Ion Identification

This is part of the HSC Chemistry course under the topic Analysis of Inorganic Compounds.

HSC Chemistry Syllabus

Conduct qualitative investigations – using flame tests, precipitation and complexation reactions as appropriate – to test for the presence in aqueous solution of the following

cations: barium (Ba2+), calcium (Ca2+), magnesium (Mg2+), lead(II) (Pb2+), silver ion (Ag+), copper(II) (Cu2+), iron(II) (Fe2+), iron(III) (Fe3+)
anions: chloride (Cl–), bromide (Br–), iodide (I–), hydroxide (OH–), acetate (CH3COO–), carbonate (CO32–), sulfate (SO42–), phosphate (PO43–)

Identifying Cations Using Flame Test, Precipitation and Complexation

This video explores a range of qualitative tests that can be conducted in a school laboratory to identify metal cations including:

Barium
Calcium
Magnesium
Lead
Silver
Copper(II)
Iron(II) and Iron(III)

Common Tests for Metal Cations

Precipitation Tests

Distinguishes cations based on the aqueous solubility of ionic compound formed when an anion is added.
In general, large divalent cations are more likely to produce precipitate e.g. Ba²⁺, Pb²⁺
Most effective when used to distinguish cations that produce compounds with vastly different solubility products (K_sp) with a given anion.
For example, adding SO₄^2– to distinguish between Ba²⁺ and Cu²⁺ BaSO₄(s) is a white precipitate whereas CuSO₄(aq) is a light blue aqueous solution.

Flame Tests

Flame test is a qualitative application of spectroscopy which involves the interaction between matter e.g. metals and electromagnetic radiation e.g. light.
Spectroscopy: when atomic electrons absorb heat from the flame, they are excited to a higher orbit (excited state).
- Electrons only absorb discrete amounts of energy that equal to the difference in energy levels between orbits.
- Excited electrons can return to their normal orbits (ground state) and release energy in form of electromagnetic waves. In flame tests, the electromagnetic waves have wavelengths in the visible light spectrum.

The colour of emitted light depends on its wavelength (l) and thus, the discrete amount of energy originally absorbed during electronic excitation.
- Electron orbits of metal ions have different energy levels due to their different ionic structure. This causes different amounts of energy to be absorbed when they are exposed to a Bunsen burner flame. Consequently, when excited electrons return to their normal orbits, each metal ion produces a different wavelength of visible light, and thus shows a different colour.
- Flame test is only useful for identifying certain metal ions, specifically metal ions that emit visible light.

Complexation Tests

A metal complex refers to a centrally position metal ion surrounded by molecules, bound by coordinate bonds.
- For example, Co²⁺ ion forms a pink complex with six water molecules and a blue complex with four chloride ions.

Fe³⁺ ion forms a yellow complex with six water molecules and a blood red complex with a thiocyanate ion (SCN^–).

Complexation is only useful for identifying transition metal ions as non-transition metals cannot form metal complexes.
Transition metal complexes have the unique ability to absorb visible light waves.
- Perceived colour of a solution is the colour of visible light that it does not For example, a Co²⁺ ion solution is pink because it does not absorb pink visible light.

Barium (Ba²⁺)

Precipitation test: adding sulfate ions (SO₄^2–) will produce a white precipitate (BaSO₄).
🔥Flame test – pale green; cider/apple green

Calcium (Ca²⁺)

Precipitation test: adding sulfate ions (SO₄^2–) will produce a white precipitate (CaSO₄).
🔥Flame test – orange red; colour of the sun. This is the best test to differentiate between Ba²⁺ and Ca²⁺ ions.

Pale green flame of Ba²⁺

Magnesium (Mg²⁺)

Precipitation tests for Mg²⁺ ions are not specific so therefore they should not be used to differentiate Mg²⁺ from other ions.
🔥Flame test – no colour
Mg²⁺ ions are usually best identified through a process of elimination

Orange red flame of Ca²⁺.

