Analytical chemistry

Atomic absorption and emission (AAS, ICP)

How excited atoms give off light of fixed colors, from the flame test to instruments that measure metals at trace levels.

1 Lower secondary 2 Upper secondary 3 Undergraduate 4 Advanced 5 Research

IntuitionIntuition: an atom’s color fingerprint

Put a pinch of table salt in a flame and it glows yellow; copper turns it green. Heat lifts electrons to higher energy levels, and when they drop back the atom emits light. Because the levels of each element are fixed, so are the colors.

Choose a metal ion to see the approximate color of its flame. Colors are as perceived by the eye, not measured spectra.

SchoolSchool level: colors of common metals

Flame colors (approximate)
Ion	Color
Li⁺	crimson red
Na⁺	intense yellow-orange
K⁺	lilac
Ca²⁺	orange-red
Sr²⁺	red
Ba²⁺	pale green
Cu²⁺	blue-green

UndergraduateUndergraduate: from color to measurement

\Delta E = h\nu = \frac{hc}{\lambda}

Example: The sodium D line

Sodium emits near $\lambda = 589$ nm. Find the photon energy in eV and per mole.

Solution

$E = hc/\lambda = 6.626\times10^{-34} \times 2.998\times10^{8} / 589\times10^{-9} = 3.37\times10^{-19}$ J $= 2.10$ eV, and $\times N_A$ gives about 203 kJ/mol.

In atomic emission (flame photometry, ICP) the intensity of a chosen line is proportional to the concentration of the element. In atomic absorption (AAS) a lamp of the same element sends light through the flame and the attenuation is measured, following the Beer–Lambert law. Both use a calibration with standards; ICP reaches much higher temperatures than a flame and can measure many elements at once at trace levels.