SELECTED PRINCIPLES: INTRODUCTION - ELECTRONIC STRUCTURES OF ATOMS
The ground-state electronic structure is the lowest-energy arrangement
of the electrons in each free, gaseous, neutral atom of an element.

Each neutral atom of an element consists of a nucleus of Z protons and
(A - Z) neutrons, together with Z extra-nuclear electrons; where Z is
the atomic number and A is the mass number. The extra-nuclear electrons 
are arranged in atomic energy levels; the first four of these can hold
a maximum of 2, 8, 18, and 32 electrons, respectively. The ground-state
electronic structure of an atom is determined by a simple principle:
namely, electrons are arranged in such a way that their total energy is
a minimum. This Table shows two different methods of representing such 
electronic structures for all 36 elements in Periods 1 to 4.

Occupancies of the
main energy levels

  Occupancies of the sub-levels of
  the main energy levels   (*) (†)

Group

 ZAtom 

 1st 2nd 3rd 4th

  1st    2nd        3rd       4th

  1 

  1H 

  1

  1s1

 18 

  2He 

  2

  1s2

  1 

  3Li 

  2,  1

  1s2, 2s1

  2 

  4Be 

  2,  2

  1s2, 2s2

 13 

  5B 

  2,  3

  1s2, 2s2 2p1

 14 

  6C 

  2,  4

  1s2, 2s2 2p2

 15 

  7N 

  2,  5

  1s2, 2s2 2p3

 16 

  8O 

  2,  6

  1s2, 2s2 2p4

 17 

  9F 

  2,  7

  1s2, 2s2 2p5

 18 

 10Ne 

  2,  8

  1s2, 2s2 2p6

  1 

 11Na 

  2,  8,  1

  1s2, 2s2 2p6, 3s1

  2 

 12Mg 

  2,  8,  2

  1s2, 2s2 2p6, 3s2

 13 

 13Al 

  2,  8,  3

  1s2, 2s2 2p6, 3s2 3p1

 14 

 14Si 

  2,  8,  4

  1s2, 2s2 2p6, 3s2 3p2

 15 

 15P 

  2,  8,  5

  1s2, 2s2 2p6, 3s2 3p3

 16 

 16S 

  2,  8,  6

  1s2, 2s2 2p6, 3s2 3p4

 17 

 17Cl 

  2,  8,  7

  1s2, 2s2 2p6, 3s2 3p5

 18 

 18Ar 

  2,  8,  8

  1s2, 2s2 2p6, 3s2 3p6

  1 

 19K 

  2,  8,  8,  1

  1s2, 2s2 2p6, 3s2 3p6 3d0,  4s1

  2 

 20Ca 

  2,  8,  8,  2

  1s2, 2s2 2p6, 3s2 3p6 3d0,  4s2

  3 

 21Sc 

  2,  8,  9,  2

  1s2, 2s2 2p6, 3s2 3p6 3d1,  4s2

  4 

 22Ti 

  2,  8, 10,  2

  1s2, 2s2 2p6, 3s2 3p6 3d2,  4s2

  5 

 23V 

  2,  8, 11,  2

  1s2, 2s2 2p6, 3s2 3p6 3d3,  4s2

  6 

 24Cr 

  2,  8, 13,  1

  1s2, 2s2 2p6, 3s2 3p6 3d5,  4s1

  7 

 25Mn 

  2,  8, 13,  2

  1s2, 2s2 2p6, 3s2 3p6 3d5,  4s2

  8 

 26Fe 

  2,  8, 14,  2

  1s2, 2s2 2p6, 3s2 3p6 3d6,  4s2

  9 

 27Co 

  2,  8, 15,  2

  1s2, 2s2 2p6, 3s2 3p6 3d7,  4s2

 10 

 28Ni 

  2,  8, 16,  2

  1s2, 2s2 2p6, 3s2 3p6 3d8,  4s2

 11 

 29Cu 

  2,  8, 18,  1

  1s2, 2s2 2p6, 3s2 3p6 3d10, 4s1

 12 

 30Zn 

  2,  8, 18,  2

  1s2, 2s2 2p6, 3s2 3p6 3d10, 4s2

 13 

 31Ga 

  2,  8, 18,  3

  1s2, 2s2 2p6, 3s2 3p6 3d10, 4s2 4p1 

 14 

 32Ge 

  2,  8, 18,  4

  1s2, 2s2 2p6, 3s2 3p6 3d10, 4s2 4p2 

 15 

 33As 

  2,  8, 18,  5

  1s2, 2s2 2p6, 3s2 3p6 3d10, 4s2 4p3 

 16 

 34Se 

  2,  8, 18,  6

  1s2, 2s2 2p6, 3s2 3p6 3d10, 4s2 4p4 

 17 

 35Br 

  2,  8, 18,  7

  1s2, 2s2 2p6, 3s2 3p6 3d10, 4s2 4p5 

 18 

 36Kr 

  2,  8, 18,  8

  1s2, 2s2 2p6, 3s2 3p6 3d10, 4s2 4p6 

* The superscripts indicate the number of electrons in each sub-level;
for example, 3p5 means that there are 5 electrons in the 3p sub-level.
† Introductory courses invariably focus on the occupancies of the 
main energy levels; so, the student is well advised to view all the
entries in this column as exotic species of tasty but unripened pears
(... which maybe either plucked at some later date or allowed to rot).

The ground-state electronic structures of gaseous atoms, as summarized