Lead(II) (Pb²⁺)

Precipitation tests

Adding a strong base e.g. NaOH will produce a white precipitate (Pb(OH)₂).
Adding iodide ions e.g. NaI will produce a yellow precipitate (Pbl₂)
Adding chloride ions e.g. NaCl will produce a white precipitate (PbCl₂)

☠ Solutions containing Pb²⁺ should NOT be tested with flame test as vaporised Pb²⁺ is a health hazard.

PbI₂

Identifying Transition Metals

Silver (Ag⁺)

Precipitation tests

Adding chloride ions e.g. NaCl or HCl will produce a white precipitate (AgCl)
Adding bromide ions e.g. NaBr or HBr will produce a cream precipitate (AgBr)
Adding iodide ions e.g. NaI or HI will produce a cream precipitate (AgI)

Complexation – silver chloride [AgCl(s)] dissolves in ammonia as it forms a diamminesilver ion complex. In contrast, PbCl₂ precipitate remains undissolved in ammonia. This is the best test to differentiate between Ag⁺and Pb²⁺ ions.

$$AgCl(s) + 2NH_3(aq) \rightarrow [Ag(NH_3)_2]^+(aq) + Cl^-(aq)$$

Adding hydroxide ions e.g. NaOH will produce AgOH which quickly decomposes to form brown Ag₂O. Ag₂O is soluble in acid and alkaline solutions.

$$Ag^+(aq) + OH^-(aq) \rightarrow AgOH(s)$$

$$2AgOH(s) \rightarrow Ag_2O(aq) + H_2O(l)$$

AgCl, AgBr, AgI

Brown appearance of silver oxide Ag₂O

Copper(II) (Cu²⁺)

Complexation

Cu²⁺ in aqueous solution is light blue due to the formation of copper hydrate complex (Cu(H₂O)₆²⁺). For example, copper(II) sulfate (CuSO₄).
Adding ammonia (NH₃) will produce a deep blue complex

Precipitation test

Adding a strong base e.g. NaOH will produce a blue precipitate (Cu(OH)₂). Formation of the precipitate deepends the light blue colour of aqueous Cu²⁺ ions.

🔥Flame test – green/blue

Light blue appearance of CuSO₄ (aq) solution

Blue/green flame of Cu²⁺ions.

Iron(II) (Fe²⁺)

Precipitation test: adding a strong base e.g. NaOH will produce a green precipitate Fe(OH)₂(s).

Oxidation test: adding acidified permanganate (MnO₄^–) will oxidise Fe²⁺ to Fe³⁺, causing the purple permanganate solution to decolourise.

Purple permanganate decolorises when Fe²⁺ ions (pale green) are present.

Complexation: Iron(II) hexahydrate Fe(H₂O)₆²⁺ has a pale green colour. This may be easy to identify if [Fe³⁺] is high enough. For example, iron(II) chloride (see image on right)

Precipitation of Fe(OH)₂(s)

Pale green appearance of Fe(H₂O)₆²⁺

Iron(III) (Fe³⁺)

Precipitation test – adding a strong base e.g. NaOH will produce a brown precipitate Fe(OH)₃(s).

Complexation

Iron(III) hexahydrate Fe(H₂O)₆³⁺ has a pale yellow colour. This may be easy to identify if [Fe³⁺] is high enough. For example, iron(III) chloride (see image below)
Adding thiocyanate ion (SCN^–) will produce a blood-red solution.

Precipitation of Fe(OH)₃(s)

Concentrated solution of iron(III) chloride. The yellow appearance is due to iron(III) hexahydrate complex

Iron(III) thiocyanate complex

Sequence of Tests for Identifying Cations

Pb²⁺ and Ag⁺are identified with Cl^– Ag⁺precipitates are usually white/cream in appearance. PbI₂ is distinctively bright yellow. The precipitation tests for these two ions should be conducted first within a sequence of tests.
Ca²⁺ and Ba²⁺ are identified with SO₄^2– When both ions are present, a flame test should be conducted. Ca²⁺ and Ba²⁺ions produce red and pale-green coloured flames respectively.
Cu²⁺, Fe²⁺ and Fe³⁺ are identified with OH^– Cu(OH)₂ is a blue, Fe(OH)₂ is green and Fe(OH)₃ is brown.
Mg²⁺ ions are best identified last within a sequence of test. It should not produce a colour in the flame test and precipitate with Cl^– ions in the first step. However, it does produce a white precipitate with SO₄^2– and OH^–

