Qualitative Analysis

Define qualitative analysis.

An analytical technique used to determine the identity of chemical species present in a sample; it indicates whether a particular ion or substance is present, but does not measure how much is present or its concentration.

Define a precipitate.

An insoluble solid formed when two solutions are mixed together.

Define a cation.

A positively charged ion.

Define an anion.

A negatively charged ion.

State the test for carbonate ions (CO₃²⁻).

Add dilute nitric acid (HNO₃) to the solid or aqueous sample; effervescence/bubbles of gas produced; bubble the gas through limewater which turns cloudy/milky due to formation of a white precipitate of calcium carbonate.

State the ionic equation for the test for carbonate ions.

CO₃²⁻(aq) + 2H⁺(aq) → CO₂(g) + H₂O(l).

State the equation for the confirmation of carbon dioxide with limewater.

Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l); OR CO₂(g) + Ca²⁺(aq) + 2OH⁻(aq) → CaCO₃(s) + H₂O(l).

State the observation when carbonate ions react with dilute acid.

Effervescence/fizzing/bubbles of gas due to carbon dioxide being produced; the gas turns limewater milky/cloudy.

Explain why nitric acid is used for the carbonate test rather than sulfuric or hydrochloric acid.

Sulfuric acid contains sulfate ions (SO₄²⁻) which would give a false positive with barium nitrate in the sulfate test; hydrochloric acid contains chloride ions (Cl⁻) which would give a false positive with silver nitrate in the halide test; nitric acid does not introduce interfering ions.

State the test for sulfate ions (SO₄²⁻).

Acidify the sample with dilute hydrochloric acid or nitric acid to remove interfering ions; then add a few drops of aqueous barium chloride (BaCl₂) or barium nitrate (Ba(NO₃)₂); a dense white precipitate of barium sulfate forms.

State the ionic equation for the test for sulfate ions.

Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s).

State the observation when sulfate ions react with barium chloride/nitrate.

A dense white precipitate forms; barium sulfate is insoluble in acid.

Explain why acid is added before testing for sulfate ions.

To remove any carbonate or sulfite ions that may interfere with the result; carbonate and sulfite ions would also form precipitates with Ba²⁺, causing false positive results; the acid dissolves these precipitates.

Explain why barium nitrate is used rather than barium chloride when testing a mixture of anions.

Barium chloride would introduce chloride ions (Cl⁻) into the solution; chloride ions would give a false positive in the subsequent halide test with silver nitrate; barium nitrate does not introduce interfering ions.

Explain why barium chloride is acceptable when testing a single salt.

If only testing for sulfate ions in a single salt, there is no subsequent halide test; introducing chloride ions does not matter as the halide test will not be carried out.

State the test for halide ions (Cl⁻, Br⁻, I⁻).

Add dilute nitric acid followed by aqueous silver nitrate (AgNO₃); a coloured precipitate of silver halide forms; add dilute then concentrated ammonia to confirm the halide present.

State the ionic equation for the formation of silver halide precipitates.

Ag⁺(aq) + X⁻(aq) → AgX(s) where X = Cl, Br or I.

State the colour of silver chloride precipitate (AgCl).

White.

State the colour of silver bromide precipitate (AgBr).

Cream.

State the colour of silver iodide precipitate (AgI).

Pale yellow.

State the ionic equation for the formation of silver chloride.

Ag⁺(aq) + Cl⁻(aq) → AgCl(s).

State the ionic equation for the formation of silver bromide.

Ag⁺(aq) + Br⁻(aq) → AgBr(s).

State the ionic equation for the formation of silver iodide.

Ag⁺(aq) + I⁻(aq) → AgI(s).

State the solubility of silver chloride in ammonia.

Soluble in dilute ammonia; dissolves to form a colourless solution containing the [Ag(NH₃)₂]⁺ complex ion.

State the solubility of silver bromide in ammonia.

Insoluble in dilute ammonia; soluble in concentrated ammonia; dissolves to form a colourless solution.

State the solubility of silver iodide in ammonia.

Insoluble in both dilute and concentrated ammonia; pale yellow precipitate remains.

Explain how to distinguish between chloride, bromide and iodide ions using ammonia.

