Notes on electron subshell configuration for Year 11 Chemistry students
from Chapter 16, VCE Chemistry Units 3 and 4 (Ellett)
16.4 Electron configuration
In his model of the atom, Rutherford suggested that electrons move around the nucleus in orbits, much as the planets move around the sun. While his model of atomic structure was able to explain many of the experimental observations made up until that time it did have a number of serious flaws, chief among which were
- as electrons move around the nucleus in circular paths they should emit electromagnetic radiation. This loss of energy would inevitably lead to the electrons collapsing into the nucleus
- the ‘emission spectra’ of elements could not be adequately explained by Rutherford’s model.
16.4.1 Bohr’s model of electron configuration
It was Neils Bohr, the Danish physicist who was to propose modifications to the Rutherford model which helped explain the anomalies listed above. He proposed that the electrons could move only in fixed orbits of specific energy levels. He suggested that the electrons closest to the nucleus had the lowest energies and that the further from the nucleus the electrons were found, the greater their energy. Significantly, Bohr was able to incorporate the new principles of quantum physics (as opposed to the classical theories first developed by Newton in the 17th century which Rutherford had used) that had recently been proposed by the German physicists and mathematicians Max Planck and Albert Einstein. According to this new quantum theory, the electrons were able to orbit the nucleus without loss of energy.
16.4.2 Emission and absorption spectra
A fundamental principle in Einstein’s model of quantum mechanics was the notion that all forms of electromagnetic radiation (visible light, IR, UV, X-rays etc) exist in small parcels or quanta of energy. When electrons move from one energy level to another they can only absorb or emit energy in these quanta. Bohr was able to explain emission and absorption spectra in terms of these quantum energy changes:
When electrons receive energy from an external source they will accept the exact amount of energy required to jump from a lower energy level to a higher one; we say the electrons are excited. As they drop back to their lowest energy level states (the ground state) the electrons emit the absorbed energy as the identical amount to that which they accepted in becoming excited. The amount of energy emitted or absorbed is directly proportional to the frequency of electromagnetic radiation emitted, where
E a n Þ E = hn where E = energy of emitted radiation
h = Planck’s constant
n = frequency of emitted radiation
Note: the electromagnetic radiation emitted from the movement of electrons within atoms is generally within the visible light section of the electromagnetic spectrum.
Emission spectra are generated when a sample to be analysed is heated to excite electrons within the atoms. This energy input excites the electrons, which swiftly drop back to their ground states, thus emitting the energy they had previously absorbed. By recording the frequencies of the light emitted from the sample an emission spectrum is generated, which appears as discrete lines of colour against a black background. Each coloured line represents the emission of energy as electrons fall from one particular energy level to another.
Similarly, an absorption spectrum is generated when white light is passed through a sample and the excitation of electrons results in energy being absorbed at particular frequencies. An absorption spectrum appears as a full spectrum with black lines representing where absorption has taken place.
Typical emission spectrum:
red light green light blue light
Typical absorption spectrum for the same element
red light green light blue light
16.4.3 Ionisation energies
The ionisation energy of an atom is defined as “the minimum amount of energy that must be supplied to remove an electron from an atom or ion in the gaseous state”. With its single electron, hydrogen has only one ionisation energy. In the case of helium it is possible to measure the ionisation energy of each electron; lithium has 3 electrons and so 3 successive ionisation energies etc.
The successive ionisation energies for the removal of electrons from sodium are given in Table 16.1 (below):
Table 16.1 Ionisation energies of sodium
Ionisation energy number / Process / Ionisation energy (MJ mol-1)First / Na ---> Na+ + e / 0.502
Second / Na+ ---> Na2+ + e / 4.569
Third / Na2+ ---> Na3+ + e / 6.919
Fourth / Na3+ ---> Na4+ + e / 9.550
Fifth / Na4+ ---> Na5+ + e / 13.356
Sixth / Na5+ ---> Na6+ + e / 16.618
Seventh / Na6+ ---> Na7+ + e / 20.121
Eighth / Na7+ ---> Na8+ + e / 25.497
Ninth / Na8+ ---> Na9+ + e / 28.941
Tenth / Na9+ ---> Na10+ + e / 141.373
Eleventh / Na10+ ---> Na11+ + e / 159.086
From this Table it can be noted that the first electron was removed with relative ease. The next 8 electrons all had reasonably similar ionisation energies, but the last two electrons required a very substantial amount of energy to ensure their removal.
