Chemguide: Core Chemistry 14 - 16

Making soluble salts

This page gives some experimental details for the preparation of soluble salts by reacting acids with metals, metal oxides or hydroxides It assumes that you already know about the reactions of metals and bases with acids.

Insoluble salts are made by an entirely different method, and you will find a link to them at the bottom of the page.

Two different methods

There are two completely different experimental ways of doing this.

In most cases, you can add an excess of a solid metal or metal oxide to an acid. Then filter off the excess and crystallise the solution you have made.

By doing that you have made sure that you have used up all the acid, and your solution will contain the pure salt you are trying to make.

But there is a problem with making sodium, potassium or ammonium salts in this way. You would normally choose to make these from sodium or potassium hydroxide - both of which are soluble in water, and you use them as a solution. Ammonium hydroxide only exists in solution.

That means you can't add an excess and then filter that excess off. Neither can you tell just by looking at a reaction between them and an acid when you have got exactly the right proportions.

So making sodium, potassium and ammonium salts needs a completely different experimental method.

Making soluble salts of most metals

This is about making soluble salts of everything except sodium, potassium or ammonium.

The reactions normally used are

• metal + acid    salt + hydrogen

• metal + metal oxide    salt + water

The reaction between an acid and a metal oxide is more useful because it can be used for any metal without having to bother about where it is in the reactivity series. You could, however, easily use the reaction between magnesium or zinc and possibly iron and an acid.

You could use a solid metal hydroxide, but these are less commonly used in the lab than oxides.

You could also use a carbonate, and we will talk about those on another page.

Making copper(II) sulfate crystals from copper(II) oxide and dilute sulfuric acid.

Start by watching this YouTube video.

There is one slight oddity about this video in the way he demonstrates folding a filter paper. He over-folds it in the demo, but when he fits it into the filter funnel, it is folded properly!

In case you missed it, the thermometer reads about 60°C.

The reaction creating the copper(II) sulfate solution is

CuO(s) + H₂SO₄(aq)     CuSO₄(aq) + H₂O(l)

During the formation of crystals from the solution, water from the solution becomes bound to the copper(II) sulfate as the crystals form. This is called water of crystallisation. We sometimes describe the crystals as hydrated copper(II) sulfate.

CuSO₄(aq) + 5H₂O(l)     CuSO₄.5H₂O(s)

The dot (sometimes a comma) between the copper(II) sulfate and the water shows that they are attached to each other.

Making magnesium chloride crystals from magnesium and dilute hydrochloric acid.

You could use exactly the same technique to make magnesium chloride crystals from magnesium metal and dilute hydrochloric acid.

This time you don't need to heat the reaction - in fact, the reaction gives out a lot of heat itself. You just add enough magnesium to the acid so that when the reaction stops there is some metal left. That shows that all the hydrochloric acid is used up.

You can filter that off, concentrate the solution, and then leave it for crystals to form. You will get colourless crystals.

Mg(s) + 2HCl(aq)     MgCl₂(aq) + H₂(g)

MgCl₂(aq) + 6H₂O(l)     MgCl₂.6H₂O(s)

Making magnesium sulfate crystals from magnesium oxide and dilute sulfuric acid.

Here is another bit of video from the same source as the first one showing the reaction between magnesium oxide and sulfuric acid to produce magnesium sulfate crystals.

MgO(s) + H₂SO₄(aq)     MgSO₄(aq) + H₂O(l)

MgSO₄(aq) + 7H₂O(l)     MgSO₄.7H₂O(s)

Making sodium, potassium and ammonium salts

You have to use a different method with these because all sodium, potassium and ammonium compounds are soluble. That means that you can't add an excess of, say, a hydroxide, and then filter off the excess.

Any excess will have dissolved in the solution you are making. You can't tell when the proportions are just right.

So if you wanted to made sodium chloride, you would use a solution of sodium hydroxide and find out how much hydrochloric acid you have to add in order to make a neutral solution.

Originally, the sodium hydroxide solution is alkaline. When you have added too much acid, it will become acidic. You need to aim for a neutral solution. To do this you have to do a titration.

This uses two accurate pieces of measuring equipment - a pipette and a burette

A pipette measures a fixed amount of liquid; a burette measures a variable amount.

