Source: Michigan State University Chemistry Lab / Brown, Lemay & Bursten, The Central Science
Tags: flame test, metal ions, electron excitation, emission spectrum, ground state, excited state, cation identification, qualitative analysis, electromagnetic radiation, visible spectrum, alkali metals, alkaline earth metals
Difficulty: Introductory | Prerequisites: Basic atomic structure (protons, neutrons, electrons, energy levels)
Flame tests sit at the intersection of atomic theory and analytical chemistry. The lab builds on the idea that electrons occupy discrete energy levels around the nucleus, and that the energy gap between those levels is unique to each element. If you are comfortable with the concepts of ground state, excited state and electromagnetic radiation, you have the background you need. This is typically one of the first qualitative analysis techniques students encounter, and it connects directly to emission spectra and the Bohr model of the atom.
When you heat a metal salt in a flame, the metal's electrons absorb energy, jump to a higher level, then fall back down and release that energy as visible light. Each metal produces a characteristic colour (sodium = orange-yellow, potassium = lilac, calcium = red-orange), so the flame colour tells you which metal cation is present. The anion has no effect on the colour.
Ground state
The lowest-energy arrangement of an atom's electrons, where they normally reside when undisturbed. In simple terms, this is an atom "at rest" with its electrons sitting in their default positions.
Excited state
The condition of an atom whose electrons have absorbed energy and jumped to a higher orbital energy level. Think of it as the atom being temporarily pumped up with extra energy, which it cannot hold onto for long.
Electromagnetic radiation
Energy that travels as waves across the electromagnetic spectrum, including visible light, UV and infrared. In simple terms, it is the energy an electron releases as light when it drops back to a lower level.
Emission spectrum
The specific set of wavelengths (colours) of light emitted by an element's electrons as they return from excited states to lower energy levels. Think of it as a fingerprint made of light: each element has a unique pattern.
Flame test
A qualitative analytical technique in which a sample is introduced into a flame and the resulting colour is used to identify the metal cation present. In simple terms, you hold a chemical in a flame and the colour tells you which metal is in it.
Cation
A positively charged ion formed when an atom loses one or more electrons. Think of it as the metal side of an ionic compound, the part responsible for flame colour.
Anion
A negatively charged ion formed when an atom gains one or more electrons. In simple terms, it is the non-metal partner in an ionic compound, and it does not affect flame colour.
Negative control
A test run without the variable of interest, used to confirm that the experimental setup itself does not produce the observed effect. Think of it as a blank test: burning a plain cotton swab to make sure the swab alone does not change the flame colour.
Ionic compound
A chemical compound held together by electrostatic attraction between cations and anions. In simple terms, a substance made of a metal bonded to a non-metal through the transfer of electrons.
Every atom has discrete orbital energy levels. Electrons normally sit in the lowest available level (ground state).
When energy is added (e.g. from a flame), electrons absorb it and jump to a higher energy level (excited state).
Excited electrons are unstable and quickly fall back to ground state, releasing the absorbed energy as electromagnetic radiation.
The wavelength of that radiation depends on the size of the energy gap, which is unique to each element.
When the emitted wavelength falls in the visible spectrum (roughly 400 to 700 nm), you see it as a specific colour of light.
A sample of the compound is introduced into a hot flame, typically on a cotton swab or nichrome wire.
The heat excites the metal cation's electrons.
As those electrons return to ground state, they emit light at a wavelength characteristic of that metal.
You observe and record the flame colour, then match it to known values to identify the cation.
A negative control was run first: a plain cotton swab was burned to confirm it produced no colour change.
For liquid compounds, the swab was dipped in hydrochloric acid, then into the liquid sample, and held over the hottest part of the burner.
For solid compounds, the solid was pulverised and dissolved in a small amount of water before following the same dipping procedure.
HCl was used because it is a volatile acid that helps convert the metal into a chloride salt, which vaporises more easily in the flame.
A fresh cotton swab was used for each compound to avoid cross-contamination.
Calcium compounds (CaCO₃, CaCl₂, CaSO₄): all produced a red-orange flame, confirming Ca²⁺ as the source of the colour.
Sodium compounds (NaCl, NaPO₄, NaCO₃): all produced an orange-yellow flame, confirming Na⁺ as the source.
Potassium compounds (KC₄H₅O₆, KCl): both produced a lilac flame, confirming K⁺ as the source.
Magnesium compound (MgSO₄): no colour change. Mg²⁺ is too small and holds its electrons too tightly for a Bunsen burner flame to excite them.
Compounds sharing the same anion (e.g. Cl⁻ in NaCl, KCl, CaCl₂) produced different flame colours, so the anion is not responsible.
Compounds sharing the same cation (e.g. Na⁺ in NaCl, NaPO₄, NaCO₃) produced the same flame colour, confirming that the metal cation determines the colour.