	Flame test	Test using NaCl(aq)	Test using Na₂SO₄(aq)	Test using NaOH(aq)
Lead Pb²⁺	X	White precipitate	White precipitate	White precipitate
Silver Ag⁺	X	White precipitate	White precipitate	Brown precipitate (Ag₂O)
Barium Ba²⁺		X	White precipitate	X
Calcium Ca²⁺		X	White precipitate	X
Magnesium Mg²⁺	X	X	White precipitate	White precipitate
Copper(II) Cu²⁺		X	X	Blue precipitate
Iron(II) Fe²⁺	X	X	X	Green precipitate
Iron(III) Fe³⁺	X	X	X	Brown precipitate

Identifying Anions Using Precipitation and pH Indicators

This video explores a range of qualitative tests that can be conducted in a school laboratory to identify anions including:

Halides: chloride, bromide and iodide
Acetate
Phosphate
Sulfate
Hydroxide
Carbonate

Common Tests for Identifying Anions

Chloride (Cl^–)

Precipitation test – adding silver ions e.g. AgNO₃ will produce a white precipitate (AgCl).

Bromide (Br^–)

Precipitation test – adding silver ions e.g. AgNO₃ will produce a cream-coloured precipitate (AgBr).

Iodide (I^–)

Precipitation tests:

Adding silver ions e.g. AgNO₃ will produce a yellowish precipitate (AgI).
Adding lead(II) ions e.g. Pb(NO₃)₂ will produce a bright yellow precipitate (PbI₂)

AgCl, AgBr, AgI

PbI₂

Hydroxide (OH^–)

Precipitation test – hydroxide ions produce soluble ionic compounds with most cations except:

Cu²⁺: blue precipitate
Fe²⁺: green precipitate
Fe³⁺: brown precipitate

Indicator test – hydroxide ions cause solution to become alkaline (pH > 7).

Red litmus paper turns blue
Universal indicator turns blue

Fe(OH)₂, Fe(OH)₃, Cu(OH)₂.

Source: Collins UK

Acetate (CH₃COO^–)

Precipitation tests are not effective for identifying acetate ions as salts containing acetate are usually quite soluble.
Indicator test – acetate is the conjugate base of acetic acid (weak acid). When dissolved in water, it exists in the following equilibrium:

Red litmus paper turns blue
Universal indicator turns blue

Adding nitric acid shifts acetate ion’s ionisation equilibrium to the right (Le Chatelier’s principle) to form more acetic acid. Production of acetic acid at high concentrations produces a distinct vinegar smell.

Red litmus paper turns blue in basic solution

Carbonate (CO₃^2–)

Adding nitric acid to produce carbon dioxide, causing bubbles to form.

HNO₃ is used instead of HCl & H₂SO₄ because the latter two introduce precipitable anions into the unknown solution which may produce a false positive result.

$$HNO_3(aq) + CO_3^{2-}(aq) \rightarrow NO_3^-(aq) + H_2O(l) + CO_2(g)$$

Identify of carbon dioxide can be verified by the limewater test (bubbling the gas in calcium oxide solution). Formation of calcium carbonate turns the solution milky.

$$Ca(OH)_2(aq) + CO_2(aq) \rightarrow CaCO_3(aq) + H_2O(l)$$

Indicator test: carbonate is the conjugate base of hydrogen carbonate (weak acid). When dissolved in water, it exists in the following equilibrium:

$$CO_3^{2-}(aq) + H_2O(l) \leftrightharpoons HCO_3^-(Aq) + OH^-(aq)$$

Red litmus paper turns blue
Universal indicator turns blue

Precipitation test: carbonate ions produce precipitates with most cations Thus, this is usually an ineffective option for identifying carbonate.

Limewater test for CO₂.

Sulfate (SO₄^2–)

Precipitation test: adding acidified barium ions (Ba²⁺) will produce a white precipitate.

‘acidified’: a few drops of dilute nitric acid are added prior to precipitation test to remove any carbonate ions that could also form a precipitate with Ba²⁺.

