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.
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