Precipitation Reactions and Testing for Ions, Grade 10
You cannot see a chloride ion. Add a few drops of silver nitrate and you can.
A solution of sodium chloride and a solution of sodium nitrate look exactly the same. Both are clear, both are colourless, and no amount of staring will separate them. Add silver nitrate to one of them and a white solid appears that you can point at, filter out and weigh.
That is what this whole unit is for. A precipitation reaction is the tool that turns an invisible ion into something visible, and everything else on this page is the machinery that makes it predictable.
What a precipitate is, briefly
A precipitate is an insoluble solid that appears when two solutions are mixed. The two solutions each contained dissolved ions. When they meet, one pair of those ions forms a compound that will not dissolve, so it drops out of solution as a solid.
AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)
The (s) is the entire result of the experiment. It is one letter and it is the thing that gets left off in exams more than any other symbol on the page.
How to predict one, in three steps
This is a method, not a guess, and learners who follow it get every prediction question right.
- Write down the four ions that are in the two solutions
- Swap the partners. The cation from one goes with the anion from the other, and the other way round
- Check both new pairings against the solubility rules. Anything insoluble is your precipitate. Anything soluble was never anywhere except in solution
Worked, with silver nitrate and sodium chloride
| Step | |
|---|---|
| The four ions | Ag+, NO3-, Na+, Cl- |
| Swap the partners | AgCl and NaNO3 |
| Check the rules | Chlorides are soluble except silver, so AgCl comes out as a solid. Every sodium salt and every nitrate is soluble, so NaNO3 stayed dissolved |
The solubility rules
| Ion | Rule |
|---|---|
| Na+, K+ and NH4+ salts | Always soluble |
| Nitrates and acetates | Always soluble |
| Chlorides, bromides and iodides | Soluble, except silver, lead and mercury(I) |
| Sulfates | Soluble, except barium, lead and strontium |
| Hydroxides | Mostly insoluble. Group 1 dissolves, and barium and calcium are moderately soluble |
| Carbonates | Mostly insoluble, except Group 1 and ammonium |
| Sulfides | Mostly insoluble, except Groups 1 and 2 and ammonium |
The first two rules do most of the work. If a compound has sodium, potassium, ammonium or nitrate in it, it dissolves, and you can stop reading the table. Rules 3 to 7 are where every precipitate in this chapter comes from.
Spectator ions, and the equation that tells the truth
There are three ways to write the same reaction, and they are not equally honest.
| Silver nitrate and sodium chloride | |
|---|---|
| Balanced equation | AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq) |
| Full ionic | Ag+ + NO3- + Na+ + Cl- → AgCl(s) + Na+ + NO3- |
| Net ionic | Ag+(aq) + Cl-(aq) → AgCl(s) |
Look at the full ionic equation and find what appears unchanged on both sides. Na+ and NO3- were free ions in solution before the mixing and they are free ions in solution after it. Nothing happened to them at all.
Those are the spectator ions. Cross them out and what is left is the net ionic equation, which says what actually took place: a silver ion met a chloride ion and the two of them stopped being dissolved.
A common way to lose a mark is to write out the full ionic equation when the question asked for the net one. If your answer still has an ion sitting on both sides, you have not finished.
Where this fits in the curriculum
| Subject | Physical Sciences |
|---|---|
| Grade | 10 |
| Term | 2 |
| Topic | Chemical change, reactions in aqueous solution |
| Status | Class experiments. Not a formal assessment |
| Marks | 55 on our worksheet. None is prescribed |
Practical 1: testing for chloride, bromide and iodide
One reagent identifies all three, and it does it by colour.
Put each of the three solutions into its own labelled test tube to a depth of about a centimetre, using a separate dropper for each. Add an equal amount of silver nitrate to all three, shake each one gently from side to side, and let them settle.
| Ion present | Precipitate | Colour | In dilute nitric acid |
|---|---|---|---|
| Cl- | AgCl | White | Does not dissolve |
| Br- | AgBr | Cream | Does not dissolve |
| I- | AgI | Yellow, and clearly deeper | Does not dissolve |
Cl-(aq) + Ag+(aq) → AgCl(s)
The three tubes have to stand side by side or the test does not work. "Cream" means nothing on its own. It means something next to a white one and a yellow one, and the order down the group is the whole point: the further down you go, the deeper the colour.