Add dilute ammonia: if precipitate dissolves it is Cl⁻; if not add concentrated ammonia: if precipitate dissolves it is Br⁻; if precipitate does not dissolve in either it is I⁻.

Explain the solubility trend of silver halides in ammonia.

Solubility decreases down the group: AgCl dissolves in dilute NH₃; AgBr requires concentrated NH₃; AgI is insoluble even in concentrated NH₃; due to decreasing solubility product down the group.

Explain why silver chloride dissolves in ammonia but silver iodide does not.

AgCl has a higher solubility product than AgI; AgCl is more soluble; the formation of the [Ag(NH₃)₂]⁺ complex ion is sufficient to dissolve AgCl but not the less soluble AgI; AgI has the lowest solubility product.

State the equation for the dissolution of silver chloride in ammonia.

AgCl(s) + 2NH₃(aq) → [Ag(NH₃)₂]⁺(aq) + Cl⁻(aq).

State the test for ammonium ions (NH₄⁺).

Add aqueous sodium hydroxide (NaOH) to the sample and warm gently in a water bath; ammonia gas is produced; test with damp red litmus paper which turns blue OR use a glass rod dipped in concentrated HCl which forms white fumes of NH₄Cl.

State the ionic equation for the test for ammonium ions.

NH₄⁺(aq) + OH⁻(aq) → NH₃(g) + H₂O(l).

State the observation when ammonium ions react with warm NaOH.

Ammonia gas is produced (pungent smell); damp red litmus paper turns blue (alkaline gas); a glass rod dipped in concentrated HCl produces white fumes of ammonium chloride.

Explain why the ammonium ion test requires warming.

The reaction between NH₄⁺ and OH⁻ produces ammonia gas; warming drives off the NH₃ gas; the equilibrium shifts to the right; ammonia is very soluble in water so warming is needed to evolve the gas.

State the test for hydroxide ions (OH⁻).

Test the pH using red litmus paper (turns blue), universal indicator (turns blue/purple), or a pH probe (gives a high pH reading).

State the test for hydrogen ions (H⁺).

Test the pH using blue litmus paper (turns red), universal indicator (turns red), or a pH probe (gives a low pH reading); OR add carbonate which produces effervescence (CO₂).

State the sequence of tests for anions in an unknown compound.

1. Carbonate (CO₃²⁻); 2. Sulfate (SO₄²⁻); 3. Halide (Cl⁻, Br⁻, I⁻).

Explain why the carbonate test must be carried out first.

Carbonates produce CO₂ with acid; if there is no effervescence, no carbonate is present; carbonates would interfere with the sulfate test as BaCO₃ is also insoluble and would form a white precipitate, giving a false positive.

Explain why the sulfate test must be carried out before the halide test.

Silver sulfate (Ag₂SO₄) is slightly insoluble and would form a white precipitate; this would interfere with the halide test causing a false positive for chloride; removing sulfate first ensures accurate results.

Explain why the halide test is carried out last.

Silver carbonate (Ag₂CO₃) and silver sulfate (Ag₂SO₄) are also insoluble; carrying out the halide test last ensures carbonate and sulfate have been removed; prevents false positive results.

Explain how to test a mixture of anions.

Step 1: Carbonate test - add dilute HNO₃ until bubbling stops to remove all CO₃²⁻; Step 2: Sulfate test - add excess Ba(NO₃)₂, filter to remove BaSO₄ precipitate; Step 3: Halide test - test the filtrate with AgNO₃ and NH₃.

Explain why you must add excess barium nitrate in the sulfate test.

To ensure all sulfate ions are precipitated; any remaining sulfate would interfere with the halide test by forming Ag₂SO₄.

Explain why you must filter after the sulfate test before testing for halides.

To remove the barium sulfate precipitate; the precipitate would interfere with the halide test; the filtrate contains the remaining ions for testing.

Explain why nitric acid is used for acidification in the sequential anion tests.

Nitric acid does not introduce interfering ions; sulfuric acid would introduce SO₄²⁻; hydrochloric acid would introduce Cl⁻; both would cause false positives in subsequent tests.

State the acid used for acidifying solutions in anion tests.

Dilute nitric acid (HNO₃).

Explain why silver nitrate is used to test for halide ions.

Ag⁺ ions react with halide ions to form insoluble precipitates; the precipitates have different colours (white, cream, yellow); the precipitates have different solubilities in ammonia which confirms their identity.

