Periodic Table Song
Development of the Periodic Table
Chapter 5
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Mid-1800’s, several scientists placed known
elements in order based on different criteria.
Mendeleev’s and Meyer’s versions, 1869
Periodicity
y and
Atomic Structure
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Modern Arrangement
What’s Going on Inside the Atom?
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Light
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Once scientists developed a logical order for
the elements they began studying the structure
and composition of individual atoms.
They used substances’
substances interactions with light to
explain the structure of atoms and develop a
model to explain how atoms affected properties
of light.
In order to understand interactions, we must
understand behavior of light.
Light and the EM Spectrum
EM waves
Light is typically described as traveling in waves
(similar to water); All electromagnetic (EM) waves
(including light) are made of two components:
electric and magnetic
EM waves travel at the speed of light, c (2.997924 x
108 m/s ≈ 3.00 x 108 m/s)
c = λν (Know these variables!)
c = speed of light;
λ (lambda) = wavelength (m, nm);
ν (nu) = frequency (1/s, s-1, Hz)
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EM Spectrum
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Describing Atoms
Different colors of light correspond to different
wavelengths in the visible portion of the EM
spectrum. Two wavelengths (λ) are shown below.
Determine the frequency (ν) for each wave.
Blue light
Red light
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Classical descriptions:
 Dalton: atoms are hard particles, all atoms of the
same element are the same
 Thomson: atoms are divisible (electrons in atoms)
 Rutherford: positively charged nucleus
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New view of atomic behavior
 Planck: Blackbody radiation – heat solids to red or
white heat, matter did not emit energy continuously; in
whole-number multiples of certain quantities
 Matter absorbs or emits energies in packets - quanta
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1 nm = 1 x 10-9 m
OR
1 x 109 nm = 1 m
From Classical to Quantum Theory
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Quantum has come to mean small; originated
from Planck’s observation of quantized energy
Einstein used this theory to observe metals
reacting to different colors of light –
Photoelectric Effect: electrons are ejected from
the surface of certain metals exposed to light at
a certain minimum frequency
 Blue light (ν = 6.7 x
Hz) causes Na to emit
electrons, red light (ν = 4.0 x 1014 Hz) does not
Wave-Particle Duality
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Based on photoelectric effect, light acts as a
wave but also exists as a stream of particles
called photons
Energy
gy of p
photons is p
proportional
p
to frequency,
q
y,
inversely proportional to wavelength
E = hν =
1014
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hc
λ
h = 6.626 x 10-34 J•s
J = kg • m2 / s2
Photoelectric Effect
Calculation Practice c = λν; E = hν
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1) Which has a higher frequency: light from a red
stoplight with a wavelength of 750 nm or a yellow light
with a wavelength of 600 nm?
2) What is the wavelength of a radio station’s waves
transmitting at a frequency of 101.5 MHz (megahertz)?
(FM radio waves range from 30 – 300 MHz.)
3) Red lights at traffic stops have wavelengths of about
650 nm. What is the frequency (in Hz) of this light?
4) Calculate the energy (in Joules) of a photon with a
wavelength of 5.00 x 104 nm (infrared region).
Worked Ex. 5.1, 5.4; Problems 5.1, 5.2, 5.3, 5.7, 5.8
Wavelike Properties of Matter
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de Broglie: If light can behave like a wave and
a particle, then matter (i.e., electrons) can
behave like a wave
If an electron behaves like a standing wave
wave,
then it can only have specific wavelengths
Can calculate wavelength for matter if we know
its velocity (use v instead of c): λ = h / m v
(This is the de Broglie equation.)
 h = Planck’s constant, m = mass (electron’s have
constant mass: 9.11x10-31 kg), v = velocity (speed)
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Worked Ex. 5.5; Problem 5.9
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Calculation Practice
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The energy of a photon is 5.87 x 10-20 J. What
is the frequency of the photon?
What is the wavelength of an electron that
travels at 34
34.7
7 m/s and has a mass of 9
9.11
11 x
10-31 kg?
