Chemical Tests

Test for OXYGEN (O₂).

• Test: insert a GLOWING wooden splint (one that has just been blown out) into a test tube of the gas.
• Observation: the splint RELIGHTS (the smouldering ember bursts into flame).
• Reason: O₂ supports combustion. The smouldering wood has fuel and heat; only O₂ is missing → adding O₂ completes the fire triangle → splint reignites.

Test for HYDROGEN (H₂).

• Test: insert a LIT wooden splint at the mouth of the test tube.
• Observation: hydrogen burns with a SQUEAKY POP sound.
• Equation: $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$ — a small explosion produces the 'pop'. Used to distinguish H₂ from any other gas — uniquely makes that sound.

Test for CARBON DIOXIDE (CO₂).

• Test: bubble the gas through LIMEWATER (a saturated solution of calcium hydroxide, Ca(OH)₂(aq)).
• Observation: limewater turns MILKY / CLOUDY / WHITE (white precipitate of calcium carbonate, CaCO₃).
• Equation: $\text{Ca(OH)}_2\text{(aq)} + \text{CO}_2\text{(g)} \rightarrow \text{CaCO}_3\text{(s)} + \text{H}_2\text{O(l)}$.
• Excess CO₂ caveat: with prolonged CO₂, the milkiness disappears as soluble calcium hydrogencarbonate forms — but this is rarely tested in 4CH1.

Test for AMMONIA (NH₃).

• Test: hold a piece of DAMP RED LITMUS PAPER at the mouth of the test tube. The word 'DAMP' is essential (gas dissolves in the water).
• Observation: red litmus turns BLUE.
• Reason: NH₃ dissolves in the water on the litmus paper to give an alkaline solution: $\text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^-$.
• Smell (with caution): NH₃ has a sharp, pungent, choking smell — never sniff directly; waft gently.

Test for CHLORINE (Cl₂).

• Test: hold a piece of DAMP LITMUS PAPER (red or blue) at the mouth of the test tube.
• Observation:
  • First, the litmus turns red (Cl₂ + H₂O → acidic solution: HCl + HOCl).
  • Then the litmus is BLEACHED (turns white) — the HOCl is a bleaching agent.
• Distinction from NH₃: NH₃ turns RED→BLUE (alkaline gas); Cl₂ turns BLUE→RED then bleaches (acidic + oxidising).

Summary table.

Gas	Test	Observation
O₂	Glowing splint	Splint relights
H₂	Lit splint	Squeaky pop
CO₂	Limewater	Turns milky
NH₃	Damp red litmus	Turns blue
Cl₂	Damp litmus	Bleaches white (after turning red)

Procedure.

1. Clean the wire: dip a platinum (or nichrome) wire into concentrated HCl, then hold in a Bunsen flame. Repeat until no colour appears in the flame (removes contaminants, especially Na⁺).
2. Dip in HCl again, then in the solid (or solution) to be tested.
3. Hold in the hot (blue) part of a Bunsen flame (air hole open).
4. Observe the flame colour. For potassium, viewing through cobalt-blue glass filters out yellow Na⁺ contamination (lets the lilac K⁺ show).

Why HCl? It helps convert metal salts to volatile metal chlorides, which give brighter, more characteristic flame colours.

Flame colours.

Cation	Colour	Mnemonic
Lithium (Li⁺)	Crimson red	'Lithium Loves Crimson'
Sodium (Na⁺)	Bright yellow / orange-yellow	Common street lights (sodium vapour)
Potassium (K⁺)	Lilac / pale purple	View through cobalt glass
Calcium (Ca²⁺)	Brick-red / orange-red	(Distinguish from Li⁺ which is more crimson)
Copper (Cu²⁺)	Green / blue-green	Like a copper-green colour

Why flame tests work. In the hot flame, electrons in the metal cation absorb energy and jump to higher energy levels. As they fall back to lower levels, they emit photons of specific wavelengths (colours). Each metal has a unique set of energy levels, so each gives a unique colour pattern. This is the basis of atomic emission spectroscopy.

Limitations.

• Only works for compounds that vaporise in the flame.
• Mixtures give blended colours that are hard to interpret.
• Sodium contamination is everywhere (sweat, dust) → cobalt glass needed for potassium.
• Some colours overlap (Li and Ca can be confused — Li is more crimson, Ca more brick-red).

Example. A white solid is added to the wire and gives a brick-red flame. Conclusion: contains Ca²⁺ (e.g. CaCl₂, CaCO₃, Ca(NO₃)₂). To find the anion, follow up with HCl test (for CO₃²⁻), BaCl₂ test (for SO₄²⁻), etc.