Explain why barium chloride/nitrate is used to test for sulfate ions.

Ba²⁺ ions react with sulfate ions to form insoluble barium sulfate; BaSO₄ is a dense white precipitate that is insoluble in acid; the white precipitate is characteristic and confirms sulfate ions.

Explain why a white precipitate forms when barium chloride is added to a sulfate solution.

Ba²⁺ ions and SO₄²⁻ ions combine; BaSO₄ is insoluble in water; Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s); the insoluble compound precipitates out of solution.

Describe the method for testing Group 2 metal ions (PAG 4).

Place clean test tubes in a rack; add 10 drops of 0.1 mol dm⁻³ metal chloride solution (Mg²⁺, Ca²⁺, Sr²⁺, Ba²⁺); add 10 drops of dilute NaOH using a pipette; swirl gently to mix; record observations; repeat with dilute ammonia solution; repeat with dilute sulfuric acid.

State the results of adding excess NaOH to Group 2 ions.

Mg²⁺: white precipitate of Mg(OH)₂ (insoluble in excess NaOH); Ca²⁺: white precipitate of Ca(OH)₂ (slightly soluble); Sr²⁺: slight white precipitate of Sr(OH)₂; Ba²⁺: no change (Ba(OH)₂ is soluble).

State the results of adding ammonia solution to Group 2 ions.

Mg²⁺: white precipitate of Mg(OH)₂ (insoluble in excess ammonia); Ca²⁺: no change; Sr²⁺: no change; Ba²⁺: no change.

State the results of adding excess sulfuric acid to Group 2 ions.

Mg²⁺: colourless solution (MgSO₄ is soluble); Ca²⁺: slight white precipitate (CaSO₄ is sparingly soluble); Sr²⁺: white precipitate (SrSO₄ is insoluble); Ba²⁺: white precipitate (BaSO₄ is insoluble).

Explain why magnesium hydroxide precipitates with ammonia but calcium hydroxide does not.

Mg(OH)₂ has a lower solubility product than Ca(OH)₂; the concentration of OH⁻ from ammonia is sufficient to precipitate Mg²⁺ but not Ca²⁺; Mg(OH)₂ is less soluble than Ca(OH)₂.

State the general ionic equation for the formation of metal hydroxide precipitates.

M²⁺(aq) + 2OH⁻(aq) → M(OH)₂(s) for 2+ ions; M³⁺(aq) + 3OH⁻(aq) → M(OH)₃(s) for 3+ ions.

Explain the importance of the correct sequence of tests in qualitative analysis.

Some tests interfere with others; carrying out tests in the correct order avoids false positives; ensures accurate identification of all ions present; the sequence is carbonate → sulfate → halide.

Explain why negative test results are important in qualitative analysis.

Negative results are just as important as positive results; they eliminate possibilities; they help identify the ions present by elimination; they narrow down the possible identities.

Explain why distilled water is used for preparing solutions in qualitative analysis.

Distilled water contains no ions; it does not interfere with the tests; it ensures accurate results; tap water contains ions (e.g. Cl⁻, Ca²⁺, Mg²⁺) that would interfere with tests.

Explain why test tubes are used for qualitative analysis.

Small samples are used; test tubes allow easy observation of colour changes and precipitates; test tubes are inexpensive and disposable; multiple tests can be carried out simultaneously.

Explain why a dropping pipette is used for adding reagents.

To add small amounts of reagent precisely; to control the volume added; to avoid over-addition and contamination; to observe the effect of adding dropwise.

Explain why solutions are mixed thoroughly after adding reagents.

To ensure complete reaction; to ensure all ions have reacted; to observe the full colour change; to ensure the precipitate forms completely.

Explain why precipitates are observed against a white background.

To see the colour of the precipitate more clearly; to distinguish between white and slightly coloured precipitates; to improve accuracy of observation; a white tile or white paper is used.

Explain why the test for ammonium ions uses NaOH and not another base.

NaOH is a strong soluble base; it provides OH⁻ ions readily; other soluble bases could introduce other ions that interfere; potassium hydroxide would introduce K⁺ ions but is also acceptable.

Explain why the gas in the ammonium ion test is tested with damp red litmus paper.