Compound | Cation | Anion | Flame Colour |
|---|---|---|---|
CaCO₃ | Ca²⁺ | CO₃²⁻ | Red-orange |
CaCl₂ | Ca²⁺ | Cl⁻ | Red-orange |
KC₄H₅O₆ | K⁺ | C₄H₅O₆⁻ | Lilac |
MgSO₄ | Mg²⁺ | SO₄²⁻ | No colour |
CaSO₄ | Ca²⁺ | SO₄²⁻ | Red-orange |
NaCl | Na⁺ | Cl⁻ | Orange-yellow |
KCl | K⁺ | Cl⁻ | Lilac |
NaPO₄ | Na⁺ | PO₄³⁻ | Orange-yellow |
NaCO₃ | Na⁺ | CO₃²⁻ | Orange-yellow |
The key relationship: same cation = same flame colour, regardless of the anion paired with it. The energy gap between electron orbitals is an intrinsic property of each element, so two compounds containing Ca²⁺ will always emit the same wavelength of light when heated, whether the anion is a chloride, a carbonate or a sulphate.
Fireworks get their colours from metal salts: potassium compounds produce purple/lilac effects, calcium compounds produce red, sodium compounds produce yellow-orange, and copper compounds produce green or blue. Firework designers choose their metal salts on exactly the same principle tested in this lab.
Astronomers use emission spectra (the same physics, scaled up) to determine the chemical composition of distant stars and nebulae by analysing the wavelengths of light they emit.
Students often think the anion determines the flame colour. It does not. The data show that compounds with the same anion but different cations produce different colours, while compounds with the same cation but different anions produce the same colour.
Students sometimes assume that magnesium "failed" or that the test was done incorrectly because MgSO₄ showed no colour. The explanation is physical: magnesium's small atomic radius means it holds its electrons very tightly, so a Bunsen burner does not supply enough energy to excite them into the visible emission range.
A common error is to confuse "excited state" with "ionisation." In a flame test the electrons are promoted to a higher orbital within the same atom. They are not removed from the atom entirely.
Students sometimes think a non-ionic compound should produce a flame colour. Only ionic compounds with metal cations produce characteristic colours, because the test depends on the excitation of metal electrons specifically.
⚠️ Be prepared to explain why compounds with the same cation produce the same flame colour (unique energy gaps between electron orbitals for each element).
⚠️ Be ready to explain why magnesium shows no colour in a standard flame test (small atom, tightly held electrons, insufficient energy from the burner).
⚠️ Expect a question that gives you an unknown compound and asks you to predict its flame colour based on the cation present.
⚠️ Know the purpose of the negative control (plain cotton swab) and why HCl is used (volatile acid, helps vaporise the metal salt).
⚠️ Be able to distinguish between ionic and non-ionic compounds in the context of flame tests. Non-ionic compounds do not produce characteristic flame colours.
True or false: The anion in an ionic compound determines the flame colour. _____ (False)
Sodium compounds produce a __________ flame colour. (Orange-yellow)
True or false: Magnesium produces no flame colour because MgSO₄ is not an ionic compound. _____ (False, it is ionic; Mg just holds its electrons too tightly for the flame to excite them)
The purpose of burning a plain cotton swab before testing any compounds is to serve as a __________. (Negative control)
True or false: Copper(II) gluconate would produce a green flame because the cation (Cu²⁺) determines the colour. _____ (True)
Q: Why do CaCO₃, CaCl₂ and CaSO₄ all produce the same red-orange flame, even though their anions differ?
A: The flame colour is determined by the metal cation, not the anion. All three compounds contain Ca²⁺, whose electrons emit light at the same characteristic wavelength when they return to ground state.
Q: MgSO₄ showed no colour in the flame test. Does this mean it is not an ionic compound?
A: No. MgSO₄ is ionic. The lack of colour occurs because the magnesium atom is very small and holds its electrons tightly, so the energy from a Bunsen burner is insufficient to excite them into the visible emission range.
Q: You are given three white powders and told one contains Na⁺, one contains K⁺, and one is non-ionic. How would you identify each using a flame test?
A: The Na⁺ compound will produce an orange-yellow flame. The K⁺ compound will produce a lilac flame. The non-ionic compound will produce no characteristic flame colour.
Q: Potassium gluconate (KC₆H₁₁O₇) gives a lilac flame and copper(II) sulphate (CuSO₄) gives a green flame. What colour would you expect from copper(II) gluconate (CuC₆H₁₁O₇)?
A: Green. The metal cation determines the colour. Cu²⁺ produces a green flame regardless of the anion.
Q: Why is hydrochloric acid used in the procedure?
A: HCl is a volatile acid. Dipping the swab in HCl helps convert the metal compound into a metal chloride, which vaporises more readily in the flame, producing a clearer and more reliable colour.
This material connects directly to the Bohr model of the atom, which explains why electrons exist in quantised energy levels and why transitions between them produce specific wavelengths. If your course covers atomic emission spectroscopy later, the flame test is the manual, visible-light version of the same principle.
The concept of unique emission wavelengths also appears in astrophysics (identifying elements in stars by their spectral lines) and in forensic chemistry (identifying unknown residues). If you cover quantum mechanics, the energy-level transitions tested here are described more precisely by the Schrödinger equation.
Flame test, flame test colours, flame test lab report, metal ion identification, cation flame colour, electron excitation, ground state vs excited state, emission spectrum, visible spectrum, electromagnetic radiation, qualitative analysis, Bohr model, atomic emission, sodium flame test, potassium flame test, calcium flame test, magnesium no flame colour, negative control chemistry, HCl in flame test, fireworks chemistry, spectral lines, alkali metal flame colours, alkaline earth metal flame colours, ionic compound identification