: can magnesium replace hydrogen over here? Here hydrogen is the cation, so can it replace hydrogen? For that, I go back to my activity series and I see magnesium is on top of hydrogen, so yes, it can replace hydrogen.

So magnesium is going to get oxidized, meaning magnesium will become the cation; it becomes two plus. And again, over here it was zero; here hydrogen had a positive charge, and chlorine had a negative charge. So it'll combine with chloride, but since it has two plus charge, it has to combine with two chloride ions, so you get MgCl₂, and hydrogen will accept the electrons, and it will bubble out as hydrogen gas. Remember, hydrogen is diatomic, so it comes out as gas.

Okay, and we can now try to balance this reaction. There's two chloride ions over here, so I put a two, and so you have two H, and everything gets balanced. And again, we can see that over here, magnesium is replacing hydrogen—hydrogen is the one that's bubbling out over here.

Chemical reactions become much more fascinating when we understand what's going on, isn't it? And now we can answer some of our original questions from this reaction. We can see how to get hydrogen gas, and that's what people did. They did not use magnesium, but they used iron. Iron is also more reactive than hydrogen, as you can see, and therefore if you put iron in hydrochloric acid, hydrogen will bubble out—a very similar reaction as over here.

And this also explains why we cannot use metallic containers to store acids, because a lot of metals are more reactive than hydrogen—they'll just react with the container. This is why we use glass containers to store acids.

Anyways, before we end this video, here's the last interesting piece. So far, in all the examples, it was the cations replacing other cations, right? And nothing was happening to the anions. Well, guess what? There are some rare types of single replacement reactions in which anions can replace other anions. Now, this usually happens in the halogens. For example, if you were to take chlorine gas and pass it through potassium bromide—an aqueous solution of potassium bromide—then the chlorine can replace the bromide ions over here to give you potassium chloride and bromine, which is liquid at room temperature.

Now again, if you were to look at the charges, this is neutral potassium; that is the cation. Bromide ion over here, again potassium is the cation. Now we have a chloride ion, and this becomes neutral. So look, this time nothing happened to the cation, but it was the an that got replaced. This mostly happens with the halogens.

And if you're wondering now what is the activity series for this, well then we can just look at the periodic table; we don't need another activity series for halogens. The ones on the top can replace the ones on the bottom—that's why chlorine can replace bromine, but bromine, for example, cannot replace fluorine.

So for our final example, let's put iodine in potassium bromide aqueous solution. Will we get any reaction? What do you think? Well, iodine is over here; it is below bromine, and therefore it will not be able to kick out bromine over here. And also, in such anion replacements, you can see that it's the chlorine that is getting reduced. You see that? It went from zero to negative; it gained electrons and becomes reduced, and it's the bromine that gets oxidized and comes out as a liquid.