Ammonia is alkaline (the only common alkaline gas); damp red litmus paper turns blue; dry paper would not work as the gas needs to dissolve to form OH⁻ ions; the paper must be damp to dissolve the NH₃.

Explain why a salt must be dissolved in water before testing.

The ions must be in solution to react; solid salts may not react as the ions are fixed in the lattice; aqueous ions are free to move and react; the solution must be transparent to observe precipitates.

Explain why silver carbonate would form in the halide test if carbonates are present.

Ag⁺ reacts with CO₃²⁻ to form Ag₂CO₃; Ag₂CO₃ is insoluble and white; this would interfere with the halide test causing a false positive; removing carbonate first prevents this.

Explain why silver sulfate would form in the halide test if sulfates are present.

Ag⁺ reacts with SO₄²⁻ to form Ag₂SO₄; Ag₂SO₄ is slightly insoluble; it would form a white precipitate; this would interfere with the halide test; removing sulfate first prevents this.

Explain why barium carbonate would give a false positive in the sulfate test.

Ba²⁺ reacts with CO₃²⁻ to form BaCO₃; BaCO₃ is insoluble and white; it looks identical to BaSO₄; removing carbonate first prevents this.

Explain why the filtrate is used for the halide test after the sulfate test.

The precipitate contains the sulfate; the filtrate contains the remaining ions; testing the filtrate avoids interference from the sulfate precipitate; the precipitate would be visible and interfere with observation.

State the solubility rule for nitrate salts.

All nitrate salts are soluble in water.

State the solubility rule for chloride salts.

Most chloride salts are soluble; silver chloride (AgCl) and lead(II) chloride (PbCl₂) are insoluble.

State the solubility rule for sulfate salts.

Most sulfate salts are soluble; barium sulfate (BaSO₄), lead(II) sulfate (PbSO₄), and calcium sulfate (CaSO₄) are insoluble (CaSO₄ is sparingly soluble).

State the solubility rule for carbonate salts.

Most carbonate salts are insoluble; sodium, potassium and ammonium carbonates are soluble.

State the solubility rule for hydroxide salts.

Most hydroxide salts are insoluble; sodium, potassium and barium hydroxides are soluble.

State the practical skills required for PAG 4 qualitative analysis.

Measure and transfer precise volumes using pipettes and droppers; observe and record colour changes, precipitates, and gas formation accurately; use test tubes, racks, droppers, and water baths safely; apply knowledge of ion testing to interpret observations correctly; identify unknown ions by systematic testing.

Describe how to carry out a test for carbonate ions in the laboratory.

Add 1 cm³ of dilute hydrochloric acid to a test tube; add an equal amount of sodium carbonate solution; quickly fit a bung and delivery tube; bubble the gas into another test tube containing limewater; limewater turns milky/cloudy if CO₂ is present.

Describe how to carry out a test for sulfate ions in the laboratory.

Add the sample to a clean test tube; add a few drops of dilute hydrochloric acid (to remove carbonates); add a few drops of aqueous barium chloride solution; a white precipitate of BaSO₄ forms if sulfate is present; Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s).

Describe how to carry out a test for halide ions in the laboratory.

Add the halide solution to a clean test tube; add a few drops of nitric acid; add a few drops of silver nitrate solution; observe the colour of the precipitate; use dilute then concentrated ammonia to confirm identity; Ag⁺(aq) + X⁻(aq) → AgX(s).

Describe how to carry out a test for ammonium ions in the laboratory.

Add 10 drops of the ammonium solution to a clean test tube; add 10 drops of sodium hydroxide; swirl to mix; place in a warm water bath; hold damp red litmus paper near the mouth of the test tube (do not touch the liquid); paper turns blue if NH₃ is present.

Describe how to carry out a test for transition metal ions in the laboratory.

Add 10 drops of the metal ion solution to a clean test tube; add 10 drops of dilute NaOH; swirl gently; observe any precipitate colour; continue adding NaOH in excess; record observations; repeat with dilute ammonia solution.

State the role of nitric acid in halide tests.

Nitric acid prevents false positive results from carbonate ions precipitating with silver ions; it removes carbonate and sulfite ions which would also form precipitates with Ag⁺; it ensures the precipitate observed is only a silver halide.

State the role of hydrochloric acid in sulfate tests.

