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ATOMIC STRUCTURE
Historical Background.
The ancient Greeks had two theories concerning the nature of matter, both based on
philosophical beliefs rather than on scientific studies. According to the continuous theory,
a sample of matter could be divided and subdivided into smaller and smaller parts
indefinitely. Each part, no matter how small, would retain the characteristics of the
original sample. According to the discontinuous theory, all matter is made up of tiny
particles, or atoms, which could not be broken down into smaller particles.
The first modern atomic theory was developed in the early 1800s by the English chemist
John Dalton. Dalton’s theory was based on laws of chemistry established by scientific
experiments and observations. In Dalton’s theory, as in the discontinuous theory of the
ancient Greeks, the atom was considered to be a solid, indivisible particle.
With the discovery of electrons in the late 1890s, the accepted model of the atom as a
solid particle began to change. It is now known that atoms are made up of many different
kinds of smaller particles.
Subatomic Particles.
The three major kinds of particles in an atom are electrons, protons, and neutrons. .
The electron is the fundamental unit of negative charge in an atom. The mass of
an electron, which equals 1/1836 the mass of a proton, is negligibly small.
The proton is the fundamental unit of positive charge in an atom. Its mass is
approximately 1 atomic mass unit (amu).
The neutron has zero charge. Its mass is also approximately 1 amu, but is slightly
greater than the mass of a proton. It is sometimes helpful to think of a neutron as a
particle formed by the combination of a proton and an electron.
Characteristics of the three major subatomic particles are given in Table 1.
Table 1. Subatomic Particles
_____________________________________________________________________________
Symbol
_________________________________________
Particle
In General Use
In Nuclear Equations Mass
Charge
o
Electron
e
1/1836 amu -1
-1e
1
Proton
p
1 amu
+1
1H
1
Neutron
n
n
1
amu
0
0
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Because protons and neutrons are found in the nucleus of atoms, they are called nucleons.
These are the only particles that have been found in a stable nucleus. However, many
other particles, such as positrons and mesons, have been observed during experiments
involving nuclear disintegration.
Structure of Atoms
In 1911, Ernest Rutherford showed that the positive charge and mass of an atom are
concentrated in the very small, dense nucleus, and that most of .the atom is empty space.
Rutherford developed this model of the atom after conducting experiments in which he
bombarded very thin gold foil with a beam of positively charged particles called alpha
particles. Most of the alpha particles passed straight through the foil without being
deflected, indicating that there is mostly empty space in atoms. A few of the alpha
particles were deflected at angles of more than 90°, These deflections occurred when the
positively charged alpha particles came near the positively charged nuclei and were
repelled by the like electrical charges.
Differences Between Atoms
A major principle of the atomic theory states that atoms of different elements are
different. They have different chemical and physical properties, which can be attributed
to differences in the number of protons and neutrons, and in the number and arrangement
of electrons.
Atomic Numbers.
1.
2.
3.
1.
2.
3.
The atomic number of an atom is equal to the number of protons in its nucleus. Since an
atom has no net charge, the atomic number also equals the number of electrons. The
number of protons in the nucleus distinguishes one element from another and identifies a
particular element.
Mass Numbers and Isotopes.
The mass number of an atom is equal to the number of protons plus the neutrons in its
nucleus. Every atom of a particular element has the same number of protons,
but the number of neutrons may differ. These different forms of an element having the
same number of protons but a different number of neutrons are called isotopes. Thus,
while all atoms of an element have the same atomic number, their mass numbers can
vary.
The mass number of an atom can be written as a number following the name of the
element, for example, carbon-12. The atomic number and mass number of an atom are
also shown by a sub- script and superscript preceding the chemical symbol for the
element, as in 126C and 146C.
In 126C, note that:
The mass number, or the number of nucleons (protons + neutrons), is 12.
The atomic number, or the number of protons (equal to the number of electrons), is 6.
The number of neutrons (mass number - atomic number) is 12 -6 = 6.
