E. Solutions and Solubility · Solutions and SolubilityThe Solubility Table, Precipitates and Testing for Ions
E3.4 E2 E2.1 A1.2 A1.5About 75 minutes
Learning goals- I can use a solubility table to classify ionic compounds as high or low solubility.
- I can predict whether a precipitate forms when two solutions are mixed, and name it.
- I can explain how precipitation is used to remove ions (such as phosphate) from water.
- I can plan simple qualitative tests (precipitates, flame colours, solution colours) to identify ions in a solution.
Words to know
precipitate
An insoluble solid that forms when two solutions are mixed; written with (s).
precipitation reaction
A double displacement reaction in which ions combine to form a low-solubility solid.
solubility table
A chart that shows which combinations of ions have high or low solubility in water.
qualitative analysis
Identifying which substances (e.g. which ions) are present, without measuring how much.
quantitative analysis
Measuring how much of a substance is present.
flame test
Identifying certain metal ions by the colour they give a flame.
filtrate
The liquid that passes through a filter; the solid trapped on the filter is the residue.
Which ions refuse to stay dissolved?
When you mix two clear solutions and a cloudy solid suddenly appears, some ions have found partners they cannot let go of. In Grade 10 you met precipitates in double displacement reactions and used a simple solubility guide. Now you will use a full solubility table to predict precipitates, to remove unwanted ions from water, and to identify unknown ions, as water-testing labs do.
Remember from the Chemical Reactions unit
In a double displacement reaction, the cations of two compounds swap partners: AB + CD → AD + CB. A double displacement between two ionic solutions happens when a product is a precipitate, a gas, or water.
Reading the solubility table
The table is organized by anion. To decide whether a compound has high or low solubility:
- Find the compound's anion in the left column.
- Check whether its cation is listed in the high-solubility column or the low-solubility column.
- "Most" means: any cation not named in the other column.
Two shortcuts are worth memorizing: all nitrates and almost all Group 1 (alkali metal) and ammonium compounds have high solubility.
In this course, high solubility means at least 0.1 mol/L dissolves at 25 °C; low solubility means less than 0.1 mol/L. Low-solubility compounds are written with (s) in equations; high-solubility ones dissolved in water are written with (aq).
| Anion | High solubility (aq): ≥ 0.1 mol/L at 25 °C | Low solubility (s): < 0.1 mol/L at 25 °C |
|---|
| Cl⁻, Br⁻, I⁻ | most | Ag⁺, Pb²⁺, Tl⁺, Hg₂²⁺, Cu⁺ |
| SO₄²⁻ | most | Ag⁺, Ca²⁺, Sr²⁺, Ba²⁺, Pb²⁺, Ra²⁺ |
| S²⁻ | Group 1 ions, Group 2 ions, NH₄⁺ | most others |
| OH⁻ | Group 1 ions, NH₄⁺, Sr²⁺, Ba²⁺, Tl⁺ | most others (including Ca²⁺, Mg²⁺) |
| PO₄³⁻, SO₃²⁻, CO₃²⁻ | Group 1 ions, NH₄⁺ | most others |
| CH₃COO⁻ (acetate) | most | Ag⁺ |
| NO₃⁻ | all | none |
| any anion with Group 1 ions, H⁺ or NH₄⁺ | almost all | — |
Worked exampleExample 1: Using the table
Classify each compound as high or low solubility: (a) CaCO₃, (b) K₃PO₄, (c) BaSO₄, (d) PbI₂, (e) Mg(OH)₂, (f) Cu(NO₃)₂.
- (a) CaCO₃: anion CO₃²⁻; Ca²⁺ is not Group 1 or NH₄⁺, so low (limestone).
- (b) K₃PO₄: K⁺ is Group 1, so high.
- (c) BaSO₄: Ba²⁺ is an exception for sulfate, so low.
- (d) PbI₂: Pb²⁺ is an exception for iodide, so low (a yellow solid).
- (e) Mg(OH)₂: hydroxides are high only for Group 1, NH₄⁺, Sr²⁺, Ba²⁺, Tl⁺; Mg²⁺ is not listed, so low.
- (f) Cu(NO₃)₂: all nitrates are high.
Low: CaCO₃, BaSO₄, PbI₂, Mg(OH)₂. High: K₃PO₄, Cu(NO₃)₂.
Predicting precipitates
- Write the ions present in each solution.
- Swap partners: pair each cation with the other solution's anion.
- Check each new combination in the solubility table. A low-solubility pair is the precipitate.
- Write a balanced equation with states. If both new pairs have high solubility, no reaction is observed (the ions just mix).
Worked exampleExample 2: Predicting the product
Predict whether a precipitate forms, and write the balanced equation: (a) Pb(NO₃)₂(aq) + KI(aq); (b) NaCl(aq) + KNO₃(aq); (c) FeCl₃(aq) + Na₃PO₄(aq).
- (a) Swap: PbI₂ and KNO₃. PbI₂ is low (Pb²⁺ with I⁻). Pb(NO₃)₂(aq) + 2 KI(aq) → PbI₂(s) + 2 KNO₃(aq). A bright yellow precipitate.