by their occupancies of the main energy levels, can be represented by 
simple electron-structure diagrams (as in the examples shown below).

The Table shows the electronic structures of the first 26 elements 
placed in the so-called main groups (i.e., 1, 2, and 13 to 18). Apart 
from the single exception of helium, the elements in each group contain 
the same number of valence (or bonding) electrons in their outermost 
energy level; in part, it is this periodicity of electronic structures 
which is the basis of the (modern) Periodic Table of the Elements.

The Table also shows the electronic structures of the first 10 elements 
placed in Groups 3 to 12. Self-evidently, the occupancies of the two 
outermost energy levels show no regularity; but, regrettably, there is 
no simple explanation for the observed anomalies. Nevertheless, three
points are worth noting ... First, scandium (Group 3) behaves as if its
electronic structure was effectively 2,8,8,3; and so its chemistry has
some similarity to the elements in Group 13. Second, in contrast to the 
main group elements, those in Groups 4 to 11, known as the transition
elements, have valence electrons in the two outermost energy levels or,
more specifically, in the 3d and 4s sub-levels. Because the necessarily 
detailed descriptions of these sub-levels are (usually) reserved for an
advanced course, one's initial understanding of the chemistry of the 
transition elements will inevitably be quite limited. And third, zinc 
(Group 12) has two valence electrons in its outermost energy level; and 
so its chemistry has some similarity to the elements in Group 2.

Currently, the received wisdom is that emphasis should be given to the
correlation of the periodicity of electronic structures with chemical
reactivity; an aspect which is readily achieved by limiting studies to
elements of Groups 1, 2, 17, and 18. However, many authors have drawn
attention to a plethora of caveats; just two of these are noted here. §  
First, Cotton and Wilkinson have written: "Little of the chemistry of 
silicon can be inferred from that of carbon." A particularly important 
caveat when one reflects upon two observations: silicon is the second 
commonest element in the Earth's crust, but carbon forms more compounds 
than any other element apart from hydrogen. The following 'cluette' may
provide some perspective. A student, having compared the occupancies of 
the main energy levels, would correctly reason that 'carbon and silicon
should show similar chemistry because they both have the same number of 
valence electrons'. Contrastingly, a mature scientist, by drawing upon
advanced theoretical models, would reason that 'their chemistry should
differ because of the different characteristics of their occupied and 
unoccupied sub-levels'. Thus, the two individuals will be perceiving
electronic structures (and the Periodic Table) in very different ways.
And second, elements in a group do not always form ions with the same
charge; e.g., in Group 11, the commonest ions of copper, silver, and
gold are, respectively, Cu(II), Ag(I), and Au(III) - admittedly, these 
are transition metals: but then so are over half of all those known.

§  See, for example, F. A. Cotton and G. Wilkinson, Advanced Inorganic
Chemistry, Wiley, New York, 1988; and J. E. Huheey, Inorganic Chemistry
(Chapter 17), Harper Row, New York, 1983.

Dr. R. Peters                   Next                     Contents' List