Paul explains about Elemental Spectra.
When an electron in an atom is excited in various forms, it will commence vibrating in a higher, previously vacant state.
As it reverts to its previous state, it emits its excess energy as a light photon whose frequency corresponds to a line in that element’s spectrum.
Niels Bohr showed that individual atoms have specific orbitals (stationary states) in which electrons can exist without radiating electromagnetic radiation. This theory of stationary states was only applicable to hydrogen.
For elements which have more than one electron the theory becomes complex as each stationary state consists of a number of sub-levels (sub-orbitals), which can only accommodate a specific number of electrons.
The atoms of each different chemical element have different electron orbitals and sub-orbitals, so that the electron energy levels E1, E2, E3, … take on different values. Remembering that the electrons can only absorb or emit photons corresponding to the difference between two energy levels, we can see that the elements of different atoms will absorb and emit photons of different energies. (The energy of a photon is related to its colour.)
Every different chemical element has a different set of photon energies it can absorb or emit, called its spectrum. The absorption spectrum of an element shows what photons it absorbs; it consists of dark lines of absorption on a coloured background of the photons it does not absorb. The element’s emission spectrum shows what photons it emits when its electrons are given energy, and consists of a number of coloured lines on a black background. The emitted and absorbed photons both correspond to differences between the energy levels of the same atom, so the lines on one match up with the lines on the other.
Electrons in general fill the lowest energy (ground state) orbitals first. For example, the first principal orbital can only hold two electrons so the lowest orbital in helium is full, because helium atoms have two electrons.
As stated before, it is possible to move one or more electrons into higher energy orbits by exciting the atom with energy.
However, the electrons will rapidly lose this energy and return to their ground state. Excited atoms lose energy by colliding with other particles and/or by emitting electromagnetic radiation likewise seen in the element’s spectrum.
The de Broglie model of electron orbitals explained why electrons in an atom can have only certain “quantified” energies.
Any atoms of elements without a full valence shell (the outermost shell) will readily form chemical bonds with other elements until they achieve full shells.
They may do this by:
- giving or gaining electrons from other atoms of an element, forming ionic bonds. (Whether it gains or gives depends on the element’s electronegativity.)
- sharing electrons with other atoms by forming covalent bonds.
The atom’s electrons tend to form bonds in a way that increases the stability of the atom.
A pure element in the solid state will conduct electricity if it has electrons in its outermost orbital (valence shell) that are free to move from one atom to the next. This occurs when the valence shell is not full; that is, when there are less than eight electrons in the outermost orbital. This is the case for all elements on the periodic table except for the six noble gases in Group 8 (which have filled-up valence shells).
At first sight you might be tempted to think that all elements apart from the noble gases would conduct electricity. This would be true if solids were just closely packed individual atoms. Many solids, however, consist of a network of ionic or covalent bonds, so that their valence shells are filled (and thus they cannot conduct electricity).