In 146C, note that:
The mass number is 14.
2. The atomic number is 6.
The number of neutrons is 8.
The number of neutrons in an atom can be found by subtracting the atomic number
from the mass number.
Atomic Mass. The mass of a neutral atom is measured in atomic mass units (amu),
which are based on the mass of the carbon-12 atom. By definition, the mass of the
carbon-12 atom is 12.000 amu, and I amu equals n the mass of a carbon- 12 atom.
Isotopes and Average Atomic Mass. Most elements occur in nature as mixtures of
isotopes. In general, the atomic mass for an element given in tables is an average atomic
mass that is the average of the masses of the naturally occurring isotopes of the element.
For example, most chlorine exists as two isotopes, chlorine-35 and chlorine- 37. About
75 percent of naturally occurring chlorine is chlorine-35. Most other chlorine atoms,
about 25 percent, are chlorine-37. The weighted average atomic mass given for chlorine
in reference tables is 35.453 amu, or 35.5 amu.
Gram Atomic Mass. A sample of an element with amass in grams numerically equal
to the atomic mass is a gram atomic mass. For example, 12 grams of carbon, 16 grams of
oxygen, and 32 grams of sulfur are equal to I gram atomic mass of each element. One
gram atomic mass of an element contains 6.02 x 1023 atoms. This is Avogadro's number.
The Bohr Model of the Atom.
Niels Bohr developed a model of the atom in which electrons are described as revolving
around the nucleus in concentric circular orbits, or shells (see Figure 1). The shells, at
increasing distances from the nucleus, are designated by the letters K, L, M, N, 0, P, and
Q, or by the numbers I through 7. The energy of the electrons in the shells increases with
distance from the nucleus; that is, the electrons nearest the nucleus have the lowest
energy, while those most distant have the highest energy.
4
N
Figure 1. The Bohr atom
Quanta and Spectral Lines.
When all the electrons of an atom are at their lowest energy levels, the atom is said to be
in the ground state. When an electron absorbs energy and moves to a higher energy level,
the atom is said to be in an excited state, which is unstable. When the electron falls from
a higher energy level to a lower energy level, it gives off energy.
Electrons absorb and give off energy only in discrete amounts called quanta. An electron
can move from one energy level to a higher energy level only by absorbing a quantum of
energy equal to the difference in energy between the two levels. When an electron falls
from a higher to a lower energy level, it gives off a quantum of energy equal to the
difference in energy between the two levels.
When the electrons of an atom in the excited state return to the ground state, they give off
energy in the form of radiant energy of specific frequencies, or colors. By viewing
radiation from various sources with a spectroscope, scientists have found that each
element produces a characteristic set of spectral lines by which it can be identified.
Orbital Model of the Atom.
Much of the Bohr model of the atom was based on analysis of the emission of energy by
hydrogen atoms. Careful analysis of energy emission by atoms of other elements revealed flaws
in the Bohr model and resulted in the development of the orbital model of the atom. In the orbital
model, the locations of electrons are described in terms of the average regions of most probable
electron location. These regions, called orbitals, differ in size, shape, and orientation in space
from the circular orbits of the Bohr atom.
Table 2. Subshells and Orbitals
Energy
Sublevels
orbitals
7s
O
6p
OOO
5d
OOOOO
4f
6s
OOOOOOO
O
5p
OOO
4d
5s
OOOOO
O
4p
OOO
3d
4s
OOOOO
O
3p
3s
OOO
O
2p
OOO
2s
O
1s
O
Arrangement of Electrons. In the orbital model, the positions of electrons are
represented by quantum numbers. The electron shells pro- posed by Bohr became the
principal energy levels of the orbital model. Each principal energy level is represented by
a principal quantum number (n), which is the same as the shell designation in the Bohr
model. For any atom in the ground state, the number of energy levels that have at least
one electron is equal to the number of the period of the element in the periodic table.
Sublevels. The principal energy levels of the orbital model are divided into sublevels.
