This page describes how to perform a flame test for a range of metal ions, and briefly discusses how the flame color arises. Flame tests are used to identify the presence of a relatively small number of metal ions in a compound. Not all metal ions give flame colors. For Group 1 compounds, flame tests are usually by far the easiest way of identifying which metal you have got. For other metals, there are usually other easy methods which are more reliable - but the flame test can give a useful hint as to where to look.
Figure 1: Colored flames from strontium, cesium, sodium and lithium (from left to right). Picture courtesy of the Claire Murray and Annabelle Baker from the Diamond Light Source. Confirm the colors with the elements on Table 1.
Practical Details in Carrying out Flame Tests
- Clean a platinum or nichrome (a nickel-chromium alloy) wire by dipping it into concentrated hydrochloric acid and then holding it in a hot Bunsen flame. Repeat this until the wire produces no color in the flame.
- When the wire is clean, moisten it again in the acid and then dip it into a small amount of the solid to be tested so that some sticks to the wire. Place the wire back in the flame.
- If the flame color is weak, it is often helpful to dip the wire back in the acid and put it back into the flame as if cleaning it. This should produce a very short but intense flash of color.
The table below gives a rough color guide for the elements. As people see and describe colors differently, it is impossible to provide a definitive guide.
|Sodium||strong, persistent orange|
|Cesium||blue/violet (see below)|
|Copper||blue-green (often with white flashes)|
If a red flame is produced, the flame test must be repeated with known samples of lithium, strontium, etc, comparing the colors of the unknown and the samples (compare colors in Figure 1).
The Origin of Flame Colors
Flame colors are produced from the movement of the electrons in the metal ions present in the compounds. For example, a sodium ion in an unexcited state has the electron configuration 1s22s22p6. When heated, the electrons gain energy and can be excited into any of the empty higher-energy orbitals—7s, 6p, 4d, or any other, depending on the amount of energy a particular electron happens to absorb from the flame. Because the electron is now at a higher and more energetically unstable level, it falls back down to the original level, but not necessarily in one transition.
Figure 2: (left): Na+ ion emits yellow flame when an electron gets excited and drops back to its ground state. (right): Submicroscopic view of how electrons move between different energy levels in Na+ ion
An electron excited from the 2p level to an orbital in the 7 level, for example, might relax directly back to the 2p level. This transition would release a certain amount of energy equal to the energy difference between the two orbitals, which would be seen as light of a particular color. However, it might fall back in two (or more) stages. For example, it could move first to the 5 level and then back to the 2 level.
Each of these jumps involves a specific amount of energy emitted as light energy, and each corresponds to a particular color. As a result of these transitions, a spectrum of colored lines is produced. The color detected is a combination of all the individual colors. The exact energy differences of the possible transitions vary from one metal ion to another; therefore, each unique ion has a different pattern of spectral lines, creating a different flame color.
Jim Clark (Chemguide.co.uk)