Chemguide: Core Chemistry 14 - 16

Making insoluble salts

This page looks at the preparation of insoluble salts using precipitation reactions. You may need to refer to the table of solubility on the page about salts and solubility.

What are precipitation reactions?

This is NOT a precipitation reaction!

Imagine that you have two solutions which contain soluble salts:

  • KI solution containing K+(aq) and I-(aq) ions,

  • Mg(NO3)2 solution containing Mg2+(aq) and NO3-(aq) ions.

There will, of course, be two nitrate ions for every one magnesium ion in the second solution to balance the charges.

There are also lots of water molecules in both solutions.

Now suppose you mix the solutions. What will you have then?

The mixed solutions contain

  • K+(aq) ions

  • I-(aq) ions

  • Mg2+(aq) ions

  • NO3-(aq) ions

What happens when the various ions bump into each other?

If two positive ions meet, they will repel each other - so nothing happens. The same is true of the negative ions.

What happens if a potassium ion meets a nitrate ion?

Potassium nitrate is a soluble salt. The attractions between potassium ions and nitrate ions aren't strong enough to overcome the attractions between the ions and the water molecules. Nothing happens when they meet.

What happens if a magnesium ion meets an iodide ion?

Again nothing happens. Magnesium iodide is also a soluble salt. The attractions between magnesium ions and iodide ions aren't strong enough for them to overcome the attractions between the ions and the water molecules.

What you end up with is simply a solution containing all four ions.

Note:  You won't find iodides listed in the table of salt solubilities on the page mentioned above. The only time you are likely to come across them is in this topic.

This IS a precipitation reaction!

Now let's make a slight change to what we have just described. We will replace magnesium nitrate solution by lead(II) nitrate solution.

If you mix this solution with potassium iodide solution, the ions in the mixture are now

  • K+(aq) ions

  • I-(aq) ions

  • Pb2+(aq) ions

  • NO3-(aq) ions

The only difference is what happens when a lead(II) ion meets iodide ions.

Lead(II) iodide is insoluble in water. The attractions between lead(II) ions and iodide ions are strong enough to overcome attractions between the various ions and water molecules, and they stick together.

Solid lead(II) iodide is formed as a bright yellow precipitate.

A precipitate is simply a solid formed by the reaction between two liquids or between a gas and a liquid.

Pb2+(aq) + 2I-(aq)     PbI2(s)

Using precipitation reactions to make insoluble salts

The video below shows how you would use the reaction we have been talking about to make a pure sample of lead(II) iodide.

Summary of the method

  • Mix the two solutions which give the precipitate.

  • Filter the mixture to leave the precipitate on the filter paper.

  • Rinse the precipitate with distilled water to remove all the soluble compounds which will be contaminating the precipitate at this stage. Several small rinses are much more effective than the one big one that the video shows.

  • Leave the filter paper and precipitate to dry, and then scrape the dry solid off the paper.

Note:  The video wasn't totally complete in listing what soluble compounds might have to be removed from the precipitate on the filter paper by rinsing it.

It correctly mentioned that there is almost certainly an excess of one of the solutions you mixed - in this case, either excess potassium iodide or lead(II) nitrate.

But you are certain to have potassium ions and nitrate ions (as potassium nitrate) left after the reaction. These were in the original solutions, and nothing has happened to them - they are still there after the formation of the lead(II) iodide, and have to be removed.

After the filtration, the precipitate isn't damp with water, but with potassium nitrate solution (plus probably an excess of one of the solutions you mixed as well).

Two more examples

The precipitates of insoluble salts that you will meet later in the course are silver chloride and barium sulfate. These are produced during the chemical tests for chlorides and sulfates in solution.

Silver chloride

A white precipitate of silver chloride is formed if you add silver nitrate solution to any solution of a chloride. The silver ions clump together with chloride ions to make insoluble silver chloride.

Ag+(aq) + Cl-(aq)     AgCl(s)

There is a simple bit of animation showing this happening. The video is very short and contains a lot of information, so you will probably need to watch more than once - possibly with lots of pauses. I will pick out a few things to look at after you have watched it once.

You may need to pause the video and look at the key on the right-hand side so that you know exactly which ion is which.

You may be unsure about the conductivity charts at the top of the screen. Conductivity is a measure of the solution's ability to conduct electricity.

Electricity is conducted through solutions by the movement of ions. The conductivity falls when the reaction is complete because there are fewer ions in the solution. You could safely ignore this!

Barium sulfate

A white precipitate of barium sulfate is formed if you add barium chloride solution to any solution of a sulfate. The barium ions clump together with sulfate ions to make insoluble barium sulfate.

Ba2+(aq) + SO42-(aq)     BaSO4(s)

Using ionic equations for these reactions

You will have noticed that I have only used ionic equations on this page, and not made any attempt at writing the full equations. There are two reasons for this.

The ionic equations go to the heart of the reaction, and show that the precipitate is formed by two ions clumping together.

But the main reason is that it is far less bother writing the ionic equations. You are just making unnecessary work for yourself if you write the full equation, and increase the chances of getting it wrong.

Suppose you had to write an equation for the formation of a precipitate of magnesium carbonate in the reaction between magnesium sulfate solution and sodium carbonate solution (or any other soluble magnesium salt and soluble carbonate).

Here's what you do:

Work out the formula for the solid precipitate:


On the left-hand side of your equation write down the formulae for the ions that combined to make it, add an arrow, and then the formula you have just worked out:

Mg2+(aq) + CO32-(aq)     MgCO3(s)

And that's it!

Note:  You must, of course, know or be able to work out, the charges on ions and the formulae of complex ions like carbonate or sulfate or nitrate. If you can't do that then you must spend enough time on the page about the formulae of ionic compounds until you can.

You can't succeed in chemistry without this skill.

What this equation says is that solutions of any soluble magnesium salt and any soluble carbonate will give you a precipitate of magnesium carbonate if you mix them.

It doesn't matter whether it is magnesium sulfate or magnesium nitrate or magnesium chloride solution. And similarly, it doesn't matter whether you add sodium carbonate, potassium carbonate or ammonium carbonate solution.

So . . . if you had to make an insoluble salt like barium sulfate, what solutions would you mix?

You need a soluble barium salt. If you refer back to the page about salts and solubility, you can choose either barium nitrate or barium chloride solution.

You need a soluble sulfate. You have a much wider choice here - you could choose anything other than the ones marked as insoluble or slightly soluble.

Whatever you choose, the ionic equation is always the same.

If you want to make life easy for yourself, just remember that all nitrates are soluble, as are all sodium, potassium and ammonium salts. Make your choices from them.

And finally . . .

A piece of video which makes no demands on you whatsoever - nothing to learn, no new skills. A drop of one solution is added to another solution so that a precipitate is formed.

It is probably best watched at full screen.

Where would you like to go now?

To the acid, bases and salts menu . . .

To the Chemistry 14-16 menu . . .

To Chemguide Main Menu . . .

© Jim Clark 2020