The number of possible sublevels in each principal energy level is equal to the principal
quantum number for that energy level. Thus, the first principal energy level has one
sublevel; the second principal energy level has two sublevels, and so on. The sublevels
are designated s, p, d, and I. The s sublevel has the lowest energy in that energy level,
while the I sublevel has the highest. In an atom in the ground state, no principal energy
level has more than four occupied sublevels.
Sublevels contain one or more orbitals (see Table 2). Each orbital has a different spatial
orientation and can accommodate one or two electrons. Electrons behave like tiny
magnets spinning on their axes. When two electrons occupy the same orbital, they have
opposite spin.
Table 2-3 shows the distribution of electrons in principal energy levels I through 4.
Electron Configuration. The electron configuration of an atom describes the
arrangement of electrons in principal energy levels, sublevels, and orbitals. Beginning
with the hydrogen atom, and in order of atomic number, the electron con- figurations of
the atoms of each element can be built up by adding one electron at a time according to
the following rules.
1. Any orbital can accommodate one or two electrons (if two, the electrons have opposite
spins).
2. An added electron is placed in the unfilled orbital of the lowest energy.
3. Within a given sublevel, a second electron is not added to an orbital until each orbital
in the sublevel contains one electron.
4. Only the s sublevel or the s and p sublevels (four orbitals for a maximum of eight
electrons) can be occupied in the outermost energy level of any atom.
Table 3. Electron Arrangement
1
one
s
one
--
2
2
2
two
s, p
four
(1 + 3)
8
(2 + 6)
3
three
s, p, d
nine
(1 + 3 + 5)
18
(2 + 6 + 10)
4
four
s, p, d, f
sixteen (1 + 3 + 5 + 7)
32
(2 + 6 + 10 + 14)
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Table 2-3 also shows the order in which electron levels fill. Beginning with the lowest
energy, 1s to 2s to 2p to 3s to 3p to 4s to 3d, which has lower energy than 4p. Again, 4d
has lower energy than 5p. Then 5p to 6s.
The electron configuration of an atom can be written out as shown below. The principal
energy levels, sublevels, and numbers of electrons are each specified. For example, the
electron con- figuration of magnesium (12Mg) is written as:
A simplified electron configuration can be written as 2-8-2, which indicates the number
of electrons in the principal energy levels, K-L-M.
Sometimes, for atoms with many electrons, it is convenient to represent the arrangement
of inner electrons by referring to the last preceding noble gas in the periodic table. For
example, the electron configuration for 56Ba is:
[1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2] 6s2
xenon core
Thus, the simplified electron configuration for
56Ba may be written as:
[Xe core] 6s2
The term “Xe core” refers to the configuration of 54 electrons as found in an atom of the
noble gas xenon, 54Xe.
Valence Electrons.
The electrons found in the outermost principal energy level of an atom are the valence
electrons. The remainder of the atom-its nucleus and other electrons-is the core, or kernel,
of the atom. Most of the chemical properties of an atom are related to the valence
electrons.
Electron-dot symbols are used to represent valence electrons. The electron configuration
for an atom gives the correct number of valence electrons, generally not more than eight.
Examples of electron-dot diagrams are shown in
Figure 2-2.
Note that electrons are often paired to reflect their arrangement within the orbitals of the
8 and p sublevels. In these diagrams, the chemical symbol for an atom represents the
kernel.
Ionization Energy.
The amount of energy required to remove the most loosely bound electron from an atom
in the gas phase is called the ionization energy (IE).
M(g) + IE  M+(g) + e-The second ionization energy refers to the amount of energy needed to remove the second
most loosely bound electron. Removal of the outermost electron of an atom requires the
least amount of energy because the attractive force of the positively charged nucleus is
least. Removal of the electron nearest the nucleus requires the greatest amount of energy
because the attractive force of the nucleus is greatest. Thus, beginning with the outermost
electron, each successive ionization energy is greater than the previous one. Ionization
energies are used to compare certain chemical properties of elements, such as chemical
reactivity.