2.2 - Electronic arrangement

2.2.1: Describe and explain the difference between a continuous spectrum and a line spectrum.

What is a spectrum?

A spectrum is produce when a light source (sunbeam, torch, laser etc) passes through a refracting prism (piece of glass, or a diffraction grating) and the light is bent through an angle that depends on the wavength of the light passing through.

If the light wavelength is long ( for instance, red light, wavelength 700nm) it is not deviated as much as a short wavelength (e.g. blue light, wavelength 400nm).

Hence, any source of light consisting of several different wavelengths may be separated and displayed on a screen or the different wavelengths may be detected electronically and displayed.

If the light source contains all possible wavelengths (e.g. white light) then a continuous spectrum results (eg a rainbow)

Continuous, Emission, and Absorption Spectra

A spectrum may be continuous, or may comprise bright lines (an emission spectrum) or dark lines (an absorption spectrum) superimposed on a background, as illustrated in the following figure.

Continuous, Emission, and Absorption Spectra

Thus, emission spectra are produced by thin gases in which the atoms do not experience many collisions (because of the low density). The emission linescorrespond to photons of discrete energies that are emitted when excited atomic states in the gas make transitions back to lower-lying levels.

A continuous spectrum results when the gas pressures are higher, so that lines are broadened by collisions between the atoms until they are smeared into a continuum. We may view a continuum spectrum as an emission spectrum in which the lines overlap with each other and can no longer be distinguished as individual emission lines.

An absorption spectrum occurs when light passes through a cold, dilute gas and atoms in the gas absorb at characteristic frequencies; since the re-emitted light is unlikely to be emitted in the same direction as the absorbed photon, this gives rise to dark lines (absence of light) in the spectrum.

2.2.2: Explain how the lines in the emission spectrum of hydrogen are related to the energy levels of electrons. Students should be able to draw an energy-level diagram, show transitions between different energy levels and recognize that the lines in a line spectrum are directly related to these differences. An understanding of convergence is expected. Series should be considered in the ultraviolet, visible and infrared regions of the spectrum. Calculations, knowledge of quantum numbers and historical references are not required.

Hydrogen Emission and Absorption Series

Electrons in their shells can receive energy in the form of heat or electricity and jump to higher energy shells (promotion). They cannot remain at these higher levels (excited state) for very long and soon fall back to their original shell (or other shells). When they fall back (relax) they have to lose the energy difference between the two shells. his loss of energy is performed by releasing electromagnetic energy in the form of infrared, visible light or ultraviolet radiation.

This movement of electrons between the shells is called electron transitions.

When electron transitions take place the energy emitted can be detected and its wavelength measured. This provides information about the relative energes of the energy shells.

In the hydrogen atom (the simplest case with only one electron to 'jump' between shells) the energy emitted appears in several series of lines each series corresponding to electrons falling back to different levels. This is shown in the diagram below.

The Lyman series corresponds to transitions between the higher shells and the lowest shell (ground state)

The energy shells are usually given a letter 'n' to describe the specific energy level. The lowest level is n=1 thesecond level is n=2 etc.

Transitions from higher shells (n >2) to n=2 produce radiation in the visible region of the spectrum and we can actually see it by splitting the light using a prism or diffraction grating and projecting it onto a screen. As in the following demonstration (right click and select rewind , then select 'play')

The hot gas in the bulb on the left emits the radiation (all mixed together) from all the transitions taking place. The prism separates the radiation according to wavelength (the shorter the wavelength the more it refracts passing through the glass) and it can then be projected onto screen or an electronic detector.

2.2.3: Describe the electron arrangement of atoms in terms of main energy levels. Students should know the maximum number of electrons that can occupy a main energy level (up to Z = 18). No knowledge of sub-levels s, p, d and f is required. The term valence electrons is used to describe the electrons in the highest main energy level.

Electronic configuration

The electrons are arranged acording to energy levels. The lowest energy (most stable) energy level is the one closest to the nucleus. The first energy level can hold up to two electrons. Nce it is full the next energy level may then start to fill up.

The second energy level can house up to eight electrons.

The outer energy level is often called the valence shell as it holds the electrons that affect bonding and chemical reactivity.

For example sodium has an electronic configuration of 2,8,1

The outer electron (valence electron) is easily lost and so it is a metal with high reactivity and a valency of 1

2.2.4: Determine the electron arrangement up to Z = 20. For example, 2.8.7 or 2,8,7 for Z = 17.

Electronic configurations (arrangements)

Resources

How line spectra are made

Useful links

Spectra

Emission and absorption spectra