Download amu (atomic mass unit): a unit used to express very small masses

Table 1 Electrical charge and relative mass of electrons, protons, and neutrons.
Particle
Symbol
Relative Charge
Mass
Proton
P
+1
1.0073 amu
Neutron
N
0
1.0087 amu
Electron
-
-1
5.468 x 10-4 amu
e
amu (atomic mass unit):
 a unit used to express very small masses
 1.66054 x 10-24 g
Rutherford’s conclusions
 Most of the mass and all of the positive charge of the atom are
contained in a small space called the nucleus.
 Most of the volume of the atom is empty space occupied by tiny
negatively charged electrons.
 There are as many negatively charged electrons outside the
nucleus as units of positive charge inside the nucleus
 the atom is electrically neutral
Rutherford Model Nuclear
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Atomic Theory
His research led him to believe that electrons in an atom exist in specific
regions at various distances from the nucleus. He visualized the electrons
as rotating in orbits around the nucleus like planets rotating around the
sun. Bohr's first paper in this field dealt with the hydrogen atom, which he
described as a single electron rotating in an orbit about a relatively heavy
nucleus. He applied the concept of energy quanta, proposed in 1900 by the
German physicist Max Planck (1858-1947) to the observed spectra of
hydrogen. Planck stated that energy is never emitted in a continuous stream,
but only in small discrete packets called "quanta" (Latin, quantus, how
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Atomic Theory
much). Bohr theorized that there are several possible orbits for electrons at
different distances from the nucleus. But an electron had to be in one specific
orbit or another; it could not exist between orbits. Bohr also stated that when
a hydrogen atom absorbed one or more quanta of energy, its electron
"jumped" to another orbit a greater distance from the nucleus. When the
electron fell back to lower orbits, it emitted quanta of energy as light, giving
rise to the spectrum of hydrogen. Each orbit, he said, is at a different energy
level. An electron in the orbit closest to the nucleus is in the first energy level; at
greater distances it may be in the second, third, or fourth energy level.
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Atomic Theory
Niels Bohr's ideas of electron distribution within the atom are useful
concepts and laid the foundation for much of the later progress in
understanding atomic structure. But, as is the case with many theories, Bohr's
assumptions have had to be modified. Difficulty arose in applying the theory to
atoms containing many electrons. Bohr's concept of the atom has been
replaced by quantum mechanics theory one of the chief difference
between these two theories is that in the quantum mechanics theory
electrons are not considered to be revolving around the nucleus in orbits, but
to occupy "orbitals"—somewhat cloudlike regions surrounding the nucleus
and corresponding to energy levels. The concept of electrons being in
specific energy levels is still retained in the modern theory.
In 1926, the Austrian physicist Erwin Schrodinger (1887-1961) introduced
his now-famous wave equation and a new method of calculation— quantum
mechanics, or wave mechanics. Schrodinger’s equation, which is a
complex mathematical expression, describes an electron as simultaneously
having properties of a wave and a particle. Thus, the electron was given
dual characteristics—some of its properties are best described in terms of
waves (like light) and other properties, like those of a particle, having
mass.
The solution of the Schrodinger equation is complex. However, the
solution leads to the introduction, of four quantum numbers which define
the probabilities of location and spatial properties of electrons in atoms.
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Atomic Theory
The four quantum numbers—n, l, m, and s. specify the energy and probable
location of each electron in an atom.
1. Electrons exist in energy levels at different distances from the nucleus. The
principal quantum number, n, indicates the energy levels of the electrons
relative to their distance from the nucleus. The number n may have any
positive integral value up to infinity (n = 1,2,3,4,... ), but only values of 1
to 7 have been established for atoms of known elements in their ground
state (lowest energy slate). Energy level n = 1 is closest to the nucleus and
is the lowest principal energy level.
2. Electrons exist in orbitals having specific shapes. The principal energy
levels (except the first) contain sublevels closely grouped together. These
sublevels consist of orbitals of specific shape. In quantum mechanics
the term orbital (not orbit) is used and refers to the region around the
nucleus in which we may expect to find a particular electron. The orbital
quantum number l (ell) specifics the shape of the electron "cloud" about
the nucleus. The four common sublevels (orbitals) normally encountered are
designated here by the lower-case italicized letters s, p, d and f
3. Electron orbitals have specific orientation in space. The magnetic
quantum number, m, designates this orientation. This quantum number
accounts for the number of s, p, d, and f orbitals that can be present in
the principal energy levels. There can be at most one s orbital, three p
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Atomic Theory
orbitals, five d orbitals; and seven f orbitals in any given principal energy
level.