Never put a thumb over the top of a test tube and turn it upside down. Shake it side to side. That instruction is in the textbook and it is right.
A white precipitate is not automatically chloride
Carbonate gives one too, and this is where the test needs a second step.
CO32-(aq) + 2Ag+(aq) → Ag2CO3(s)
Most textbooks call silver carbonate white and leave it at that. It is actually pale yellow, so next to a genuinely white silver chloride there is already a visible difference. The colour is a clue. The acid is the proof.
Pour off the liquid above the precipitate and add a few drops of dilute nitric acid to the solid. Silver chloride sits there and does nothing. Silver carbonate fizzes, gives off carbon dioxide and dissolves.
Ag2CO3(s) + 2HNO3(aq) → 2Ag+(aq) + CO2(g) + H2O(l) + 2NO3-(aq)
Practical 2: testing for sulfate and carbonate
Barium nitrate does for these two what silver nitrate did for the halides, and it runs into exactly the same problem.
Quarter-fill one tube with sodium sulfate and another with sodium carbonate. Add an equal amount of barium nitrate to each.
| Ion present | Precipitate | Colour | In dilute nitric acid |
|---|---|---|---|
| SO42- | BaSO4 | White | Nothing. It sits there |
| CO32- | BaCO3 | White | Fizzes and dissolves |
BaCO3(s) + 2HNO3(aq) → Ba2+(aq) + 2NO3-(aq) + H2O(l) + CO2(g)
One reagent solves both problems, and this is the thing worth remembering
Textbooks teach these two ambiguities four pages apart, as though they were separate facts. They are the same fact.
| Test | Two white solids | The one that fizzes and dissolves |
|---|---|---|
| Silver nitrate | AgCl or Ag2CO3 | The carbonate |
| Barium nitrate | BaSO4 or BaCO3 | The carbonate |
Carbonate reacts with any acid to give carbon dioxide. That is why it fizzes, and the fizzing is why the solid disappears. Nothing else in either test does that.
Learn the mechanism once instead of memorising four outcomes. If a white precipitate dissolves in acid with bubbles, it was a carbonate, whichever reagent made it.
Practical 3: making a salt on purpose
The same reaction that identifies an ion also manufactures a compound.
- Measure 10 ml of sodium carbonate solution into a small beaker, then measure 10 ml of copper(II) nitrate solution with a second cylinder and add it
- Stir. A blue-green solid appears at once
- Fold a filter paper in half, then in quarters, open it into a cone and sit it in the funnel, with the funnel in a conical flask
- Pour the mixture through. The solid stays on the paper, the sodium nitrate solution runs into the flask below
- Rinse the solid on the paper with a little distilled water, then leave it to dry in the air and transfer the powder into a labelled sample tube
What you actually made is not what the textbook says
Most Grade 10 books call this copper(II) carbonate, CuCO3. It is not.
Mixing a copper solution with a carbonate solution in water gives basic copper carbonate, Cu2(OH)2CO3. Plain CuCO3 cannot be precipitated from water at all. That is why the solid comes out blue-green rather than the pale blue you were promised, and it is why a drop of acid on it fizzes.
The correction is a better lesson than the original. Cu2(OH)2CO3 is malachite: the green stone in copper jewellery, and the same compound that turns old copper roofs and church spires green. Your class has just made a gemstone out of two clear liquids.
