Geology Reference
In-Depth Information
environments as Sn(II) and Sn(IV) compounds. Another
example is arsenic which, as Figure 4.2 shows, occurs
naturally as both As(III) and As(V).
making this downward transition must dispose of an
amount of energy Δ
E
(Figure 6.4) exactly equal to the
difference between its initial and final energy levels,
and this energy output takes the form of electromag-
netic radiation (Box 6.3.). Excited atoms therefore emit
a series of sharply defined wavelength peaks
(Figure 6.5) that provide detailed information about
their electronic energy structure: these wavelengths
constitute the
emission spectrum
of the element(s)
concerned. Because the electron energy levels in an
atom, and therefore the wavelengths it emits, are
Z
-dependent, the spectrum of one element is readily
distinguishable from that of another (Box 6.3.). Atomic
spectra thus provide an important practical means of
identifying elements present in a sample - be it a rock
powder or a solution - when suitably excited, and of
determining their relative abundances.
The success in explaining why an atomic emission
spectrum consists of a series of sharp lines rather than a
continuum is one of the triumphs of wave mechanics.
Just as an atom
emits
radiation at characteristic
wavelengths when electrons fall from excited energy
levels to lower ones, so the excitation of an electron
from the ground state to an excited energy level may
be associated with (
i.e.
can be caused by)
1
the
absorption
of radiation at the same distinctive wavelengths. The
astronomical application of such atomic
absorption
spectra
will be discussed in Chapter 11.
Can an electron leap from any energy level to
any other level within an atom? Analysis of the peaks
present in an atomic spectrum indicates that the answer
must be 'no'. Certain transitions are 'forbidden' because
they would violate basic physical principles such as the
conservation of angular momentum. Wave mechanics
recognizes such restrictions in the form of a number of
'selection rules'. For example, a radiative transition in
an atom must satisfy the two conditions:
Atomic spectra
An atom is said to be in its
ground state
when all of its
electrons occupy the lowest energy levels allowed to
them by the Pauli Principle (Chapter 5). This lowest-
energy configuration (Figure 6.4a) is the one normally
encountered at room temperature. But atoms can
absorb energy from their surroundings, for example
when they are heated or exposed to energetic rad-
iation, and this causes one or more electrons to jump
from a stable, low energy level into one of the vacant
orbitals at higher energy, or perhaps even to be ejected
from the atom altogether (Figure 6.4a). The unsta-
ble
excited state
of the atom so produced, with a
vacancy in a low energy level (Figure 6.4b), soon
reverts to the stable ground state by filling the vacancy
with an electron from a higher level. The electron
(a)
(b)
Magnesium atom
in ground state
Magnesium atom
in excited state
0
3p
3p
3s
3s
2p
2p
2s
2s
Energy
loss
Δ
E
∆
l
=±1
∆
n
≠ 0
Vacancy
1s
1s
Thus element spectra do not include lines that corre-
spond to transitions between 3 s and 2 s states (for
Empty orbital
Half-occupied orbital
Full occupied orbital
1
But excitation may be caused by forms of energy other than
electromagnetic radiation. For example, in the electron
microprobe (Box 6.4) a high-energy electron beam is the
agent of excitation.
Figure 6.4
Ground states and excited states. The transition
shown emits an MgK
β
quantum as an X-ray photon.
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