Lesson Video: Tests for Cations Chemistry

Video Transcript

In this video, we’re going to learn how to identify a few different positively charged aqueous ions based on the colors of their solutions and how they react with sodium hydroxide and ammonium hydroxide solutions. We’ll also look at a special test just for the ammonium cation. We’ll only be looking at some of the most common ions you might find in solution in the lab. There are, of course, many more cations than this, but this selection shows an interesting range of responses. If you use your detective skills, you may be able to figure out which one of these there is in a particular solution.

Now, let’s meet the ions. First, we have the aluminum ion with a charge of three plus. Next, we have the calcium ion with a charge of two plus. Then, there’s the chromium three ion with a charge of three plus and two different types of iron ion: iron two and iron three, with charges of two plus and three plus, respectively. Next, the copper two ion with a charge of two plus and the last of our metal cations zinc two plus. We’ll also look at the ammonium ion, NH4+. Of course, to have any of these in solution, we must be dealing with a dissolved ionic compound. This means if we have some of these positively charged cations, we’re going to need some anions, which are negatively charged, for instance, chloride or nitrate anions.

For this video, I’m going to assume we’re dealing with the chloride salts of all these cations. The chlorides are all soluble in water or at least slightly soluble. We’ll start off by looking at each ion individually and then we’ll compare them. But before that, let’s have a look at the tests. The tests we’ll be using are the color of the solution of the ions, although we’ll be focusing more on the results of the chemical tests. Our first chemical test will be treatment with sodium hydroxide solution, first in dilute form and then to excess. The metal cations we’ll be looking at will produce precipitates when treated with dilute sodium hydroxide solution. Some of these precipitates will then disappear if treated with an excess. And our second chemical test involves treatment with first dilute ammonium hydroxide and then an excess.

Some of the ions respond differently to ammonium hydroxide solution than sodium hydroxide solution, allowing us to distinguish one from another. There will also be a special test for the ammonium ion, which we’ll come to later. In the lab, you may come across sodium hydroxide in various forms, for instance, as a white powder or pellets or as a clear, colorless solution. Ammonium hydroxide looks like this. It’s also a clear, colorless solution. If you add ammonia, NH3, to water, you’ll produce ammonium hydroxide. Sometimes you’ll see ammonia solution, ammonium hydroxide, or ammonium hydroxide solution; these all refer to the same thing.

Meanwhile, the term dilute is a relative term. However, with sodium hydroxide solution or ammonium hydroxide solution, you might be looking at a concentration of about 0.1 molar, while a concentrated solution may have a concentration of one or more molar. If we used a concentrated solution rather than a dilute solution, we might jump directly to the results of the excess test. Nonetheless, be wary that these concentrations are just guidelines. The results of these tests will depend on the amounts added and the concentration of cation in solution.

Now, it’s time to move on to our first candidate. A solution of aluminum chloride will be clear and colorless. If we have a solution that contains Al3+ ions and we add dilute sodium hydroxide solution, we’ll produce a white precipitate of AlOH3, aluminum hydroxide. But something slightly different happens when we add an excess of a sodium hydroxide. If there’s enough sodium hydroxide in the solution, the aluminum hydroxide precipitate will react further and disappear. In this video, we won’t be looking at the nature of the products involved when precipitates redissolve. What we’re focusing on here is the pattern of precipitates and their color.

Now let’s start with a fresh test tube of our solution and do the next test. If we add a dilute solution of ammonium hydroxide to a solution containing Al3+ ions, we’ll form the same white precipitate of aluminum hydroxide we formed with sodium hydroxide. Just like a sodium hydroxide solution, ammonium hydroxide acts as a source of the hydroxide ion. However, if you add an excess of ammonium hydroxide or concentrated ammonium hydroxide, the precipitate will not redissolve.

Again, we’re not going to go into the reasons for this; we just have to remember it. In both chemical tests, hydroxide ions have been added to the aluminum ions, forming a precipitate of aluminum hydroxide. The hydroxide anion only has a charge of one minus, so we need three of them to counterbalance the aluminum’s three plus charge. So here is our balanced ionic equation, and we need to add in state symbols, taking care to include a solid symbol for our precipitate.