Hydrochloric acid removes carbonate ions that may interfere with the result; carbonate ions would also form a precipitate with Ba²⁺; the acid dissolves any BaCO₃ that might form.

Explain the common mistake of using sulfuric acid for acidification.

Sulfuric acid introduces sulfate ions into the test tube; this gives an immediate false-positive white precipitate of barium sulfate; always use nitric acid instead (or hydrochloric acid only for sulfate tests).

Explain the common mistake of using hydrochloric acid in the sequential anion test.

Hydrochloric acid introduces chloride ions into the mixture; this guarantees a false-positive white precipitate when silver nitrate is added; use nitric acid instead for acidification in the halide test.

Explain the common mistake of using barium chloride in the sulfate test when testing a mixture.

Barium chloride introduces chloride ions; this will give a false positive in the subsequent halide test with silver nitrate; use barium nitrate instead to avoid introducing chloride ions.

Explain why limewater is used to confirm the presence of carbon dioxide.

Limewater is calcium hydroxide solution; CO₂ reacts with Ca(OH)₂ to form CaCO₃ which is insoluble and forms a white precipitate; Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l); limewater turning milky is a positive test for CO₂.

Explain why damp red litmus paper is used to test for ammonia.

Ammonia is an alkaline gas; it dissolves in the moisture on the paper to form NH₄⁺ and OH⁻; OH⁻ ions cause the red litmus to turn blue; ammonia is the only common alkaline gas.

Explain why concentrated HCl and ammonia produce white fumes.

HCl vapour reacts with NH₃ vapour to form solid NH₄Cl (ammonium chloride); NH₃(g) + HCl(g) → NH₄Cl(s); the white solid particles form a fume/mist.

Explain why ammonia is tested with a glass rod dipped in concentrated HCl.

The HCl reacts with any NH₃ gas to form white fumes of NH₄Cl; this confirms the presence of ammonia; the test is more specific than litmus paper.

Explain why the water bath is used for warming in the ammonium ion test.

A water bath provides gentle and even heating; prevents overheating which could break the test tube or cause spitting; ensures controlled release of ammonia gas.

Explain why the test tube is not heated directly in the ammonium ion test.

Direct heating could cause vigorous boiling; the solution could spit or bump; ammonia would be evolved too rapidly; a water bath provides safer, controlled heating.

Explain why a fume cupboard is used for the ammonium ion test.

Ammonia gas is toxic and has a pungent smell; a fume cupboard removes the gas safely; protects the experimenter from inhaling the gas.

Explain why safety goggles are worn during qualitative analysis.

Protects eyes from splashes of corrosive chemicals (acids, alkalis, silver nitrate); prevents injury from glass breakage; required for all practical work.

Explain why gloves are worn when handling silver nitrate.

Silver nitrate stains skin black; the stain is difficult to remove; gloves prevent staining and protect from toxicity.

Explain why concentrated acids are handled with care.

Concentrated acids are corrosive; they can cause severe burns; they release toxic fumes; they should be handled in a fume cupboard.

Explain why the test for transition metal ions with NaOH gives coloured precipitates.

Transition metal ions have partially filled d sub-shells; when OH⁻ reacts to form the hydroxide precipitate, the colour is due to d-d electron transitions; different metal ions give different colours.

Explain why the colour of a precipitate is important in qualitative analysis.

The colour helps identify the ion present; different ions give different coloured precipitates; the colour is characteristic of the specific metal hydroxide.

Explain why the colour of silver halides is important in identification.

AgCl is white, AgBr is cream, AgI is pale yellow; these colours are characteristic and help identify the halide present; solubility in ammonia confirms the identity.

Explain why the test for carbonate ions is usually done on the solid sample.

Carbonates react with acids to produce CO₂; using the solid gives a quicker and more vigorous reaction; the effervescence is easily observed; some carbonates are insoluble in water.

Explain why the test for sulfate ions is done on an aqueous solution.

The sulfate ions must be in solution to react with Ba²⁺; the salt must dissolve to release SO₄²⁻ ions; the precipitate forms in solution and is visible.

Explain why the test for halide ions is done on an aqueous solution.

The halide ions must be in solution to react with Ag⁺; the salt must dissolve to release X⁻ ions; the precipitate forms in solution and is visible.