What you need
| Item | Qty | Why |
|---|---|---|
| Test tubes, borosilicate 3.3, rimmed | 12 | Three halides, two anions, spares, and the comparison needs them together |
| Test tube rack, 6 hole, with pegs | 2 | The rack is the comparison. Six tubes side by side against white paper |
| Dropper bottles, 5 ml, pack of 10 | 1 pk | One per reagent, labelled. Every step in this practical is "add a few drops" |
| Disposable pipettes, pack of 100 | 1 pk | A clean one per tube. Cross-contamination is the commonest failure |
| Beaker, borosilicate low form, 250 ml | 2 | Mixing for the salt preparation |
| Conical flask, narrow neck | 1 | Catches the filtrate |
| Filter funnel, borosilicate glass | 1 | Separating the precipitate |
| Qualitative filter paper, 90 mm | 1 pk | 100 circles. One per group per run |
| Measuring cylinder, glass class B, 100 ml | 2 | Two, so the copper solution never touches the carbonate cylinder |
| Stirring rod, borosilicate | 2 | |
| Spatula, stainless steel, spoon end | 1 | |
| Watch glass | 1 | Drying the finished salt |
The dropper bottles are the method, not an accessory. Ten labelled squeezable bottles are the difference between a lesson that runs and thirty learners queueing at one reagent with one pipette. Never swap the tops. One silver nitrate top in the potassium iodide bottle ruins the reagent for the next four classes.
The solutions you supply yourself
| Reagent | Concentration | Used in |
|---|---|---|
| Silver nitrate | 0,02 mol·dm-3 | Practical 1 |
| Sodium chloride, potassium bromide, potassium iodide | 0,1 mol·dm-3 | Practical 1 |
| Dilute nitric acid | 0,5 mol·dm-3 | Practicals 1 and 2 |
| Barium nitrate | 0,1 mol·dm-3 | Practical 2 |
| Sodium sulfate, sodium carbonate | 0,1 mol·dm-3 | Practicals 2 and 3 |
| Copper(II) nitrate | 0,1 mol·dm-3 | Practical 3 |
Some books specify 2 mol·dm-3 nitric acid for the second practical. Use 0,5. The job is to dissolve a carbonate and produce visible bubbles, which happens easily at 0,5, and 2 mol·dm-3 nitric acid in a Grade 10 class is more hazard than the result requires.
Two reagents we will not use, and what to use instead
| Refused | Why | Use instead |
|---|---|---|
| Lead(II) nitrate | Soluble lead salts are cumulative toxins, restricted in school laboratories, and the waste cannot go down a drain. It appears in the standard paper exercise on this topic | Barium nitrate and sodium sulfate. Same rule, same prediction, same white precipitate |
| Sodium sulfide | Releases hydrogen sulfide on contact with any acid. It is the worked example in several textbooks | Keep it as a paper example only. Never take it off the page |
A written exercise that names a banned reagent teaches the class to reach for it later. Predicting the lead iodide precipitate on paper is chemically fine and pedagogically careless.
What you should see
- Three different colours in the halide tubes, in order: white, cream, yellow
- None of the three silver halides dissolves in nitric acid. All three stay put
- Two identical white solids in the barium tubes, and only one of them fizzes
- A blue-green solid in the beaker within seconds of the two clear solutions meeting
- A colourless filtrate running into the conical flask. If it is blue, the filtering has gone wrong
No precipitate is also a result
Mix sodium nitrate with potassium chloride and nothing happens. Both possible products, sodium chloride and potassium nitrate, are soluble, so there is nothing to come out of solution.
A learner who writes "nothing happened, so it failed" has thrown away a correct answer. In qualitative analysis a negative result carries exactly as much information as a positive one: the ion you were testing for was not there. The test worked. The answer was no.