Now we’ve done one cation, we can go through the others more quickly. Here is everything we need to fill in for the next ion, which is calcium two plus. A fresh solution of calcium chloride is clear and colorless. Treating a solution of calcium two plus ions with dilute sodium hydroxide will produce a white precipitate of calcium hydroxide. Adding an excess will not affect the precipitate; it will not redissolve. On the other hand, adding ammonium hydroxide solution to a solution of calcium two plus ions will not produce a precipitate in the first place. Adding an excess will have no extra effect. Unlike aluminum ions, calcium ions only have a charge of two plus, so we only need two hydroxide ions to balance the charges.

Next up is chromium three. A solution of chromium three chloride will be a dark green, although it’s only slightly soluble. Adding dilute sodium hydroxide solution to a solution of chromium three plus ions will produce a gray-green precipitate of chromium three hydroxide. Adding an excess of sodium hydroxide solution will cause the gray-green precipitate to disappear. If we use dilute ammonium hydroxide instead, we’ll produce the same gray-green precipitate of chromium three hydroxide. But adding an excess will have no effect and the precipitate will remain. Since chromium three ions have a charge of three plus, we end up with a formula of CrOH3 for the hydroxide.

Next, we have iron two, Fe2+. Solutions of iron two chloride are green. Adding dilute sodium hydroxide solution to a solution of Fe2+ will produce a pale green precipitate of iron two hydroxide but adding an excess will have no further effect. Adding dilute ammonium hydroxide solution produces the same pale green precipitate, with an excess also having no further effect. As with calcium, we’ll only need two hydroxide ions per every two plus ion to balance the charges. The next candidate Fe3+ exhibits very different colors to Fe2+, so pay careful attention to the differences.

Instead of being green, a solution of iron three chloride is yellow or brown. The result of a test with dilute sodium hydroxide solution is a red-brown precipitate of iron three hydroxide, with an excess producing no further effect. Adding dilute ammonium hydroxide produces the same red-brown precipitate of iron three hydroxide. And an excess has no further effect. So we see the same pattern of precipitations with iron two and iron three but very different colors. As with aluminum and chromium three, we need three hydroxide ions to balance the charge of Fe3+, giving us FeOH3.

Next, we’ll look at copper two. Solutions of copper two chloride are blue. Adding dilute sodium hydroxide solution to a solution of copper two plus will produce a pale blue precipitate of copper two hydroxide. And an excess will have no further effect. Adding dilute ammonium hydroxide instead will produce the same pale blue precipitate of copper two hydroxide. However, in an excess of ammonium hydroxide, the precipitate reacts, dissolves, and the solution turns a darker blue. Our ionic equation for the production of the initial precipitate is simply Cu2+ plus 2OH− react to form CuOH2. Again, we’re not going to focus on the nature of the product when the precipitate dissolves.

Next up, we have the last of our metal cations. Solutions of zinc chloride are colorless. Adding a dilute solution of sodium hydroxide to a solution of Zn2+ ions will produce a white precipitate of zinc hydroxide. With excess sodium hydroxide solution, the precipitate will react and dissolve. We get exactly the same result with ammonium hydroxide solution and excess ammonium hydroxide solution, with the white precipitate of zinc hydroxide dissolving in excess ammonium hydroxide. And in our balanced ionic equation, the zinc two plus ions and OH− ions are in a ratio of one to two.

Now let’s move on to the only nonmetal cation in this video. Ammonium chloride forms colorless solutions. We can begin to detect ammonium ions by adding sodium hydroxide either in solid or solution form. What happens is essentially an acid base reaction, where the acidic ammonium ion reacts with the basic hydroxide ion producing ammonia gas and liquid water. You may see bubbles produced, but ammonia doesn’t have a color, so how do we know its ammonia? The most obvious way is via the smell because the smell of ammonia is very distinctive. However, if you have the option, you should not inhale ammonia. Instead, the best option is to check whether the gas is basic using wet, red litmus paper. If the gas is ammonia, the ammonia will dissolve into the water, turning the red litmus paper blue. In some cases, you may need to provide some heat to produce enough ammonia for the test.