- (b) Swap: NaNO₃ and KCl. Both are high (nitrate; Group 1). No precipitate, no reaction.
- (c) Swap: FePO₄ and NaCl. FePO₄ is low (phosphate with a non-Group 1 cation). FeCl₃(aq) + Na₃PO₄(aq) → FePO₄(s) + 3 NaCl(aq).
(a) PbI₂(s); (b) no reaction; (c) FePO₄(s).
Using precipitation to clean water
Phosphate in sewage causes eutrophication (see the water pollution lesson). Many Ontario sewage treatment plants remove it by adding a metal salt such as iron(III) chloride, iron(II) or iron(III) sulfate, or aluminium sulfate (alum). The metal ions form low-solubility phosphates that settle out with the sludge:
Fe³⁺(aq) + PO₄³⁻(aq) → FePO₄(s)
Al³⁺(aq) + PO₄³⁻(aq) → AlPO₄(s)
Precipitation is also used to remove hardness ions (Ca²⁺, Mg²⁺) from water and to remove dissolved metals from mine and factory wastewater, for example as low-solubility hydroxides.
Qualitative analysis: which ions are in there?
Chemists identify ions using properties they can observe (part of overall expectation E2: investigating the properties of solutions):
1. Colour of the solution. Many transition-metal ions are coloured: Cu²⁺(aq) is blue, Ni²⁺(aq) green, Fe³⁺(aq) yellow-orange to brown, MnO₄⁻(aq) purple, CrO₄²⁻(aq) yellow, Cr₂O₇²⁻(aq) orange. Ions of Group 1 and 2 metals and most common anions are colourless.
2. Flame tests. Heated metal ions give characteristic colours: Li⁺ red, Na⁺ yellow-orange, K⁺ violet (lilac), Ca²⁺ orange-red, Sr²⁺ red, Cu²⁺ blue-green. (These are the colours of fireworks.)
3. Precipitation tests. Add a solution that forms a precipitate with only the ion you are looking for:
- Cl⁻: add AgNO₃(aq) → white AgCl(s).
- SO₄²⁻: add BaCl₂(aq) → white BaSO₄(s).
- Pb²⁺: add KI(aq) → yellow PbI₂(s).
- CO₃²⁻: add acid → bubbles of CO₂ gas.
When several ions might be present, test in a careful sequence and remove each precipitate by filtration before the next test, so one ion does not hide another.
Just for funFireworks are flame tests with a bigger budget. The reds come from strontium compounds, the greens from barium compounds, and the yellow-orange from sodium. Chemists call it qualitative analysis; everyone else calls it Canada Day.
Hands-on activityA safe precipitate at home
You need: Epsom salt (magnesium sulfate), washing soda (sodium carbonate; it is an irritant, so a grown-up handles it, with gloves), two clear glasses, warm tap water, spoons, a coffee filter and a funnel or sieve, safety glasses if you have them.
- Glass 1: dissolve 1 tablespoon of Epsom salt in 100 mL of warm water. Glass 2: a grown-up dissolves 1 tablespoon of washing soda in 100 mL of warm water.
- Predict: swap partners in MgSO₄ + Na₂CO₃. Use the solubility table to predict which product is the precipitate.
- Pour Glass 2 slowly into Glass 1. Observe.
- Filter the mixture through the coffee filter. Look at the residue and the filtrate.
- Write the balanced equation with states.
What to look for: a thick white precipitate of magnesium carbonate forms immediately. MgSO₄(aq) + Na₂CO₃(aq) → MgCO₃(s) + Na₂SO₄(aq). The filtrate is clear: it contains dissolved sodium sulfate. Do not taste anything; rinse everything down the drain with lots of water and wash hands. Never use lab chemicals such as silver nitrate or lead nitrate at home.
Your turn
1Classify as high or low solubility: (a) AgBr, (b) (NH₄)₂CO₃, (c) Ca(OH)₂, (d) Zn(NO₃)₂.
(a) low (Ag⁺ with Br⁻); (b) high (NH₄⁺); (c) low (hydroxide, Ca²⁺ not in the high list); (d) high (nitrate).
2Predict the precipitate: BaCl₂(aq) + Na₂SO₄(aq).
Swap: BaSO₄ and NaCl. BaSO₄(s) is low solubility. BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2 NaCl(aq).
3Predict whether a precipitate forms: KBr(aq) + NH₄NO₃(aq).
Swap: KNO₃ and NH₄Br. Both high: no precipitate.
4Write the balanced equation for CuSO₄(aq) + NaOH(aq).
CuSO₄(aq) + 2 NaOH(aq) → Cu(OH)₂(s) + Na₂SO₄(aq). (Cu(OH)₂ is a blue precipitate.)
5A colourless solution gives a white precipitate with AgNO₃(aq). Which anions could be present (from Cl⁻, NO₃⁻, SO₄²⁻)?
Cl⁻ (AgCl). Silver sulfate also has low solubility in the table, so SO₄²⁻ could also give a precipitate with enough Ag⁺; a BaCl₂ test would tell them apart. Nitrate never precipitates.
