Electronic Structure of Atoms

Electron configurations

Electron Configurations

The way in which electrons are distributed among the various orbitals is called the electron configuration

Orbitals are filled in order of increasing energy, with no more than two electrons per orbital

Lithium

This element has 3 electrons. We would thus begin by placing two electrons in the 1s ground state, or lowest energy, orbital. These two electrons would have opposite magnetic spin quantum numbers. We would then place the third electron in the next highest energy level orbital - the 2s orbital:

• The arrows indicate the value of the magnetic spin (ms) quantum number (up for +1/2 and down for -1/2)
• The description of the occupation of the orbitals would be described in the following way:

1s22s1

or, "1s two, 2s one".

• Electrons having opposite spins are said to be "paired" electrons, as with the electrons occupying the Li 1s orbital
• Likewise, the single electron in the 2s orbital (for Li) is said to be "unpaired"

Writing electronic configurations

• The two electrons in He represent the complete filling of the first electronic shell. Thus, the electrons in He are in a very stable configuration
• For Boron (5 electrons) the 5th electron must be placed in a 2p orbital because the 2s orbital is filled. Because the 2p orbitals are equal energy, it doesn't matter which 2p orbital is filled

What do we do now with the next element, Carbon (6 electrons)? Do we pair it with the single 2p electron (but with opposite spin)? Or, do we place it in another 2p orbital?

The second 2p electron in Carbon is placed in another 2p orbital, but with the same spin as the first 2p electron:

Hund's rule: for degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized

Electrons repel each other, by occupying different orbitals the electrons remain as far as possible from one another

• A carbon atom in its lowest energy (ground state) has two unpaired electrons
• Ne has filled up the n=2 shell, and has a stable electronic configuration

Electronic configurations can also be written in a short hand which references the last completed orbital shell (i.e. all orbitals with the same principle quantum number 'n' have been filled)

• The electronic configuration of Na can be written as [Ne]3s1
• The electronic configuration of Li can be written as [He]2s1

The electrons in the stable (Noble gas) configuration are termed the core electrons

The electrons in the outer shell (beyond the stable core) are called the valence electrons

Something curious

The noble gas Argon (18 electrons) marks the end of the row started by Sodium

Will the next element (K with 19 electrons) put the next electron one of the 3d orbitals?

• Chemically, we know Potassium is a lot like Lithium and Sodium

• What these elements (the alkali metals) have in common is an unpaired valence electron in an s orbital
• If Potassium has an unpaired electron in an s orbital it would mean that it is in the 4s orbital
• Thus, the 4s orbital would appear to be of lower energy than the 3d orbital(s)

The 4s orbital would be filled when we have an element with 20 electrons (Calcium).

Then we go back and fill up the 3d orbitals, which can hold a maximum of 10 electrons

Thus, the 4th row of the periodic table is 10 elements wider than the previous row - we have available five 'd' orbitals we can fill (with 10 electrons). These 10 elements are the Transition Elements, or Transition Metals.

With Cerium (element 58) the 'f' orbitals enter the picture. These orbitals can hold 14 electrons.

• The first 'f' orbitals are the 4f orbitals (n=4; l=0(s), 1(p), 2(d), 3(f))
• These additional elements are represented by the 14 lanthanide (4f orbital filling) and actinide (5f orbitals) series of elements
• The energy of the 5d and 4f orbitals are very close

1996 Michael Blaber