Cheminorg

Flame Tests

In flame tests, a metal cation becomes excited when heated in a flame, causing its electrons to move to higher energy levels. When the electrons fall back to lower energy levels, they emit light at specific frequencies, which we perceive as characteristic colors. To perform a flame test, a platinum or nichrome wire is first dipped into hydrochloric acid (HCl) and then placed into the flame; this process is repeated until no color is produced, ensuring the wire is free of impurities. The Bunsen burner should be adjusted to produce a blue flame so that the color emitted by the metal cation is clearly visible. The chloride ion helps clean the wire and also forms metal chlorides, which are more volatile and dissolve more easily than many other metal salts.

Yellow / Orange

• Sodium (Na⁺): produces a very strong yellow–orange flame.
• Calcium (Ca²⁺): produces an orange-yellow flame similar to sodium, though usually less intense.
• Iron (Fe²⁺ / Fe³⁺): produces a goldish-yellow flame in both oxidation states.

Red / Pink

• Lithium (Li⁺): produces a pink to carmine-red flame.
• Strontium (Sr²⁺): produces a very bright crimson red flame.

Purple / Lilac

• Potassium (K⁺): produces a lilac or light purple flame, which is often faint and can be masked by sodium impurities.

Green

• Barium (Ba²⁺): produces a pale apple-green flame that appears briefly.
• Copper (Cu²⁺): produces a bright green flame.
• Boron: produces a vivid grass-green flame.

Blue / Blue-White

• Zinc (Zn²⁺): produces a bluish-green flame.
• Lead (Pb²⁺): produces a faint blue-white flame and releases toxic fumes.

White

• Magnesium: solid burns with an extremely bright white flame due to the formation of magnesium oxide (MgO). However, the MgCl2 doesn’t have a discernable color when performing a flame test.

Note: Sodium is extremely common and can be found on skin, laboratory surfaces, and even in trace amounts of tap water. When burned, sodium produces an intensely bright yellow flame that easily overwhelms and masks the colors of other metal ions. When testing for potassium, which produces a faint lilac flame, a cobalt blue glass square can be used. The blue glass filters out sodium’s yellow emission, allowing the violet color of potassium to become visible.

Color in Solution

• Copper(II) ions (Cu²⁺): produce a vibrant blue color when present in aqueous solution.
• Nickel(II) ions (Ni²⁺): form a greenish-blue solution.
• Iron: Iron(II), Fe²⁺, forms a pale green solution; Iron(III), Fe³⁺, produces an orange to yellow-brown solution.
• Dichromate ions (Cr₂O₇²⁻): have a highly distinctive orange color.
• Chromate ions (CrO₄²⁻): appear bright yellow.
• Permanganate ions (MnO₄⁻): produce an intense purple color, while manganese(II), Mn²⁺, forms a very pale pink solution that may appear nearly colorless at low concentrations.
• Vanadium: Vanadium(II), V²⁺, is purple; vanadium(III), V³⁺, is green; vanadium(IV), VO²⁺, is blue; and vanadium(V), typically present as VO₂⁺ or vanadate species, appears yellow.
• Cobalt(II) ions (Co²⁺): form a pink solution.
• Lead: Lead(III), Pb³⁺, produces a blue-green color, while lead(II), Pb²⁺, and lead(IV), Pb⁴⁺, are typically colorless in solution.
• The iron–thiocyanate complex, FeSCN²⁺: forms a deep reddish-brown solution.

Origin of Color in Transition Metal Ions

The colors observed in many ions are primarily due to the presence of d electrons. In transition metals, d electrons can absorb specific frequencies of visible light and undergo electronic transitions between energy levels. When a particular wavelength is absorbed, the remaining wavelengths are reflected or transmitted, causing the solution to appear the complementary color. For example, if red light is absorbed, green light is observed. For this reason, ions without d electrons are generally colorless, with only a few notable exceptions.

Common Compounds and Precipitates

Black: MnO, MnO₂, FeO, CuO, Co₃O₄, Ni₂O₃; sulfides of Ag⁺, Cu⁺, Cu²⁺, Ni²⁺, Fe²⁺, Co²⁺, Pb²⁺, Hg²⁺, Bi³⁺; Hg(s), BiI₃, Bi(s), Cu(SCN)₂, Sb(s), Hg₂O(s), Cu[C(=NH)S]₂(s)

Blue: Hydrated Cu²⁺ salts, CoHg(SCN)₄

Red: HgO, HgI₂, Pb₃O₄, Hg₂CrO₄(s), Ag₂CrO₄(s)

Green: Cr³⁺, Ni²⁺, hydrated Fe²⁺ salts, Hg₂I₂(s), Cu(C₇H₆O₂N)₂(s), CuHAsO₃(s)

Light Pink: Hydrated Mn²⁺ salts

Orange: KO₂, K₂Cr₂O₇, Sb₂S₃, ferrocyanide complexes, HgO, Sb₂S₃(s), Sb₂S₅(s)

Reddish Pink: Hydrated Co²⁺ salts

Yellow: Chromates (CrO₄²⁻), AgBr, As₂S₃, AgI, PbI₂, CdS, PbCrO₄(s), Hg₂CO₃(s), Ag₃PO₄(s), Bi(C₆H₃O₃)(s), Cu(CN)₂(s), Ag₃AsO₃(s), (NH₃)₃As(Mo₃O₁₀)₄, [SbI₆]³⁻(aq)

Brown: CdO, Fe₂O₃, PbO₂, CuCrO₄, Ag₂O(s), Ag₃AsO₄(s)

White: PbCl₂(s), Pb(OH)₂(s), PbSO₄(s), PbSO₃(s), Pb₃(PO₄)₂(s), Pb(CN)₂(s), Hg₂Cl₂(s), Hg₂HPO₄(s), Al(OH)₃(s), AgCl(s), AgCN(s), Ag₂CO₃(s), Bi(OH)₂NO₃(s), Bi(OH)₃(s), CuI(s), Cd(OH)₂(s), Cd(CN)₂(s), MgNH₄AsO₄(s), SbOCl(s), Sb₂O₃(s)

---

Flashcards: Test Your Knowledge

Click on a card to reveal the answer.