Precipitation of barium sulfate BaSO₄ (s)

Indicator test: sulfate is the conjugate base of H₂SO₄ (strong acid). Therefore, salts containing sulfate are neutral.

Red litmus paper remains red
Universal indicator remains green (pH = 7.0–7.4)

Phosphate (PO₄^3–)

Precipitation test: adding an alkaline solution of barium ions (Ba²⁺) will produce a white precipitate.

Phosphate ions do not precipitate with Ba²⁺ when there are H⁺ present because phosphate ions reacts with H⁺ ions to produce hydrogen phosphate. Ba²⁺ ions do not precipitate with hydrogen phosphate.

$$PO_4^{3-}(aq) + H^+(aq) \leftrightharpoons HPO_4^{2-}(aq)$$

Ba²⁺ will precipitate with phosphate ions when a weak base is added to reduce [H⁺] and shift the equilibrium towards the side with phosphate ions. An example of a weak base is ammonia NH₃.

Ammonium molybdate phosphate test

Nitric acid and ammonium molybdate (NH₄)₂MoO₄ will form a yellowcomplex after reacting with phosphate ions (ammonium phosphate molybdate)
Often used as a qualitative method to monitor environmental phosphate level.
Multi-step reaction (extension):

Indicator test: phosphate is the conjugate base of hydrogen phosphate (weak acid). When dissolved in water, it exists in the following equilibrium:

$$PO_4^{3-}(aq) + H_2O(l) \leftrightharpoons HPO_4^{2-}(aq) + OH^-(aq)$$

Thus, salts containing phosphate are basic

litmus paper turns blue
Universal indicator turns blue

Ammonium phosphate molybdate

Sequence of Tests for Identifying Anions

Carbonate ions are identified through neutralisation with a strong acid (HNO₃) to produce CO₂.
Hydroxide ions are identified by adding Cu²⁺ ions to produce a blue precipitate. This test does not work if phosphate ions are also present.
Sulfate ions are identified through addition of Ba(NO₃)₂.
Phosphate ions are identified through addition of alkaline Ba(NO₃)₂.
Chloride ions are identified through addition of AgNO₃.
If no precipitate forms, acetate ions are present.

	Test using HNO₃(aq)	Test using Ba(NO₃)₂(aq)	Test using AgNO₃(aq)
Chloride Cl^–	X	X	White precipitate
Bromide Br^–	X	X	Cream precipitate
Iodide I^–	X	X	Yellow precipitate
Sulfate SO₄^2–	X	White precipitate in acidic conditions	White precipitate
Phosphate PO₄^3–	X	White precipitate in alkaline conditions. Dissolves when nitric acid is added.	X
Carbonate CO₃^2–	Bubbling. Formation of CO₂	White precipitate	White precipitate
Hydroxide OH^–	Temperature increase due to neutralisation	X	Brown precipitate(Ag₂O)
Acetate CH₃COO^–	Smell of vinegar. Formation of acetic acid	X	White precipitate

Cross-solubility table between cations & anions outlined by the HSC syllabus

	Chloride Cl^–	Bromide Br^–	Iodide I^–	Sulfate SO₄^2–	Phosphate PO₄^3–	Carbonate CO₃^2–	Hydroxide OH^–	Acetate CH₃COO^–
Lead Pb²⁺	White	White	Yellow	White	White	White	White	Sparingly soluble (white)
Silver Ag⁺	White	Cream	Yellow	White	Yellow	Yellow	Brown	Sparingly soluble (white)
Barium Ba²⁺	✓	✓	✓	White	White	white	✓	✓
Calcium Ca²⁺	✓	✓	✓	White	White	White	Sparingly soluble (cloudy white)	✓
Magnesium Mg²⁺	✓	✓	✓	✓	White	White	White	✓
Copper(II) Cu²⁺	✓	✓	White	✓	Blue-green	White	Blue	✓
Iron(II) Fe²⁺	✓	✓	✓	✓	Brown	White	Green	✓
Iron(III) Fe³⁺	✓	✓	✓	✓	Brown	Orange-red	Brown	✓

Note: all nitrates, sodium, potassium and ammonia containing ionic compounds are aqueous.