Adding a few drops of NaOH(aq) to a solution of an unknown metal cation gives a coloured precipitate of the metal hydroxide. The colour identifies the cation.

Standard observations.

Cation	Precipitate colour	Notes
Cu²⁺	Pale blue Cu(OH)₂	Insoluble in excess NaOH
Fe²⁺	Green Fe(OH)₂	Turns BROWN on standing in air (oxidises to Fe(OH)₃)
Fe³⁺	Red-brown / orange-brown Fe(OH)₃	Sometimes called 'rust-coloured'
Al³⁺	White Al(OH)₃	DISSOLVES in EXCESS NaOH (amphoteric) → colourless solution
Ca²⁺	White Ca(OH)₂	Slightly faint; does NOT dissolve in excess NaOH
NH₄⁺	NO precipitate	But on WARMING, gives off NH₃ gas (damp red litmus → blue)

Ionic equations.

• $\text{Cu}^{2+} + 2\text{OH}^- \rightarrow \text{Cu(OH)}_2$ (pale blue)
• $\text{Fe}^{2+} + 2\text{OH}^- \rightarrow \text{Fe(OH)}_2$ (green)
• $\text{Fe}^{3+} + 3\text{OH}^- \rightarrow \text{Fe(OH)}_3$ (brown)
• $\text{Al}^{3+} + 3\text{OH}^- \rightarrow \text{Al(OH)}_3$ (white)
• $\text{Ca}^{2+} + 2\text{OH}^- \rightarrow \text{Ca(OH)}_2$ (white)

For ammonium:

$\text{NH}_4^+ + \text{OH}^- \xrightarrow{\text{heat}} \text{NH}_3 + \text{H}_2\text{O}$

Distinguishing the two white precipitates — Al³⁺ vs Ca²⁺.

Both Al(OH)₃ and Ca(OH)₂ are white precipitates. To tell them apart:

1. Add EXCESS NaOH to each test tube.
2. Al(OH)₃ DISSOLVES in excess (it is AMPHOTERIC — reacts as an acid with the extra hydroxide):

$\text{Al(OH)}_3 + \text{OH}^- \rightarrow [\text{Al(OH)}_4]^-$

3. Ca(OH)₂ stays as a precipitate — does NOT dissolve in excess.

Distinguishing Fe²⁺ vs Fe³⁺.

• Fe²⁺ gives an INITIAL green precipitate that SLOWLY TURNS BROWN on standing (because Fe²⁺ is oxidised to Fe³⁺ by atmospheric O₂):

$4\text{Fe(OH)}_2 + \text{O}_2 + 2\text{H}_2\text{O} \rightarrow 4\text{Fe(OH)}_3$

• Fe³⁺ gives a brown precipitate immediately and stays brown.

Test for ammonium ion (NH₄⁺).

NH₄⁺ is unique — it does NOT give a precipitate with NaOH. Instead:

1. Add NaOH(aq) to the suspected NH₄⁺ solution.
2. WARM gently with a Bunsen.
3. Hold damp red litmus paper at the mouth of the tube.
4. If NH₃ is released, the litmus turns BLUE → confirms NH₄⁺.

This is the only common cation that gives a gas with NaOH rather than a precipitate, so it is diagnostic.

Test 1: Carbonate (CO₃²⁻).

• Test: add a few drops of dilute HCl (or any dilute acid).
• Observation: fizzing / effervescence; the gas turns limewater milky → CO₂.
• Ionic equation: $\text{CO}_3^{2-} + 2\text{H}^+ \rightarrow \text{H}_2\text{O} + \text{CO}_2$.

Test 2: Halide ions (Cl⁻, Br⁻, I⁻).

• Procedure: add a few drops of dilute nitric acid (HNO₃) FIRST, then add silver nitrate solution (AgNO₃(aq)).
• Why dilute HNO₃ first? Removes any carbonate (which would give white Ag₂CO₃ precipitate — false positive) or hydroxide (Ag₂O — brown precipitate). Cannot use HCl because that would itself precipitate with AgNO₃!
• Observations:

Halide	Precipitate colour	Solubility in NH₃
Cl⁻	WHITE AgCl	Dissolves in dilute NH₃
Br⁻	CREAM AgBr	Dissolves only in concentrated NH₃
I⁻	YELLOW AgI	Insoluble in NH₃

Ionic equations:

• $\text{Ag}^+\text{(aq)} + \text{Cl}^-\text{(aq)} \rightarrow \text{AgCl(s)}$
• $\text{Ag}^+\text{(aq)} + \text{Br}^-\text{(aq)} \rightarrow \text{AgBr(s)}$
• $\text{Ag}^+\text{(aq)} + \text{I}^-\text{(aq)} \rightarrow \text{AgI(s)}$

Test 3: Sulfate (SO₄²⁻).