A 0.143 kg baseball is thrown at a velocity of
42.5 m/s. Calculate the wavelength of the
baseball. How does the baseball’s wavelength
compare to the electron from the example
above?
Bohr’s Model of the Atom
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Heisenberg Uncertainty Principle
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If electrons have wavelike properties, then we
can’t know both its position and velocity. In
order to determine the position of an electron,
we hit it with a photon of light
light, but this will
change its position and velocity.
Bohr’s Model of the Atom
Bohr sought to reconcile these views of the
electron.
 Developed the planetary analogy of atoms.
 Electrons orbit around the nucleus like planets
around the sun.
 Electrons travel in discrete, quantized circular orbits;
like going up or down stairs.
 Each orbit has a specific energy associated with it,
labeled as n = 1, 2, etc.
 Ground state is the lowest energy level for an atom
(n = 1).
Bohr’s Model of the Atom
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Transitions between Energy Levels
When an atom absorbs energy, an electron
can jump from a lower energy level to a higher
energy level.
When an atom emits (releases) energy
energy, an
electron drops from a higher energy level to a
lower energy level. This process sometimes
gives off energy as visible light.
H e- transitions
Eng. Color
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Visible Light
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Emission Spectra
White light we see consists of all colors in the
visible spectrum. Use a prism (or CD) to break
them up.
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cont. spect.
white light
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Hydrogen
Æ
Each element gives
off unique spectrum
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Demo:
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Gas Discharge
Di h
Tubes
 Each element has its
Light given off by atoms
doesn’t necessarily correspond to all visible colors.
own individual emission
spectrum. This allowed
scientists to identify
elements in different
minerals.
Flame tests
Spectra of Elements:
http://www.wwnorton.com/college/chemistry/chemconnections/BlueLight/pages/elements.html
Emission Spectra of Elements
Quantum Mechanical Model
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The Bohr model worked well for hydrogen, but
failed for elements with more than one proton
and one electron.
Quantum Mechanics was developed (by
Schrödinger in the 1920’s) to describe the
motion of subatomic particles
 Did not attempt to describe position of particles;
used mathematical equations to describe the
probability of finding the particles
 The probability density (map of likely locations) is
the “electron cloud”
Figure 7.8
Atomic Orbitals
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Quantum Numbers
Quantum Mechanics Movie
The region of highest probability for finding an
electron is an “electron cloud”. This region of
high probability is called an atomic orbital.
Each orbital holds at most 2 electrons
electrons.
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There are 4 quantum number that describe the
size, shape, and location of electrons
We use these numbers to describe where
electrons are found for an atom
atom. Can also use
the periodic table!!!
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The Principal Quantum Number, n
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Modern atom.exe
 describes distance of the electron from the nucleus;
called shells
e- orbit vs
e- cloud
 n = 1, 2, 3, etc; larger number is farther from
nucleus
 n corresponds to a row in the periodic table
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Orbitals and Quantum Numbers
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Shapes of Orbitals
The Angular Momentum Quantum Number, l
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 In each row of the periodic table are different groups of orbitals
with different shapes. These groups of orbitals are called
subshells and labeled s, p, d, and f.
 s subshells are spherical (first two columns)
 p subshells
b h ll are d
dumb-bell
b b ll shaped
h
d (l
(lastt six
i columns)
l
)
 d subshells are intersecting dumb-bells (transition metals)
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The Magnetic Quantum Number, ml
 1 s orbital in a
 describes the orientation of the orbital with respect to x, y,
subshell
and z axes
 s, p, and d orbitals have different shapes and therefore
different orientations
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s orbitals are
spherical; white rings
are nodes (regions
where an electron
won’t be found)
The Spin Quantum Number, ms
s orbital
 describes the spin of an electron in an orbital (shown as up
and down arrows in orbital diagrams)
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Shapes of Orbitals
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2px orbital 2py orbital
2pz orbital
Shapes of Orbitals
p orbitals are dumb-bells (2 lobes); node
between lobes
 3 p orbitals in a subshell
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d orbitals: intersecting dumb-bells (4 lobes);
nodes between lobes
 5 d orbitals in a subshell
p orbital
d orbital
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Shells and Subshells
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Electron Configurations
The first shell (row) has 1 subshell (s)
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 s Æ only 1 orbital
 An s subshell can hold at most 2 electrons
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Arrangement of subshells in the Periodic Table
The 2nd shell (row) has 2 subshells (s and p)
 p Æ set of 3 orbitals
 A p subshell can hold at most 6 electrons
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The 3rd shell (row) has 3 subshells (s, p, and d)
 d Æ set of 5 orbitals
 A d subshell can hold at most ? electrons
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The 4th shell (row) has 4 subshells (s, p, d, and f)
 f Æ set of 7 orbitals
 What is the maximum number of electrons allowed in the f
subshell?
