Download Ch 4 PPT - mvhs

Chemistry, The Central Science, 11th edition
Theodore L. Brown; H. Eugene LeMay, Jr.;
and Bruce E. Bursten
Chapter 4
Electronic Structure
of Atoms
John D. Bookstaver
St. Charles Community College
Cottleville, MO
© 2009, Prentice-Hall, Inc.
Waves
• To understand the electronic structure of atoms,
one must understand the nature of
electromagnetic radiation.
• The distance between corresponding points on
adjacent waves is the wavelength ().
© 2009, Prentice-Hall, Inc.
Waves
• The number of waves passing a given
point per unit of time is the
frequency ().
• For waves traveling at the same
velocity, the longer the wavelength,
the smaller the frequency.
– Wavelength: It is the distance
between two consecutive peaks
or troughs in a wave.
– Frequency: It indicates how many
waves pass a given point per
second.
– Speed: It indicates how fast a
given peak is moving through the
space.
– Speed of light ,c= ,where
=wave length and  =frequency
© 2009, Prentice-Hall, Inc.
EMR
Electromagnetic Radiation:
Electromagnetic radiation is
one of the ways in which
energy travels through
space. All forms of EMR
compose the
electromagnetic radiation
spectrum, which includes
sun rays, microwaves, Xrays, visible spectrum, UV
rays and IR rays.
• Some characteristics of EMR are:
• All electromagnetic radiation
moves at a constant speed of
about 3.0 X 108 m/s.
• All EMR exhibit wave like
behavior. Waves have three
primary characteristics:
– Wavelength: It is the distance
between two consecutive peaks
or troughs in a wave.
– Frequency: It indicates how
many waves pass a given point
per second.
– Speed: It indicates how fast a
given peak is moving through
the space.
– Speed of light ,c= ,where
=wave length and 
=frequency
Electromagnetic Radiation
• All electromagnetic
radiation travels at the
same velocity: the
speed of light (c),
3.00  108 m/s.
• Therefore,
c = 
© 2009, Prentice-Hall, Inc.
Refer to the following EMR spectrum.
(visible spectrum Roy G. BiV)
• Imagine you have
invented a machine that
allows you to see all types
of EMR. Make a list of
type of EMR you might
see in your home.
• What will happen if you
change the red light in the
dark room for photo
processing with the yellow
light and why?
Wave Nature of Light
• Before the concept
of quantization of
energy, the wave
like nature of
light/energy was
widely accepted.
• Two properties that
exhibit the wave
like behavior of
light are
interference and
diffraction.
Animation Diffraction: http://www.acoustics.salford.ac.uk/feschools/waves/diffract3.htm
Animation for Interference:
http://id.mind.net/~zona/mstm/physics/waves/interference/waveInterference1/WaveIn
terference1.html
The Nature of Energy
• The wave nature of light
does not explain how an
object can glow when its
temperature increases.
• Max Planck explained it
by assuming that energy
comes in packets called
quanta.
© 2009, Prentice-Hall, Inc.
Planck’s Theory of Quantization of Energy
• Planck’s theory: Before
Planck’s theory, the wave
model of the light was
widely accepted. But it was
unable to explain some
phenomenon for example
change in the radiation
(wave length) emitted by an
object with the change in
temperature.
• To explain this, Planck
suggested that the energy
transfer or exchange is not a
continuous process, but is
done in small packets of
energy called by him as
quantum.(Word quantum
means fixed amount.) So, he
introduced the concept of
quantization of energy.
• According to Planck’s theory,
E= hv,
where E= energy of radiation
h= Planck’s constant
v= frequency of radiation
Quantization of Energy: Max Plank
• www
•
www.physics.usceud
• Where else do you
see quantization in
real life?
The Nature of Energy
• Therefore, if one knows the
wavelength of light, one can
calculate the energy in one
photon, or packet, of that
light:
c = 
E = h
© 2009, Prentice-Hall, Inc.
The Nature of Energy
Another mystery in
the early 20th century
involved the emission
spectra observed from
energy emitted by
atoms and molecules.