6Which flame colour identifies K⁺? Na⁺?
K⁺: violet (lilac). Na⁺: yellow-orange.
7Why do sewage plants add iron(III) chloride?
Fe³⁺ forms low-solubility FePO₄ with phosphate, which settles out with the sludge, so less phosphate reaches lakes.
More practice: warm-up
1Are all nitrates high solubility?
2What does (s) after a product mean in a precipitation equation?
The product is a solid precipitate.
4Name the precipitate in the test for sulfate.
5Is NaOH high or low solubility?
More practice: on your own
1Predict the products and write balanced equations: (a) AgNO₃(aq) + K₂CrO₄(aq) (Ag₂CrO₄ is low solubility); (b) Pb(NO₃)₂(aq) + NaCl(aq).
(a) 2 AgNO₃(aq) + K₂CrO₄(aq) → Ag₂CrO₄(s) + 2 KNO₃(aq). (b) Pb(NO₃)₂(aq) + 2 NaCl(aq) → PbCl₂(s) + 2 NaNO₃(aq).
2Choose a solution to remove Pb²⁺ ions from wastewater by precipitation, without adding a toxic ion. Write the equation.
Sample: add sodium sulfate (or sodium carbonate): Pb²⁺(aq) + SO₄²⁻(aq) → PbSO₄(s). Sodium and sulfate are low-toxicity ions.
3A solution might contain Ba²⁺ or Mg²⁺ (or both). Design a test to tell them apart, using the solubility table.
Add Na₂SO₄(aq): a white precipitate means Ba²⁺ (BaSO₄); MgSO₄ is high solubility, so Mg²⁺ gives none. Filter, then add NaOH(aq) to the filtrate: a white precipitate means Mg²⁺ (Mg(OH)₂). (Ba(OH)₂ is high solubility, so any Ba²⁺ left would not interfere.)
4Hard water contains Ca²⁺. Washing soda (Na₂CO₃) is added to laundry to soften it. Explain with an equation.
Ca²⁺(aq) + CO₃²⁻(aq) → CaCO₃(s). Calcium ions are removed as a precipitate, so they cannot react with soap to form soap scum.
5Why must you use a fresh sample (or filter) between tests in a sequence?
A precipitate or added reagent from the first test could interfere with or hide the result of the next test; filtering removes the first precipitate so the next result is clear.
Challenge
1Three unlabelled bottles hold NaCl(aq), Na₂SO₄(aq) and Na₂CO₃(aq). You may use dilute HCl(aq) and BaCl₂(aq). Describe how to identify all three.
Add HCl to a sample of each: bubbles (CO₂) identify Na₂CO₃. To fresh samples of the other two add BaCl₂: a white precipitate (BaSO₄) identifies Na₂SO₄; no precipitate means NaCl.
2A solution contains Ag⁺, Ba²⁺ and Cu²⁺. Plan a sequence (using NaCl, Na₂SO₄ and NaOH solutions) that removes them one at a time. Explain the order.
- Add NaCl: only AgCl precipitates (BaCl₂ and CuCl₂ are high). Filter. 2. Add Na₂SO₄ to the filtrate: BaSO₄ precipitates (CuSO₄ is high). Filter. 3. Add NaOH: blue Cu(OH)₂ precipitates. NaOH must come last, because Ag⁺ would also form a precipitate with OH⁻ (AgOH/Ag₂O); and Na₂SO₄ before NaOH is fine because Ba(OH)₂ is high solubility.
3Why would iron(III) sulfate be a good choice for removing phosphate at a sewage plant, but barium chloride a poor choice, even though both could form a precipitate with something in the water?
Fe³⁺ forms low-solubility FePO₄, removing phosphate, and iron and sulfate are relatively low-toxicity. Barium phosphate would also precipitate, but excess Ba²⁺ is toxic and would be released into the river; a treatment chemical must not add a new pollutant.
Quiz
1. Which compound has low solubility?
2. Mixing Pb(NO₃)₂(aq) and KI(aq) produces a precipitate of
3. Which pair of solutions gives no precipitate?
5. Which ion is used in water treatment to precipitate phosphate?
6. A solution is blue and gives a blue precipitate with NaOH. It likely contains
Big idea (E3.4, E2): the solubility table predicts which ion pairs form precipitates; precipitation is used to identify ions and to remove them from water.
How to help: Give your teen two formulas and ask them to 'swap partners' out loud and check the table. Ask which ion a test is looking for and why that reagent is chosen.
Common difficulties:
- Looking up the cation row instead of the anion row.
- Forgetting that 'most' means 'all except the ones listed'.
- Unbalanced charges in new formulas (Ba²⁺ + PO₄³⁻ gives Ba₃(PO₄)₂).
- Planning tests in an order where one precipitate hides another.
Safety: washing soda is an irritant: a grown-up handles it, with gloves. Lab reagents such as silver nitrate, barium chloride and lead nitrate are toxic or corrosive and are used only in a supervised school lab.