I'm going to show you a sequence of 4 pieces of video showing how these bits of apparatus are used.

The first shows how you rinse out a pipette and burette with water.

The next one shows how you fill the burette with titrant. Titrant is a posh word for whatever solution you are putting into the burette. In my experience few people in school chemistry ever use the word.

I have two problems with this piece of video.

First she uses a glass funnel to add the solution to the burette. This is fine as long as the funnel neck is a lot smaller than the top of the burette. If it isn't, the funnel can get jammed and you risk breaking a quite expensive burette getting it out again. In school you would probably use a small plastic funnel.

More importantly, a burette must always be filled below eye level to stop any possibility of any of the solution splashing into your eyes. And the burette should be clamped, not waved around!

If you are short, and can't get the top of the burette below your eye level using a stand and clamp on the bench, place the stand on a lab stool to fill the burette.

The next video shows how to fill and use a pipette. In the olden days you sucked liquid into the pipette by putting it in your mouth. That is seriously frowned upon these days!

I have a safety problem with the way she holds the pipette as she puts it into the pipette filler. By holding the pipette so far from the filler, she risks putting enough pressure on the pipette that it breaks - and then the broken end can stab your other hand. I have seen that happen, and you will probably end up in hospital if that happens to you!

As a minor point, she doesn't specifically say that she discards the solution that she is rinsing the pipette with before filling it properly.

Finally, you find out how much of the solution in the burette is needed to produce a neutral solution by using an indicator.

The indicator she is using is called screened methyl orange, which is green in alkaline solution and turns grey when you have reached the neutral point.

She talks about adding the acid slowly when you get close to the end point of the titration, but doesn't say how you would know that.

When you open the tap on the burette, where the acid first hits the solution, the indicator changes colour before the acid has time to mix and react. Swirling the flask mixes the acid with the alkali and they react - the indicator would go back to tha alkaline colour.

Towards the end point, it takes longer and longer for the temporary colour change to disappear. When that starts to happen, you add the acid more and more slowly.

Making sodium chloride

You would take a known volume of sodium hydroxide solution in the pipette and transfer it to a conical flask. Put dilute hydrochloric acid in the burette.

The indicator used in the next video is phenolphthalein. This is pink in alkaline solution, and turns colourless when the sodium hydroxide is neutralised to make sodium chloride solution.

NaOH(aq) + HCl(aq)     NaCl(aq) + H₂O(l)

The video is accurate, but done very quickly. Watch it once and then look at it again if you need to. All the safety points I mentioned in the previous videos are done correctly here. The pipette filler used is a different type, but it has the same function - to suck liquid into the pipette.

Unfortunately this video can't be embedded on other sites, so you will have to watch it on YouTube. Use the back button on your browser to come back here afterwards.

The video mentions methyl orange as a possible alternative to phenolphthalein. The colour change for this is yellow in alkaline solution to red in acidic solution. The end point of the titration is the first trace of orange in the solution.

This final bit of video shows a titration using methyl orange. The video shows the yellow alkaline colour of the indicator and the point that you first see some orange in the solution. It then goes on to add too much acid to show the red acidic colour.

Bizarrely, the video spins itself out to fill the length of the piece of music they use. Once you have seen the colour changes, nothing else of any interest happens.

Summary

• Do a titration to find out how much hydrochloric acid you need to neutralise the sodium hydroxide solution.

• Mix exactly the same volumes of acid and alkali as you did before, but with no indicator.

• Crystallise the solution.

Other similar reactions

You could use any combination of sodium, potassium or ammonium hydroxide solution with hydrochloric, sulfuric or nitric acids. All of these are colourless solutions, and they all form colourless salts.

For example:

KOH(aq) + HNO₃(aq)     KNO₃(aq) + H₂O(l)

2NaOH(aq) + H₂SO₄(aq)     Na₂SO₄(aq) + 2H₂O(l)

NH₄OH(aq) + HCl(aq)     NH₄Cl(aq) + H₂O(l)

For complicated reasons, you can't use phenolphthalein in titrations involving ammonium hydroxide solution. Methyl orange works OK with all of these possible combinations, and that is the one you should choose if you have the option. It is important that you remember its colour changes.

Apart than that, the experimental detail, including the crystallisation, is the same for all of these.