© 2009, Prentice-Hall, Inc.
The Nature of Energy
• Einstein used this assumption
to explain the photoelectric
effect.
• He concluded that energy is
proportional to frequency:
E = h
where h is Planck’s constant,
6.626  10−34 J-s.
© 2009, Prentice-Hall, Inc.
Photoelectric Effect: It refers to the emission of electrons from a metal,
when the light shines on the metal. For each metal the frequency of
light needed to release the electrons is different. But the wave theory of
light could not explain it. The photoelectric effect led scientists to think
about the dual nature of light i.e. as a wave and a particle both.
Animation 1:
http://www.upscale.utoronto.ca/GeneralInterest/Harrison/Flash/Nuclear/XRayIntera
ct/XRayInteract.html
Animation 2: http://faculty.ucc.edu/chemistry-pankuch/Photoelectric/PE3.html
Dual Wave- Particle Nature of Light
• Dual nature (Wave- Particle nature) of light:
Albert Einstein in 1905, expanded on Planck’s
theory by introducing the idea of the dual
nature of the radiation. Albert Einstein called
the quantum of energy in light as ‘photons’.
Light as wave
Light as particle
DUAL NATURE OF LIGHT
http://www.nobel.se/physics/educational/tools/quantum/w-pimages/wp.gif
http://www.astronomy.org.nz/events/monthly_meetings/Reviews/2001/gigant4.gif
http://www.tpub.com/doenuclearphys/nuclear%20physics%20and%20reactor%20theory_files/image003.
jpg
•
Photoelectric
Effect
Max Planck suggested that the
hot object emits energy in small,
specific amounts called quanta.
• 1905, Einstein said
electromagnetic radiation has a
dual wave- particle nature.
http://www.shsu.edu/%7Echm_tgc/sounds/flashfiles/pee.s
wf
The Nature of Energy
• For atoms and molecules
one does not observe a
continuous spectrum, as
one gets from a white
light source.
• Only a line spectrum of
discrete wavelengths is
observed.
© 2009, Prentice-Hall, Inc.
Continuous and Line Spectra (Absorption and Emission
Spectrum- Line Spectra)
(http://images.google.com/imgres?imgurl=http://csep10.phys.utk.edu/astr162/lect/light/spectra.gif&imgrefurl=http://csep10.phys.utk.edu/astr162/lect/light/a
bsorption.html&h=240&w=450&sz=33&hl=en&start=2&tbnid=TaN57QO8MhMG4M:&tbnh=66&tbnw=124&prev=/images%3Fq%3Dabsorption%2Band%2Bemission%2Bspectr
um%26svnum%3D10%26hl%3Den%26lr%3D%26sa%3DN)
Bohr Model Of the Hydrogen Atom
• Ephoton=hv
The energy levels of Hydrogen ( As
explained by Bohr’s Model) :
•An excited atom can release
some or all of its excess energy
by emitting a photon, thus
moving to a lower energy state.
•The lowest possible energy
state of an atom is called the
‘ground state’.
•Different wavelengths of light
carry different amount of energy
per photon. Ex. A beam of red
light has a lower energy
photons than beam of blue light.
The Nature of Energy
•
Niels Bohr adopted Planck’s
assumption and explained
these phenomena in this way:
1. Electrons in an atom can only
occupy certain orbits
(corresponding to certain
energies).
Animation : http://pokok.mysch.net/~chemistry/AL2.htm#Animation
Animation:
http://physics.gac.edu/~chuck/PRENHALL/Chapter%2031/AABXTEJ0.html
© 2009, Prentice-Hall, Inc.
The Nature of Energy
•
Niels Bohr adopted Planck’s
assumption and explained
these phenomena in this way:
2. Electrons in permitted orbits
have specific, “allowed”
energies; these energies will not
be radiated from the atom.
© 2009, Prentice-Hall, Inc.