These data, in conjunction with other information, lent support to the theory that electrons existed in shells and that there were a specified number of electrons in each shell. But why were the shells found at these particular energies, and what influenced the number of electrons at each shell? These and other questions were addressed by Erwin Schrödinger in 1926 in his mathematical treatment of Bohr’s theory in which he proposed that electrons behaved like waves as they moved around the nucleus. The details of Schrödinger’s analysis are well beyond the requirements of this course and will not be discussed here. The most important aspect of this quantum mechanical analysis of particle/wave motion is that electrons can be considered to move in regions of space called orbitals, rather than in specific orbits around the nucleus. These orbitals make up subshells which in turn can be grouped into regions of similar energy known as shells. It is to this current model of electron configuration that we shall now turn our attention.
16.4.4 The current model of electron configuration
The broadest classification of where electrons may be found is the shell. All electrons in a particular shell have similar energy levels, though it should be recognised that the shell covers a range of similar energies and is not a discrete value. The shell closest to the nucleus can hold a maximum of 2 electrons, the second shell can hold 8 electrons, the third 18 and the ‘n’th shell can hold ‘2n2’ electrons.
Within each shell are a number of similar energy levels known as subshells. There are 5 types of subshell and each has been designated a letter; the s, p, d, f and g subshells. The number of electrons which each subshell can accommodate is 2, 6, 10, 14 and 18 respectively. Furthermore, the 1st shell contains only one subshell (the s subshell); the 2nd shell contains two subshells (the s and p); the 3rd shell three subshells (s, p and d) etc. Finally, each subshell consists of regions of space where the electrons will actually be found and these regions are called orbitals. According to the Pauli Principle, an orbital can hold only a maximum of 2 electrons, so there are, by definition, one ‘s’ orbital, three ‘p’ orbitals, five ‘d’ orbitals and seven ‘f’ orbitals. All orbitals within the same subshell have the same energies.
These rules can be summarised as in Table 16.2 below:
Table 16.2 Electron configuration
Shell number / Number of subshells in shell / Description of subshells in shell / No. of orbitals in shells / No. of electrons1 / 1 / 1s / 1 / 2
2 / 2 / 2s and 2p / 1 + 3 = 4 / 8
3 / 3 / 3s, 3p and 3d / 1 + 3 + 5 = 9 / 18
4 / 4 / 4s, 4p, 4d and 4f / 1 + 3 + 5 + 7 = 16 / 32
5 / 5 / 5s, 5p, 5d, 5f and 5g / 1 + 3 + 5 + 7 + 9 = 25 / 50
Figure 16.1 - Subshell filling order
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
Before we can write the electron configuration for an atom or ion we must first remember the correct order in which electrons fill the subshells. As we have previously noted, the shells cover a range of energy levels so that by the time we reach the third and fourth shells there is some overlap in the energy levels of the constituent subshells. This is the reason why the ‘4s’ subshell is filled before the ‘3d’ subshell. These overlaps become more frequent as the number of shells increases and so the simple memory mnemonic (Figure 16.1) should be used. By reading the order of filling subshells in the direction shown it can be seen that the order is:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s etc
Example 16.2
Write out the electron configuration of the elements
(i) carbon (Z = 6)
(ii) magnesium (Z = 12)
(iii) argon (Z = 18)
(iv) iron (Z = 26)
Solution
(i) 1s22s22p2
(ii) 1s22s22p63s2
(iii) 1s22s22p63s23p6
(iv) 1s22s22p63s23p64s23d6
Note: In example (iv) the electrons filling the same shell (in this case, the third) are generally grouped together and so the more usual way to write the electron configuration of iron is 1s22s22p63s23p63d64s2
Example 16.3
Determine which of the following species are in the excited state
(i) 1s22s22p53s1
(ii) 1s22s22p63s23p64s1
(iii) 1s22s22p3
(iv) 1s22s22p63s13p5
Solution
(i) and (iv) are in excited states
Example 16.4
Write out the electron configuration of the following species
(i) 19F9-
(ii) 23Na11+
(iii) 32S162-
(iv) 64Zn302+
Solution
(i) 1s22s22p6
(ii) 1s22s22p6
(iii) 1s22s22p63s23p6
(iv) 1s22s22p63s23p63d84s2