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Things that make you go hmmm….
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Energies of Orbitals
What is the maximum number of:
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 electrons allowed in the 2px orbital?
 subshells allowed in the 4th shell?
In hydrogen, all shells
are equivalent in
energy.
 electrons allowed in the 3d subshell?
 electrons allowed in the 4d subshell?
 electrons allowed in the 3p subshell?
 electrons allowed in the 3rd shell?
Energies of Orbitals
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Aufbau Principle
In many-electron
models, the energy
levels depend on the
shell and subshell
subshell.
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Aufbau principle: start with the nucleus and
empty orbitals, then “build” up the electron
configuration using orbitals of increasing
energy.
energy
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Electron Configurations
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Electron Configurations
Arrangement of subshells in the Periodic Table
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Arrangement of subshells in the Periodic Table
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Electron Configurations
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Electron Configurations
Arrangement of subshells in the Periodic Table
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Arrangement of subshells in the Periodic Table
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Electron Configurations
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Writing Electron Configurations
Arrangement of subshells in the Periodic Table
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Write electron
configurations for the
following atoms.
H
He
Li
Be
B
N
O
Ne
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Na
Al
S
Ar
K
Sc
Ti
Zn
Br
Worked Ex. 5.10, 5.11;
Problems 5.18, 5.19, 5.20
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Electron Configurations (Figure 5.17)
Valence Electrons
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Electrons in the outermost shell.
 1s2 2s2 2p6
 1s2 2s2 2p6 3s2 3p5
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Identify the valence electrons (v
(v. e
e.)) in the
following configurations:
 1s2 2s2 2p6 3s2
 1s2 2s2 2p33s2
1s2 2s2
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Shorthand Notation
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Electron Configuration
Rather than writing out complete electron
configurations, we can use the previously filled
shell (noble gas) and show the valence
electrons (v
(v. e
e.):
):
P: 1s2 2s2 2p6 3s2 3p3 Æ [Ne] 3s2 3p3 (5 v. e.)
Write the shorthand notation for:
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Some exceptions to the Aufbau order…
What are the expected electron configurations
for Cr and Cu?
Filled and half-filled
half filled d subshells seem to be
especially stable.
Cr: 1s2 2s2 2p6 3s2 3p6 4s1 3d5
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Cu: 1s2 2s2 2p6 3s2 3p6 4s1 3d10
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 Also true for Mo and W
 Ca
 Cl
e_config.
 Also true for Ag and Au
 Sr
 Fe
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Orbital Diagrams
Hund’s Rule
Orbital diagrams
are pictorial
representations
of electron
configurations.
Electron Configurations
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If two or more orbitals (i.e., a p or d orbital) with
the same energy are available, one electron
goes into each orbital until they have to pair up.
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For example, an atom with 2 p electrons: 1
electron will go into the first (px) orbital, the
next electron will go into the second (py) orbital.
 Fighting sibling analogy
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Pauli Exclusion Principle
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Pauli Exclusion Principle: no two electrons
can have the same values of all 4 quantum
numbers
Describes what happens when electrons share
an orbital.
 Only two electrons can occupy a single orbital and
they must have opposite spin (i.e., the 4th quantum
number). The first electron is designated as positive
spin (up arrow), the second electron in that orbital
has negative spin (down arrow).
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