If it does not work
| What happens | What caused it |
|---|---|
| All three halide tubes look the same | They are being compared one at a time. Stand all three in the rack together against white paper |
| Everything goes cloudy, including the acid | Cross-contaminated droppers. A clean dropper per reagent, every time. This is the commonest failure in the whole practical |
| No precipitate at all with silver nitrate | Old silver nitrate, or a tube that got water instead of solution. Silver nitrate decomposes in light, so keep it in a brown bottle |
| The precipitate dissolves in the acid | It was a carbonate. That is a correct result, not a failure |
| The barium tubes both fizz | The sulfate solution has carbonate in it, usually from a shared dropper |
| The filtrate runs through blue | The filter paper is torn, badly seated, or the liquid went over its top edge. The paper must sit below the rim of the funnel |
| Filtering takes the whole period | Let the solid settle and decant the clear liquid first, then pour the sludge |
| The copper salt will not dry | It is still wet with sodium nitrate solution. Rinse it on the paper with a little distilled water |
| The copper solid is green, not blue | Correct. It is basic copper carbonate |
| A learner's fingers turn black | Silver nitrate. Harmless, and it wears off in about a week |
Safety
This is not a mild practical. It contains a corrosive and a poison, and it needs to be run as though it does.
- Goggles on, without exception. Silver nitrate and nitric acid
- Silver nitrate stains skin black for about a week. It is harmless and it looks alarming. Announce it before the lesson rather than after, because a class that has been warned is careful in a way that a general instruction never achieves
- Barium nitrate is toxic if swallowed. Soluble barium salts are poisons
- Two residue bottles, clearly labelled, one for silver waste and one for barium. Neither goes down the sink
- Shake test tubes from side to side. Never invert one with a thumb over the top
- No lead salts, in the practical or in the written exercises
Disposal: silver and barium residues into their bottles for collection. Everything else down the sink with plenty of water.
How the 55 marks are made up
| Section | Marks |
|---|---|
| Predictions, made before the practical | 6 |
| The halide results | 7 |
| Equations, including the net ionic | 10 |
| Predicting a precipitate from the rules | 8 |
| Two white solids, and telling them apart | 8 |
| Carbonate and sulfate results | 4 |
| Making a salt, and why precipitation cleans water | 6 |
| Conclusion | 6 |
No mark allocation is prescribed for this practical. The worksheet and this split are ours.
The mark most often dropped is the one for a pair of solutions that gives no precipitate. Learners write an equation with a solid in it because the question looked like all the others, without ever checking both products against the rules.
Close behind: the state symbol on the precipitate, and net ionic equations that still have the spectator ions in them.
If you have time
Give each group an unknown solution and the two reagents. Nothing else. They have to work out which ion is in it and justify the answer from their own observations. It takes ten minutes and it is worth more than the three prescribed practicals put together, because the method finally has a job.
Ask why municipal water treatment uses precipitation. A dissolved metal cannot be filtered out. Turn it into an insoluble compound and it can. That is the copper experiment at industrial scale.
Leave the silver chloride tube on a sunny windowsill. It darkens to grey and then purple as the light reduces it to silver metal, which is the reaction that photographic film was built on.
Free worksheet and marking memo
Both free, no sign up, straight to the PDF.
- Learner worksheet, 55 marks, with the prediction table, the halide results, four equations to write, the two-white-solids problem and the negative-result question
- Marking memorandum, with the expected colours, how to mark a net ionic equation that still has spectators in it, and the three answers that look right and score nothing
Related practicals
- Solubility and dissolving, Grade 10. Where the solubility rules come from, and why some salts refuse to dissolve at all
- Flame tests, Grade 10. Cations there, anions here. The two together are qualitative analysis
- Separating mixtures, Grade 10. Filtration, which the salt preparation depends on
- Physical and chemical change, Grade 10. A precipitate is one of the signs of a chemical change
Buy this experiment
We are putting together a Precipitation and Ion Testing Kit for Grade 10 with the test tubes, racks, dropper bottles, pipettes, beakers, conical flask, funnel, filter paper, cylinders, rods and watch glass in one box, plus a printed teacher guide and the marking memo. The chemicals are not in it yet, and we would rather say so than ship a box that cannot run the practical.
The dropper bottles and the racks are the two lines this practical genuinely cannot do without. They are in the test tube racks range and alongside the glass test tubes.
Silver nitrate has no substitute. Unlike most Grade 10 chemistry, there is no store cupboard version of this practical, and the contents card says exactly which nine solutions a school needs to have before the box is worth opening.