WHAT WE CAN LEARN FROM THE PERIODIC TABLE - Chemical Fundamentals of Geology and Environmental Geoscience - page 100

Box 6.1 Chemical symbols

a few of the one- or two-letter codes for the chemical
elements will be familiar to most readers. the majority are
abbreviations of the english element names, but a few
refer to Latin names such as Na (natrium) for sodium and
ag (argentum) for silver, or to alternative continental names
like 'wolfram' for tungsten (W). a list of chemical symbols
is given at the end of the topic (appendix C).
Using subscripts and superscripts, chemical symbols
can be augmented to specify every detail about a particu-
lar atom (see Figure 6.1.1). It is rarely necessary to spec-
ify more than one of these numbers in a given context.
they should always be written in the positions shown, to
avoid ambiguity. For example, to specify a particular iso-
tope (Chapter 10) one writes 40 ar. (Older literature may use
the obsolete notation ar 40 .)

The charge on an ionized atom
(Alternatively the oxidation state
can be shown in capital Roman
numerals.)

The mass number
A = ( Z + N ) specifies
one isotope of the element.

57
26

3+

Fe

2

The number of atoms in the
molecule under consideration The atomic number Z

Figure 6.1.1 how key parameters of an atom are codified
in an element's chemical symbol.

the atomic number Z can generally be omitted, its value
being implied by the chemical symbol itself (although see
Figure 6.2).

lithium (Li; Z = 3, electronic configuration = ls 2 2s 1 ) and
beryllium (Be; Z = 4, 1s 2 2s 2 ) it is the 2 s level; in boron
(B; Z = 5, ls 2 2s 2 2p 1 ) it is the 2p level; and so on. If we
were to disregard the increasing nuclear charge, we
would predict that the energy needed to strip an elec-
tron from this 'outermost' level would vary with
atomic number as shown in Figure 6.1a. One would
expect a general decline in ionization energy with
increasing Z , punctuated by sudden drops marking
the large energy gaps between one 'shell' and the next
one up (Figure 5.6); the downward series of steps in
Figure 6.1a thus reflects the occupation of progres-
sively higher energy levels in Figure 5.6. There is no
suggestion of periodicity.
However, because nuclear charge - and therefore the
strength of the nuclear field - increases with atomic
number, we find that each of the level steps anticipated
in Figure 6.1a is actually a ramp (Figure 6.1b), whose
rising profile reflects the increase in nuclear attraction
experienced by each electron in the atom in passing
from one atomic number to the next. The ionization
energy of helium is (nearly) twice that of hydrogen, for
example, because its doubly charged nucleus attracts
each electron twice as strongly as the singly charged
hydrogen nucleus. The ramps are separated by the
sudden drops noted in Figure 6.la, producing a mark-
edly periodic variation of ionization energy.

At the top of each ramp is an element that hangs on
tenaciously to all of its electrons. These elements are
the noble (or inert ) gases , helium (He), neon (Ne) and
argon (Ar). (Two others, krypton (Kr) and xenon (Xe),
lie beyond the Z -range of the diagram - see Figure 6.lb
inset.) Their electronic structures are characterized by
completely filled shells, in which all electrons are held
so firmly that the exchange of electrons involved in
chemical bonding is ruled out. Noble gases therefore
exhibit no significant chemical reactivity. Indeed, the
electronic structure of the noble gases is so stable that
other elements seek to emulate it by losing electrons, or
by acquiring additional electrons from other atoms (as
happens for example with the element chlorine, Cl).
Instead of forming diatomic molecules like O 2 , N 2 and
C1 2 , the noble gases are monatomic .
Immediately to the right of each noble gas in Figure 6.lb
lies an element with a conspicuously low ionization
energy. Lithium (Li), sodium (Na) and potassium (K)
are alkali metals , whose electronic structures consist of
the filled shells of the preceding noble gas, plus one
further electron which has to occupy the next shell at
a significantly higher energy (Figure 5.7). It projects
further from the nucleus than the core electrons, and is
screened by them from the full attraction of the nuclear
charge, making it even easier to remove or involve in
bonding. The chemistry of the alkali metals is dominated