The Nature of Energy
•
Niels Bohr adopted Planck’s
assumption and explained these
phenomena in this way:
3. Energy is only absorbed or emitted
in such a way as to move an
electron from one “allowed”
energy state to another; the
energy is defined by
E = h
© 2009, Prentice-Hall, Inc.
Light Equations
•
•
•
•
•
c=  
c is the speed of light (3.0
x 108 m/s)
lambda is the wave length
(in m, cm, or nm)
 represents the
frequency (waves/s or
hertz)
http://www.astronomynot
es.com/light/s3.htm
• E=h 
• E is energy (in joules)
• h is Planck’s Constant
(h= 6.626 x 10-34J s)
•  is frequency
• E=h(c/ )
http://sunflower.bio.indiana.edu/~rhangart/courses/b373/lecturenotes/photomorph/sinewave.gif
http://hackensackhigh.org/~rkc2/diffraction.jpg
• . Heisenberg
Uncertainty
Principle: States
that it is impossible
to determine
simultaneously
both the position
and velocity of an
electron or any
other particle.
• Dual Wave-Particle
nature by De
Broglie
• Quantum Theory: Describes
mathematically the wave
properties of electrons or other
very small particles treating e as
waves and using Heisenberg’s and
De Broglie’s principles.
Quantum Mechanical Model of Atom (Schrodinger’s Model)
Erwin Schrödinger, in developing a quantum-mechanical model for the
atom, began with a classical equation for the properties of waves. He
modified this equation to take into account the mass of a particle and the de
Broglie relationship between mass and wavelength. The important
consequences of the quantum-mechanical view of atoms are the following:
(http://www.cartage.org.lb/en/themes/sciences/chemistry/Generalchemistry/Atomic/Electronicstructure/Electronicstructures/Quantum/Quantum.htm-)
1.
2.
3.
4.
The energy of electrons in atoms is quantized.
The number of possible energy levels for electrons in atoms of
different elements is a direct consequence of wave-like properties of
electrons.
The position and momentum of an electron cannot both be
determined simultaneously.
The region in space around the nucleus in which an electron is most
probably located is what can be predicted for each electron in an
atom. Electrons of different energies are likely to be found in different
regions. The region in which an electron with a specific energy will
most probably be located is called an atomic orbital.
Quantum Mechanics
• Erwin Schrödinger
developed a mathematical
treatment into which both
the wave and particle
nature of matter could be
incorporated.
• It is known as quantum
mechanics.
© 2009, Prentice-Hall, Inc.
Quantum Mechanics
• The wave equation is
designated with a lower case
Greek psi ().
• The square of the wave
equation, 2, gives a
probability density map of
where an electron has a
certain statistical likelihood of
being at any given instant in
time.
© 2009, Prentice-Hall, Inc.
Quantum Numbers
• Schrodinger arrived at three functions while solving the
wave equation for electrons and came up with three
quantum numbers, each corresponding to a property of
atomic orbital. The fourth quantum number was later
introduced to clarify the spin of electrons.
use(http://www.google.com/search?hl=en&q=wave+mechanical+model)
• Solving the wave equation gives a set of wave
functions, or orbitals, and their corresponding
energies.
• Each orbital describes a spatial distribution of electron
density.
• An orbital is described by a set of three quantum
numbers.
© 2009, Prentice-Hall, Inc.
Principal Quantum Number (n)
• The principal quantum number, n, describes
the energy level on which the orbital resides.
• The values of n are integers ≥ 1.
© 2009, Prentice-Hall, Inc.
Angular Momentum Quantum
Number (l)
• This quantum number defines the shape of
the orbital.
• Allowed values of l are integers ranging
from 0 to n − 1.
• We use letter designations to communicate
the different values of l and, therefore, the
shapes and types of orbitals.
© 2009, Prentice-Hall, Inc.
Angular Momentum Quantum Number (l)
Value of l
0
1
2
3
Type of orbital
s
p
d
f
© 2009, Prentice-Hall, Inc.
Mosby-Year Book Inc.
Magnetic Quantum Number (ml)
• The magnetic quantum number describes
the three-dimensional orientation of the
orbital.