Now that’s all the ions we’re going to look at in this video. Let’s look at them all at once. Here’re the cations and here’re the tests; I’ve isolated ammonium for now. If we get a precipitate for a particular test, we’ll put a tick; otherwise, we’ll put a cross. For all the metal ions in this video, we get a precipitate when we treat their solutions with dilute sodium hydroxide solution. For aluminum, calcium, and zinc, the precipitate is white, while the precipitates for chromium three, iron two, iron three, and copper two are all different colors. For the aluminum, chromium three, and zinc ions, adding an excess of sodium hydroxide will cause the precipitate to dissolve. For the rest, adding more sodium hydroxide solution does not affect the precipitates.

If instead of using sodium hydroxide solution we use ammonium hydroxide solution, we get exactly the same results for all the ions with dilute ammonium hydroxide apart from with calcium, where we don’t get a precipitate at all. And there’s still no precipitate if we add an excess. For copper two and zinc ions, when we add an excess of ammonium hydroxide, the precipitate will react and redissolve. In the case of the copper two ions, the solution will turn a much darker blue than before. For the remaining metal ions, adding an excess ammonium hydroxide will have no further effect on the precipitate. And finally, for the ammonium ion, all we need to do is add sodium hydroxide and a little heat. We can then treat any gas given off with wet, red litmus paper, which should turn blue. After all that, it’s time for some practice.

Which of the following metal cations does not produce a precipitate when a few drops of aqueous ammonia are added to a salt or solution of that metal cation? (A) Zn2+, (B) Al3+, (C) Ca2+, (D) Cr3+, or (E) Cu2+.

The question describes these five options as metal cations. We’re dealing with cations of zinc, aluminum, calcium, chromium, and copper. And we’re being asked which of these metal cations as a salt or in solution would not respond to aqueous ammonia and produce a precipitate. The formula for ammonia is NH3. Ammonia can react with water to produce ammonium hydroxide, a good source of the hydroxide ion. The question doesn’t tell us directly if the ammonium hydroxide is dilute, to an excess, or concentrated. But since it’s only a few drops, we can safely assume it’s dilute.

The question is, what is going to happen when we treat each of these ions with dilute aqueous ammonia? Zn2+ ions will react to produce a white precipitate of zinc hydroxide, and we get a similar reaction with Al3+ ions, producing aluminum hydroxide. However, calcium two plus ions do not react with aqueous ammonium hydroxide no matter the concentration. There isn’t an easy way to understand why; you’ll just need to remember it. Meanwhile, chromium three ions do react, forming a gray-green precipitate of chromium three hydroxide, while the precipitate we get from copper two plus is pale blue copper two hydroxide.

Now, there is one possible point of confusion with this question if we’re not quite sure if we’re dealing with dilute or excess ammonium hydroxide. In excess ammonium hydroxide, the precipitates formed from Zn2+ and Cu2+ ions dissolve. But since we are only dealing with a few drops, we can be fairly confident that we aren’t dealing with dilute conditions. Therefore, out of the five options, the only metal cation that does not produce precipitate when treated with a few drops of aqueous ammonia is calcium two plus.

Now, let’s finish up with the key points. Different cations can be detected based on the precipitates they do or do not form. And these are the ions and tests we’ve looked at. These are the colors of the precipitates we get when we treat each of these ions with dilute sodium hydroxide solution. If we treat these with excess sodium hydroxide solution, the precipitates for aluminum three plus, chromium three plus, and zinc two plus will redissolve. If instead of using dilute sodium hydroxide we use dilute ammonium hydroxide instead, otherwise known as aqueous ammonia, we get exactly the same precipitates apart from calcium, where we don’t get a precipitate at all.

If we continue to treat with ammonium hydroxide to the point of excess, the precipitates for copper two plus and zinc two plus redissolve. The rest stay as they were. And the last cation, ammonium, can be detected by treatment with sodium hydroxide and potentially a little heat. We can smell the ammonia gas given off or detected using wet, red litmus paper, which should turn blue.