4. An electron spins about its own axis in either a clockwise or
counterclockwise direction. The spin quantum number, s, relates to the
direction of spin of an electron. Because there are only two possible
directions of spin, each orbital, no matter what its designation, can
contain a maximum of two electrons.
Perspective representation of the s, px, pv and pz atomic orbitals.
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Atomic Theory
When two electrons occupy the same orbital, they must have opposite
spins. When an orbital contains two electrons, the electrons are said to be
"paired." Thus, the quantum numbers tell us relatively how far the
electrons are located from the nucleus, the shapes of the electron orbitals,
and the orientation of the orbitals in space.
The basic rules and limitations regarding the state of electrons in atoms
follow.
1. In the ground state (lowest energy state) of an atom, electrons occupy
orbitals of the lowest possible energy. Thus, an electron will occupy an s
orbital in the n = 1 level before it occupies an s orbital in the n = 2 level.
An electron will occupy an s orbital in any principal energy level before it
occupies a p orbital in that same energy level.
2. Each orbital may contain a maximum of two electrons (with opposite
spins).
3. No two electrons in an atom can have the same four quantum numbers.
This is a statement of the Pauli Exclusion Principle.
All the electrons in an atom are not located the same distance from the
nucleus. As pointed out by both the Bohr Theory and quantum mechanics,
the probability of finding the electrons is greatest at certain specified
distances, called energy levels, from the nucleus. Energy levels are also
referred to as electron shell and may contain only a limited number of
electrons. Energy levels are numbered starting with n = 1 as the shell nearest
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Atomic Theory
the nucleus and going to n = 7. They are also identified by the letters K. L. M.
N, 0, P. Q with K equivalent to the first energy level, L to the second level,
etc., as follows:
Table 2
Energy Level
1st
2nd
3rd
4th
5th
6th
7th
n
1
2
3
4
5
6
7
Letter Designation
K
L
M
N
O
P
Q
Each succeeding energy level is located farther away from the nucleus.
The maximum number of electrons that can occupy a specific energy
level can be calculated from the formula 2n2, where n is the number of the
principal energy level. For example, shell K, or energy level 1, can have a
maximum of two electrons (2 x 12 = 2); shell L, or energy level 2, can have a
maximum of eight electrons (2 x 22 = 8) etc. Table 3 shows the maximum
number of electrons that can exist in each of the first five energy levels.
Table 3 Maximum number of electrons that can occupy each principal energy level:
Energy Level n
1
2
3
4
5
Letter Designation
K
L
M
N
O
Maximum number 2n
2
2
2x1 =2
2
2x2 =4
2
2x 3 = 18
2
2x 4 = 32
2
2x 5 = 50
The theoretical value of 50 electrons in energy level 5 has never been
attained in any clement known to date.
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Atomic Theory
The principal energy levels contain sublevels, designated by the letters s, p,
d, f. The s sublevel consists of one orbital; the p sublevel consists of three
orbitals; the d sublevel consists of five orbitals; and the f sublevel consists
of seven orbitals. The maximum number of electrons that can exist in these
sublevels is 2 electrons in the s sublevel, 6 in the p sublevel, 10 in the d
sublevel, and 14 in the f sublevel. Since each orbital can contain two
electrons the 6 electrons in the p sublevel are distributed in 3 p orbitals; the
10 d electrons are determination, we need to know the maximum number of
electrons possible in an energy level and to use two of the rules:
(I) No more than two electrons can occupy one orbital, and
(2) An electron will occupy the lowest possible sublevel.
The maximum number of electrons in the first energy level is two (see
Table 3); both of these must then be s electrons. These are designated as
Is2 indicating two s electrons in the first energy level. The s orbital in the
second energy level is written as 2s, in the third energy level as 3s, etc- The
second energy level, with a maximum of eight electrons, will contain only s
and p electrons; namely, a maximum of two s and six p electrons. They are
designated as 2s22p6.
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Atomic Theory