Coaching Centre
Spectra and spectroscopy
Continuous, emission and absorption spectra and the conditions that produce them, how energy levels explain spectral lines, calculating photon energies, and using spectroscopy to identify elements.
Helium was discovered on the Sun 27 years before it was found on Earth. During the solar eclipse of 1868, astronomers saw a bright yellow line at 587.6 nm in the light from the Sun's edge that matched no known element. Norman Lockyer named the element after the Greek sun god, Helios. It wasn't isolated on Earth until 1895. Spectroscopy lets you identify what something is made of without ever touching it.
Three kinds of spectrum
A spectroscope spreads light into its wavelengths using a prism or a diffraction grating. In 1859–60 Gustav Kirchhoff, working with Robert Bunsen, summed up what it shows in three rules, now called Kirchhoff's laws of spectroscopy:
- A hot, dense object (a solid, a liquid, or a gas under high pressure) gives a continuous spectrum: every wavelength, with no gaps. A globe filament, molten metal and the Sun's dense interior all do this.
- A hot, low-density gas gives an emission spectrum: bright lines at particular wavelengths on a dark background. Neon signs and street lights are examples.
- A cool, low-density gas in front of a hotter continuous source gives an absorption spectrum: a continuous spectrum crossed by dark lines.
The crucial observation: for any element, the dark absorption lines are at exactly the same wavelengths as its bright emission lines. Each element has its own pattern of lines, like a barcode.
Why lines? Energy levels
Electrons in an atom can only have certain energies, called energy levels. (You'll see why in Module 8; for now, take it as an observed fact.) An electron can move between levels by emitting or absorbing exactly the energy difference as a single photon:
$$E = hf = \frac{hc}{\lambda} = E_{upper} - E_{lower}$$
- In a hot gas, collisions knock electrons up to higher levels. When they drop back, each drop emits a photon with one particular energy: an emission line.
- In a cool gas, electrons are mostly in low levels. A photon from behind with exactly the right energy can be absorbed, lifting an electron up. That wavelength is removed from the beam: an absorption line. The absorbed energy is re-emitted in all directions, so very little continues towards you.
- In a dense object, atoms are so close their levels are smeared into broad bands, so all wavelengths are emitted.
Hydrogen's brightest visible line has wavelength 656.3 nm. Find the energy of each photon in joules and electron-volts, and identify the transition, given the levels are −1.51 eV (n = 3) and −3.40 eV (n = 2).
The electron-volt (eV) is the energy an electron gains when it moves through a potential difference of 1 V: $1 \text{ eV} = 1.602 \times 10^{-19}$ J. It is a convenient unit for atoms and photons.
Identifying elements
To identify an element, compare the wavelengths of the lines in a sample with laboratory spectra of known elements. Every line of an element should be present (bright, or dark in absorption); a single matching line isn't enough. A mixture shows the lines of all its elements added together. Bunsen and Kirchhoff discovered caesium (Latin caesius, sky blue) and rubidium (rubidus, deep red) in 1860–61 this way, from unfamiliar lines in mineral water and a mineral.
Spectroscopy is used every day: testing for heavy metals in water and soil (atomic absorption spectroscopy, invented by Alan Walsh at CSIRO in Melbourne in the 1950s and now used in labs worldwide), checking the composition of steel, analysing fireworks and forensic samples, and of course working out what stars are made of (7D).
- Look at the hot filament with no gas cloud. Then put cool sodium in front. What kind of spectrum do you see now? Tap the dark lines to measure them.
- Now look at the sodium lamp on its own. Compare the bright lines with the dark ones you just measured.
- Choose the hydrogen tube and tap the red line. Check the photon energy readout (1.89 eV) with $E = hc/\lambda$.
- Put the cool hydrogen cloud in front of the filament, then switch to Look at the cloud from the side. Explain what you see using energy levels.
- Choose the mystery tube. Use the reference overlays to work out which two gases it contains, then reveal the answer.
Test yourself
Light from a hot filament passes through cool sodium vapour and then a spectroscope. The spectrum is
Cool gas in front of a continuous source absorbs its own characteristic wavelengths.
A photon has wavelength 486 nm. Its energy is
$E = hc/\lambda = 4.09 \times 10^{-19}$ J = 2.55 eV.
Why do emission and absorption lines of an element occur at the same wavelengths?
The photon energy equals the gap between two levels, whether the electron goes up or down.
Which source produces a continuous spectrum?
A hot, dense liquid emits all wavelengths.
Name each type of spectrum.
Which kind of spectrum?
Bright lines on a dark background form an spectrum. Each line is produced when an electron moves between two energy . A photon's energy is given by E = . The element first identified in the Sun's spectrum was . One electron-volt equals 1.602 × 10⁻¹⁹ .
A student views three sources through a spectroscope: an incandescent globe, a hydrogen discharge tube, and the globe viewed through a flask of cool hydrogen gas. Describe and explain each spectrum observed.
The incandescent globe gives a continuous spectrum: all colours from violet to red with no gaps. The filament is a hot, dense solid, so its atoms interact strongly and their energy levels are smeared into bands; the jostling charges emit photons of every energy.
The hydrogen tube gives an emission spectrum: a few bright lines (red at 656 nm, blue-green at 486 nm, and violet lines at 434 and 410 nm) on a dark background. The gas is hot and low density, so collisions excite electrons to higher energy levels; when they fall back, each drop emits a photon of energy $hf = E_{upper} - E_{lower}$. Only a few level differences are possible, so only a few wavelengths appear.
Through the cool hydrogen, the globe's continuous spectrum has dark lines at exactly the same wavelengths as the bright hydrogen lines. Electrons in the cool gas absorb photons whose energy exactly matches the gap between levels, so those wavelengths are removed from the beam. The energy is re-emitted in all directions, so little of it continues to the spectroscope, leaving dark lines.