element, and its electron energy levels allow an element to be classified. The elec-
tron is the component of the atom used in bonding (Section 2.3). During
bonding, electrons are either donated from one atom to another, or shared; in
either case the electron is prised away from the atom. One way of ordering the
elements is therefore to determine how easy it is to remove an electron from its
atom. Chemists call the energy input required to detach the loosest electron from
atoms, the ionization energy. As explained in Box 1.1, the number of positively
charged components (protons, Z) in an atom is balanced by the same number of
negatively charged electrons that form a ‘cloud’ around the nucleus. Although
electrons do not follow precise orbits around the nucleus, they do occupy spe-
cific spatial domains called orbitals. We need only think in terms of layers of these
orbitals. Those electrons in orbitals nearest the nucleus are tightly held by elec-
trostatic attraction-forming core electrons that never take part in chemical reac-
tions. Those further away from the nucleus are less tightly held and may be used
in ‘transactions’ with other atoms. These loosely held electrons are known as
valence electrons. Electrons normally occupy spaces available in the lowest energy
orbitals such that energy dictates the electron distribution around the nucleus.
The valence electrons reside in the highest occupied energy levels and are thus
the easiest to remove. For example, the element sodium (Na) has a Znumber of
- This means that sodium has 11 electrons, 10 of which are core electrons, and
one valence electron. It is this single valence electron that dictates the way sodium
behaves in chemical reactions.
Plotting the expected first ionization energy—i.e. that required to detach the
loosest valence electron from the atom—against atomic number (Fig. 2.1a), shows
that as atomic number increases the energy required to detach valence electrons
decreases from Z=1 (H) to Z=20 (Ca). In this diagram the increasing nuclear
charge between hydrogen (H) and calcium (Ca) has been disregarded. The clear
downward steps in energy mark large energy gaps where electrons occupy pro-
gressively higher energy orbitals further away from the nucleus. The steps in Fig.
2.1a predict a marked difference in atomic structure between helium (He) and
lithium (Li), between neon (Ne) and sodium (Na) and between argon (Ar) and
potassium (K). Although much simplified, this periodic repetition of the elements
has long been used as the basis to tabulate the ordering of elements on a grid
known as the Periodic Table (Fig. 2.2), first published in its modern form by
Mendeleev in 1869.
If the ionization energy is corrected to account for nuclear charge (Fig. 2.1b)
—because increasing nuclear charge makes electron removal more difficult—the
energy pattern in each period becomes more like a ramp. Each ‘period’ begins
with an element of conspicuously low ionization energy, the so-called alkali
metals (Li, Na and K). Each of these elements readily lose their single valence
electron to form singly charged or monovalent ions (Li+, Na+and K+). The
periods of elements are depicted as ‘rows’ in the Periodic Table (Fig. 2.2), and
when these rows are stacked on top of one another a series of ‘columns’ result
(Fig. 2.2). Column Ia depicts the alkali metals. Moving up the energy ramps in
Fig. 2.1b, the alkali metals are followed by the elements beryllium (Be), magne-
sium (Mg) and calcium (Ca), each with two, relatively easily removed valence
Environmental Chemist’s Toolbox 15