Chapter 10
Modern Atomic Theory
and the Periodic Table
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10.1 A brief history
atoms proposed by Greek philosopher
Dalton’s model of atom
Thomson’s model
Rutherford’s model
there remain questions can not be answered:
․ how atomic structure relates to the periodic
table (arrangement of electrons)
․ how to explain line spectrum of atom
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10.2 Electromagnetic radiation
energy and light – energy travels through space is
by electromagnetic radiation,
all radiations travel at the same speed –
v = λ • υ = 3 × 10-8 m/s
λ : wavelength
υ : frequency
electromagnetic spectrum
wavelike nature -- radiation
behaves like particle -- photon
explain the properties of electromagnetic
radiation by both wave and particle properties
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10.3 The Bohr atom
at high temperature or when subjected to high
voltage, elements in the gaseous state give off
colored light
a set of brightly colored lines – line spectrum
ex. line spectrum of hydrogen
line spectrum indicates that light is being emitted
only at certain wavelength (or frequency)
1912 Bohr model of hydrogen
• electrons exit in specific regions at various
distance from the nucleus
• the electrons as revolving in orbits around the
nucleus like planets rotating around the Sun
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Max Planck energy quanta
the energy is never emitted in a continuous
stream but only in small discrete packets called
quanta
electrons are only in several energy levels
hydrogen atom absorbs one or more quanta of
energy, the electron will jump to a higher energy
level
ground state
– the lowest energy level
excited states
– the higher energy levels
when an electron falls from a high-energy level
to a lower one, a quantum of energy is emitted
as light at a specific frequency
Bohr model
1) suggesting quantized energy levels for electrons
2) showing that spectral lines result from the
radiation of small increments of energy when
electrons shift from one energy level to another
however, Bohr model only succeeded in H atom,
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did not succeed for heavier atoms
1924 de Broglie all objects have wave properties
for small objects such as an electron, the wave
properties become significant
1926 Schrödinger Schrödinger equation
a mathematical model that described electrons
as waves
the probability of finding an electron in a certain
region around the atom can be determined
wave mechanics or quantum mechanics
forming the basis for our modern understanding
of atomic structure
we cannot locate an electron precisely within an
atom
electrons are not revolving around nucleus in
orbits as Bohr postulated
orbital a region where a high probability of
finding a given electron
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10.4 Energy levels of electron
Bohr the energy of the electron is quantized
the electron is restricted to only certain
allowed energies
wave-mechanics also predicts discrete principal
energy levels within the atom
these energy levels are
designated by the letter n
n: positive integer
as n increases, the energy of the electron
increases, and the electron is found on average
farther from the nucleus
each principal energy level
is divided into sublevels
n=1
n=2
n=3
1 sublevel
2 sublevels
3 sublevels
each of these sublevels contains space for
electrons called orbitals
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n=1
1s orbital spherical shape
the electron does not move around on the
surface of sphere, the surface encloses a space
where there is a 90% probability that the
electron may be found
how many electrons can fit into a 1s orbital?
spin a property of electron
each electron can only spin in two directions
representation of the spin ↑ or ↓
two electrons with the same spin cannot occupy
the same orbital
Pauli exclusion principle – an atomic orbital
can hold a maximum of two electrons which
must have opposite spins
n = 1 energy level contains one type of orbital
(1s) that hold a maximum of 2 electrons
n = 2 2s spherical shape
hold a maximum of two electrons
2p – 2px, 2py, 2pz
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each p orbital has two lobes and can hold a
maximum of two electrons
the total number of electrons that can reside in
all three p orbitals is 6
n = 2 energy level contains two types of
orbitals (a 2s and three 2p) that hold a
maximum of 8 electrons
n = 3 3s
3p – 3px, 3py, 3pz
3d – 3dxz, 3dxy, 3dyz, 3dz2, 3dx2 – y2
n = 3 energy level contains three types of
orbitals (a 3s, three 3p, five 3d) that hold a
maximum of 18 electrons
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n = 4 energy level contains four types of
orbitals (a 4s, three 4p, five 4d , seven 4f)
that hold a maximum of 328 electrons
hydrogen atom consists of a nucleus (one proton)
and one electron occupying a region outside of the
nucleus
10-13 cm
10-8 cm
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10.5 Atomic structures of the first 18
elements
all atoms contain orbitals similar to those found in
hydrogen
systematically placing electrons in these hydrogen
like orbitals, following the guidelines:
1. no more than two electrons can occupy one
orbital
2. electrons occupy the lowest energy orbitals
available s < p < d < f for a given n value
3. each orbital on a sublevel is occupied by a
single electron before a second electron enters
ex. atomic structure diagrams of F, Na and Mg
there are two ways to show the arrangement of
the electrons in the orbitals:
i) electron configuration
number of electrons in sublevel orbitals
2p6
principal energy level
type of orbital
ii) orbital diagram
orbital □
spin ↑ ↓
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valence electrons – the electrons un the outmost
(highest) energy level of an atom
ex. O 1s2 2s2 2p4 6 valence electrons
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2
2
6
2
Mg 1s 2s 2p 3s
2 valence electrons
10.6 Electron structures and the periodic
table
1869 Mendeleev & Meyer
periodic arrangements of the elements based on
increasing atomic masses
periodic table
period – horizontal row
the outmost energy level
group or family – vertical column
elements behave in a similar manner
IA ~ VIIA, IB ~ VIIB, VIII, noble gases
1 ~ 18
representative elements – A group elements
transition elements – B group elements
IA – alkali metals IIA – alkaline earth metals
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VIIA – halogens
the valence electron configurations for H ~ Ar
• the valence electron configuration for the
element in each column is the same, but the
number for the energy level is different
• the chemical behavior and properties of elements
in a particular family are similar and must be
associated with the electron configuration
abbreviated electron configuration
[He] 2s2 2p1
B 1s2 2s2 2p1
Cl 1s2 2s2 2p6 3s2 3p5
[Ne] 3s2 3p5
Na 1s2 2s2 2p6 3s1
[Ne] 3s1
n = 4 K 1s2 2s2 2p6 3s2 3p6 4s1 [Ar] 4s1
Ca 1s2 2s2 2p6 3s2 3p6 4s2 [Ar] 4s2
element 21 ~ 30 transition elements
electrons are placed in the 3d orbitals
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arrangement of elements according to the
sublevel being filled
inner transition elements
lanthanide series – 4f
actinide series – 5f
ex. 10.1 write the electron configuration for
P and Sn
[Ne]3s23p3
P 1s22s22p63s23p3
Sn 1s22s22p63s23p64s23d104p65s24d105p2
[Kr] 5s24d105p2
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groups of elements show similar chemical
properties because of the similarity of these
outmost electron configurations
the periodic table illustrates several important
points:
1. the number of the period corresponds with
the highest energy level occupied by
electrons
2. the group numbers for the representative
elements are equal to the total number of
outmost electrons in the atom
3. the elements of a family have the same
outmost electron configuration, but in
different energy level
4. the elements within each of the s, p, d, f
blocks are filling the s, p, d, f orbitals
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5. within the transition elements some
discrepancies in the order of filling occur