Transcript Chapter 5

C H E M I
S T R Y
Chapter 5
Periodicity and Atomic Structure
Light and the Electromagnetic
Spectrum
Electromagnetic energy (“light”) is characterized by
wavelength, frequency, and amplitude.
Light and the Electromagnetic
Spectrum
n
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l
Chapter 5/3
Light and the Electromagnetic
Spectrum
Wavelength x Frequency = Speed
l
m
x
n
= c
1
m
s
s
c is defined to be the rate of travel of all electromagnetic
energy in a vacuum and is a constant value—speed of light.
c = 3.00 x 108
m
s
Examples
 The light blue glow given off by mercury streetlamps has a
frequency of 6.88 x 1014 s-1 (or, Hz). What is the wavelength in
nanometers?
1.– Particlelike Properties of Electromagnetic Radiation: The Planck Equatio
The energy level of Hydrogen
Particlelike Properties of
Electromagnetic Energy
 Photoelectric Effect: Irradiation
of clean metal surface with light
causes electrons to be ejected from
the metal. Furthermore, the
frequency of the light used for the
irradiation must be above some
threshold value, which is different
for every metal.
Particlelike Properties of
Electromagnetic Energy
Particlelike Properties of
Electromagnetic Energy
Einstein explained the effect by assuming that a beam of light
behaves as if it were a stream of particles called photons.
E = hn
E
ν
Electromagnetic energy (light) is quantized.
h (Planck’s constant) = 6.626 x 10-34 J s
Emission of Energy by Atom
 How does atom emit light?
 Atoms absorbs energy
 Atoms become excited
 Release energy
 Higher-energy photon –>shorter wavelength
 Lower-energy photon -> longer wavelength
Examples
 What is the energy (in kJ/mol) of photons of radar waves
with ν = 3.35 x 108 Hz?
 What is the energy (in kJ/mol) of photons of an X-ray with
λ = 3.44 x 10-9 m?
Particlelike Properties of
Electromagnetic Energy
 Niels Bohr proposed in 1914 a
model of the hydrogen atom as a
nucleus with an electron circling
around it.
 In this model, the energy levels of
the orbits are quantized so that only
certain specific orbits corresponding
to certain specific energies for the
electron are available.
Wavelike Properties of Matter
Louis de Broglie in 1924 suggested that, if light can behave in some
respects like matter, then perhaps matter can behave in some respects like
light.
In other words, perhaps matter is wavelike as well as particlelike.
l= h
mv
The de Broglie equation allows the calculation of a “wavelength” of an
electron or of any particle or object of mass m and velocity v.
Examples
 What is the de Broglie wavelength (in meters) of a small car
with a mass of 11500 kg traveling at a speed of 55.0 mi/h
(24.6 m/s)?
 What velocity would an electron (mass = 9.11 x 10-31kg)
need for its de Broglie wavelength to be that of red light
(750 nm)?
Quantum Mechanics and the
Heisenberg Uncertainty Principle
In 1926 Erwin Schrödinger proposed the quantum
mechanical model of the atom which focuses on the
wavelike properties of the electron.
In 1927 Werner Heisenberg stated that it is impossible to know
precisely where an electron is and what path it follows—a
statement called the Heisenberg uncertainty principle.
Quantum Mechanics and the Heisenberg
Uncertainty Principle

