Emission Spectra with a Spectroscope: A Classroom Investigation

Few classroom experiments give such direct evidence for quantum energy levels as looking at a glowing gas through a spectroscope. Instead of a smooth rainbow, you see a handful of bright, sharp coloured lines, a unique barcode for each element. This investigation lets students observe line spectra for themselves, compare them with continuous spectra, and connect what they see to the idea that electrons occupy fixed energy levels.

Purpose

• To observe and compare continuous spectra and line (emission) spectra.
• To identify elements from the pattern of their spectral lines.
• To explain line spectra using electron energy levels and photons.

Principle

When atoms are given energy (by heating or by an electric discharge), electrons move to higher energy levels. They quickly fall back, releasing the energy as photons. Because energy levels are fixed, only certain energy differences are possible, so only certain wavelengths of light are emitted. A spectroscope spreads the light out by wavelength, showing each allowed transition as a separate bright line.

A hot solid, such as a filament lamp, emits a continuous spectrum: all colours blended together. A hot, low-density gas emits a line spectrum. For the theory, see emission vs absorption spectra and photons and energy levels.

Equipment

• Hand-held spectroscopes (a diffraction grating in a tube, with a slit), or diffraction-grating glasses; one per pair of students.
• Gas discharge tubes (commonly hydrogen, helium, neon and mercury) with a dedicated power supply, operated by the teacher.
• An ordinary filament or halogen lamp (continuous spectrum).
• Optional: a compact fluorescent or older fluorescent tube; a sodium lamp if available.
• For the flame-test extension: Bunsen burner, heat-proof mat, nichrome or platinum wire loops (or wooden splints soaked in solution), dilute hydrochloric acid for cleaning loops, and small samples of salts such as sodium chloride, lithium chloride, potassium chloride, calcium chloride, strontium chloride and copper(II) chloride.
• Coloured pencils and a results sheet with blank spectrum strips marked 400 to 700 nm.

Safety

• Discharge tubes run at high voltage. Only the teacher should handle tubes and switch supplies on and off. Students observe from a marked distance and never touch the tube holder or leads. Allow tubes to cool before changing them.
• Some discharge tubes (especially mercury) emit ultraviolet light. Avoid staring at the tube at close range for long periods; brief viewing through a spectroscope at a sensible distance is normal classroom practice.
• Never look at the Sun through a spectroscope or grating.
• Flame tests: wear eye protection; tie back hair; keep flammable materials away. Copper, strontium and barium salts are harmful if swallowed, and barium salts are toxic, so use only the salts listed above, in small amounts, and wash hands afterwards. Dilute hydrochloric acid is an irritant.
• Follow your school’s risk assessment (see risk assessment in chemistry).

Method

Part A: continuous spectrum

1. Point the spectroscope at a filament lamp (not directly at a very bright bulb; view it from a metre or two away).
2. Rotate the spectroscope until the spectrum is sharp and horizontal.
3. Draw what you see on a blank spectrum strip. You should see a smooth, unbroken band from violet to red.

Part B: line spectra from discharge tubes

4. The teacher switches on the hydrogen tube. Observe it through the spectroscope from the marked position.
5. Record each bright line: its colour and approximate position on the 400 to 700 nm scale (many spectroscopes have a built-in scale).
6. Repeat for helium, neon and mercury, recording each on a separate strip.
7. Note which element gives the most lines and which gives the fewest.

Part C: everyday lights

8. Look at a fluorescent lamp through the spectroscope. Identify any lines that match your mercury spectrum.
9. If available, look at a sodium street lamp or sodium lamp. Record the strong yellow line.

Part D (extension): flame tests with a spectroscope

10. Clean a wire loop by dipping it in dilute hydrochloric acid and holding it in a roaring blue Bunsen flame until it gives no colour.
11. Dip the loop in the acid, then into a small amount of one salt.
12. Hold it at the edge of the flame and note the flame colour by eye.
13. A partner looks at the flame through a spectroscope and records any lines.
14. Clean the loop and repeat for each salt.

Results

Typical observations

Source	What is seen
Filament lamp	Continuous spectrum, violet to red
Hydrogen tube	Four visible lines: red (656 nm), blue-green (486 nm), and two violet (434 and 410 nm, the last often faint)
Helium tube	Several lines, including a strong yellow at 588 nm
Neon tube	Many lines, mostly red and orange
Mercury tube	Strong violet, blue, green (546 nm) and yellow lines
Fluorescent lamp	Continuous glow from the coating, with mercury’s bright lines on top
Sodium flame or lamp	Strong yellow line near 589 nm

Flame colours (extension)

Metal ion	Flame colour
Lithium	Crimson
Sodium	Intense yellow-orange
Potassium	Lilac
Calcium	Orange-red
Strontium	Red
Copper	Blue-green

More detail is in flame test colours.

Analysis questions

1. Describe the difference between the spectrum of the filament lamp and the spectrum of hydrogen.
2. Explain, in terms of electrons and energy levels, why hydrogen gives lines rather than a continuous band.
3. Why does each element give a different set of lines?
4. The red hydrogen line is at 656 nm. Calculate the energy of one of these photons. (Use E = hc ⁄ λ, with h = 6.63 × 10⁻³⁴ J s and c = 3.00 × 10⁸ m s⁻¹.)
5. How could astronomers use line spectra to find out what a distant star is made of?

Answer notes: (1) The lamp shows all wavelengths; hydrogen shows only a few sharp lines. (2) Electrons can only occupy fixed energy levels; each line comes from a jump between two levels, releasing a photon of exactly that energy difference. (3) Each element has different energy levels because of its different nuclear charge and number of electrons. (4) E = (6.63 × 10⁻³⁴ × 3.00 × 10⁸) ⁄ 6.56 × 10⁻⁷ ≈ 3.0 × 10⁻¹⁹ J. (5) Match dark absorption lines in the star’s spectrum to the known lines of elements (see atomic spectra and astronomy).

Sources of error and limitations

• Scale reading: the built-in scale on a cheap spectroscope may be off by 10 to 20 nm, and lines are read by eye. Compare with a known line (the mercury green at 546 nm) to calibrate.
• Faint lines: violet and deep-red lines are hard to see, especially in a bright room. Darken the room and let eyes adjust.
• Contamination in flame tests: sodium is present almost everywhere and its yellow can swamp other colours. Clean loops thoroughly, and view potassium through blue cobalt glass if available to filter out sodium’s yellow.
• Background light: stray light from windows or screens adds a faint continuous spectrum.
• Flame tests identify only some metals and can’t distinguish mixtures well. Instrumental methods are far more reliable (see atomic emission spectroscopy).

Key takeaways

• Hot solids give continuous spectra; excited gases give line spectra.
• Each line corresponds to an electron falling between two fixed energy levels and emitting a photon of one specific energy.
• Every element has a unique line pattern, which is how spectra identify elements, from school flames to distant stars.
• Discharge tubes must be handled only by the teacher because of the high voltage.
• To connect the lines to the maths, try the atomic spectra practice problems.