• Procedure: add a few drops of dilute HCl FIRST, then add barium chloride solution (BaCl₂(aq)).
• Why dilute HCl first? Removes any carbonate (which would give white BaCO₃ — false positive). Don't use HNO₃ here as nitrate doesn't interfere but HCl is the standard.
• Observation: WHITE precipitate of barium sulfate (BaSO₄).
• Ionic equation: $\text{Ba}^{2+}\text{(aq)} + \text{SO}_4^{2-}\text{(aq)} \rightarrow \text{BaSO}_4\text{(s)}$.

Summary table.

Anion	Test reagents	Observation
CO₃²⁻	dilute HCl	Fizz; CO₂ (limewater milky)
Cl⁻	dilute HNO₃ + AgNO₃	White precipitate (AgCl)
Br⁻	dilute HNO₃ + AgNO₃	Cream precipitate (AgBr)
I⁻	dilute HNO₃ + AgNO₃	Yellow precipitate (AgI)
SO₄²⁻	dilute HCl + BaCl₂	White precipitate (BaSO₄)

Why the 'acid first' steps matter. Both the halide test and the sulfate test rely on a unique precipitate. If carbonate or hydroxide is present in the solution, those would ALSO give a white precipitate with Ag⁺ or Ba²⁺ — leading to false positives. The dilute acid (HNO₃ for halide, HCl for sulfate) destroys those interferents:

• HCl removes CO₃²⁻ as CO₂.
• HNO₃ removes CO₃²⁻ as CO₂ AND prevents AgOH from forming.

Always state the 'acid first' step — it earns its own mark in the mark scheme.

Two distinct things can be tested for: PRESENCE of water (chemical tests) and PURITY of water (physical test).

Chemical test 1: Anhydrous copper(II) sulfate.

• Procedure: add a few drops of the suspected water sample to white anhydrous CuSO₄ powder (or vice versa).
• Observation: white powder turns BLUE (forms hydrated copper(II) sulfate pentahydrate).
• Equation: $\text{CuSO}_4 + 5\text{H}_2\text{O} \rightarrow \text{CuSO}_4 \cdot 5\text{H}_2\text{O}$.
• The reaction is also exothermic (the powder warms up).

Chemical test 2: Cobalt(II) chloride paper.

• Procedure: dip a strip of blue anhydrous cobalt(II) chloride paper in the liquid.
• Observation: paper turns PINK (forms hydrated cobalt chloride hexahydrate).
• Equation: $\text{CoCl}_2 + 6\text{H}_2\text{O} \rightarrow \text{CoCl}_2 \cdot 6\text{H}_2\text{O}$.
• The reverse change (heat to dry) returns the paper to blue.

Important caveat — presence vs purity.

Both chemical tests confirm that water is PRESENT in the liquid sample. They do NOT prove the liquid is PURE water. A salt solution (NaCl in water), an ethanol-water mixture, or even pond water would also turn the CuSO₄ blue / cobalt paper pink.

Physical test for pure water — boiling point and melting point.

To prove that a sample is pure water (and not a mixture), measure:

• Melting point: pure water melts at exactly 0 °C at standard pressure (101 kPa).
• Boiling point: pure water boils at exactly 100 °C.

Why? Pure substances have SHARP, FIXED melting and boiling points. Mixtures (impure water) do NOT — they boil over a range, AND the boiling point is elevated (and freezing point depressed) by the dissolved impurities.

• Pure water: boils at exactly 100 °C; a thermometer in boiling water reads steadily at 100 °C.
• Salt water (1 M NaCl): boils at ~102 °C (depending on conc); thermometer reading drifts up as water evaporates and concentration increases.

Decision flowchart.

• Want to test 'is there ANY water in this liquid?' → use chemical test (CuSO₄ or cobalt paper).
• Want to test 'is this liquid PURE water, with NO dissolved substances?' → use physical test (mp = 0 °C, bp = 100 °C, sharp not range).

Bonus — anhydrous vs hydrated salt formulae.

• Hydrated CuSO₄·5H₂O — BLUE crystals.
• Anhydrous CuSO₄ — WHITE powder (obtained by heating hydrated CuSO₄ to drive off water of crystallisation).
• The reaction is reversible: adding water to anhydrous CuSO₄ regenerates the blue hydrate.

So 'CuSO₄ blue → white on heating' and 'CuSO₄ white → blue on adding water' are CONNECTED chemistry — these are also the basis of the test for water itself.