Heisenberg Uncertainty Principle – both the position (Δx) and the
momentum (Δmv) of an electron cannot be known beyond a certain level of
precision
1.
(Δx) (Δmv) > h
4π
2.
Cannot know both the position and the momentum of an
electron with a high degree of certainty
3. If the momentum is known with a high degree of certainty
i.
Δmv is small
ii.
Δ x (position of the electron) is large
4.
If the exact position of the electron is known
i.
Δmv is large
ii.
Δ x (position of the electron) is small
Wave Functions and Quantum
Numbers
Wave
equation
solve
Wave function
or orbital (Y)
Probability of finding
electron in a region
of space (Y 2)
A wave function is characterized by three parameters called
quantum numbers, n, l, ml.
Wave Functions and Quantum
Numbers
Principal Quantum Number (n)
• Describes the size and energy level of the orbital
• Commonly called shell
• Positive integer (n = 1, 2, 3, 4, …)
• As the value of n increases:
• The energy increases
• The average distance of the e- from the nucleus
increases
Wave Functions and Quantum
Numbers
Angular-Momentum Quantum Number (l)
• Defines the three-dimensional shape of the orbital
• Commonly called subshell
• There are n different shapes for orbitals
• If n = 1 then l = 0
• If n = 2 then l = 0 or 1
• If n = 3 then l = 0, 1, or 2
• etc.
• Commonly referred to by letter (subshell notation)
• l=0
s (sharp)
• l=1
p (principal)
• l=2
d (diffuse)
• l=3
f (fundamental)
• etc.
Wave Functions and Quantum
Numbers
Magnetic Quantum Number (ml )
• Defines the spatial orientation of the orbital
• There are 2l + 1 values of ml and they can have any
integral value from -l to +l
• If l = 0 then ml = 0
• If l = 1 then ml = -1, 0, or 1
• If l = 2 then ml = -2, -1, 0, 1, or 2
• etc.
Wave Functions and Quantum
Numbers
Wave Functions and Quantum
Numbers
Wave Functions and Quantum
Numbers
 Identify the possible values for each of the three
quantum numbers for a 4p orbital.
 Give orbital notations for electrons in orbitals with the
following quantum numbers:
a)
n = 2, l = 1, ml = 1
b) n = 4, l = 3, ml =-2
 Give the possible combinations of quantum numbers for
the following orbitals:

A 3s orbital
b) A 4d orbital
The Shapes of Orbitals
Node: A surface of zero
probability for finding the
electron.
The Shapes of Orbitals
Electron Spin and the Pauli
Exclusion Principle
Electrons have spin which gives rise to a tiny magnetic field and to a spin
quantum number (ms).
Pauli Exclusion Principle: No two electrons in an atom can have the same four
quantum numbers.
Orbital Energy Levels in
Multielectron Atoms
Electron Configurations of
Multielectron Atoms
Effective Nuclear Charge (Zeff): The nuclear charge actually
felt by an electron.
Zeff = Zactual - Electron shielding
Electron Configurations of
Multielectron Atoms
Electron Configuration: A description of which orbitals are occupied by
electrons.
Degenerate Orbitals: Orbitals that have the same energy level. For
example, the three p orbitals in a given subshell.
Ground-State Electron Configuration: The lowest-energy
configuration.
Aufbau Principle (“building up”): A guide for determining the filling
order of orbitals.
Electron Configurations of
Multielectron Atoms
Rules of the aufbau principle:
1.
2.
3.
Lower-energy orbitals fill before higher-energy orbitals.
An orbital can only hold two electrons, which must have opposite
spins (Pauli exclusion principle).
If two or more degenerate orbitals are available, follow Hund’s rule.
Hund’s Rule: If two or more orbitals with the same energy are available,
one electron goes into each until all are half-full. The electrons in the halffilled orbitals all have the same spin.
Electron Configurations of
Multielectron Atoms
Electron
Configuration
H:
1s1
1 electron
s orbital (l = 0)
n=1
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Electron Configurations of
Multielectron Atoms
Electron
Configuration
H:
1s1
He:
1s2
2 electrons
s orbital (l = 0)
n=1
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Electron Configurations of
Multielectron Atoms
Electron
Configuration
H:
1s1
He:
1s2
Lowest energy to highest energy
Li:
1s2 2s1
1 electrons
s orbital (l = 0)
n=2
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Chapter
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Electron Configurations and the
Periodic Table
Valence Shell: Outermost shell.
Li: 2s1
Na: 3s1
Cl: 3s2 3p5
Br: 4s2 4p5
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Electron Configurations and the
Periodic Table
 Give expected ground-state electron configurations for the
following atoms, draw – orbital filling diagrams and
determine the valence shell
 O (Z = 8)
 Ti (Z = 22)
 Sr (Z = 38)
 Sn (Z = 50)
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Chapter
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Periodic Properties: Atomic
Radii
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column
radius
row
radius
Chapter
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Periodic Properties: Atomic
Radii
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Chapter
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Examples
 Which atom in each of the following pairs would you expect
to be larger?
 Mg or Ba
W or Au