• Allowed values of ml are integers ranging
from -l to l:
−l ≤ ml ≤ l.
• Therefore, on any given energy level, there
can be up to 1 s orbital, 3 p orbitals, 5 d
orbitals, 7 f orbitals, etc.
© 2009, Prentice-Hall, Inc.
Magnetic Quantum Number (ml)
• Orbitals with the same value of n form a shell.
• Different orbital types within a shell are subshells.
© 2009, Prentice-Hall, Inc.
Orbitals and Electron Capacity of the First Four Principle Energy
Levels
Maximum
Principle
Number of Number of
Type of
number of
energy level
orbitals per orbitals per
sublevel
electrons
(n)
type
level(n2)
(2n2)
1
2
3
4
s
1
s
1
p
3
s
1
p
3
d
5
s
1
p
3
d
5
f
7
1
2
4
8
9
18
16
32
s Orbitals
• The value of l for s
orbitals is 0.
• They are spherical in
shape.
• The radius of the sphere
increases with the value
of n.
© 2009, Prentice-Hall, Inc.
s Orbitals
Observing a graph of
probabilities of finding an
electron versus distance
from the nucleus, we see
that s orbitals possess
n−1 nodes, or regions
where there is 0
probability of finding an
electron.
© 2009, Prentice-Hall, Inc.
p Orbitals
• The value of l for p orbitals is 1.
• They have two lobes with a node between them.
© 2009, Prentice-Hall, Inc.
d Orbitals
• The value of l for a d
orbital is 2.
• Four of the five d
orbitals have 4
lobes; the other
resembles a p
orbital with a
doughnut around
the center.
© 2009, Prentice-Hall, Inc.
Energies of Orbitals
• For a one-electron
hydrogen atom,
orbitals on the same
energy level have the
same energy.
• That is, they are
degenerate.
© 2009, Prentice-Hall, Inc.
Energies of Orbitals
• As the number of
electrons increases,
though, so does the
repulsion between
them.
• Therefore, in manyelectron atoms, orbitals
on the same energy
level are no longer
degenerate.
© 2009, Prentice-Hall, Inc.
Spin Quantum Number, ms
• In the 1920s, it was
discovered that two
electrons in the same
orbital do not have exactly
the same energy.
• The “spin” of an electron
describes its magnetic
field, which affects its
energy.
© 2009, Prentice-Hall, Inc.
Spin Quantum Number, ms
• This led to a fourth
quantum number, the spin
quantum number, ms.
• The spin quantum number
has only 2 allowed values:
+1/2 and −1/2.
© 2009, Prentice-Hall, Inc.
Pauli Exclusion Principle
• No two electrons in the
same atom can have
exactly the same energy.
• Therefore, no two
electrons in the same atom
can have identical sets of
quantum numbers.
© 2009, Prentice-Hall, Inc.
Electron Configurations
• This shows the
distribution of all
electrons in an atom.
• Each component consists
of
– A number denoting the
energy level,
© 2009, Prentice-Hall, Inc.
Electron Configurations
• This shows the
distribution of all
electrons in an atom
• Each component consists
of
– A number denoting the
energy level,
– A letter denoting the type
of orbital,
Animation: http://intro.chem.okstate.edu/WorkshopFolder/Electronconfnew.html
© 2009, Prentice-Hall, Inc.
Electron Configurations
• This shows the
distribution of all
electrons in an atom.
• Each component consists
of
– A number denoting the
energy level,
– A letter denoting the type
of orbital,
– A superscript denoting the
number of electrons in
those orbitals.
© 2009, Prentice-Hall, Inc.
Orbital Diagrams
• Each box in the diagram
represents one orbital.
• Half-arrows represent the
electrons.
• The direction of the
arrow represents the
relative spin of the
electron.
© 2009, Prentice-Hall, Inc.
Hund’s Rule
“For degenerate
orbitals, the lowest
energy is attained
when the number of
electrons with the
same spin is
maximized.”
© 2009, Prentice-Hall, Inc.
Periodic Table
• We fill orbitals in
increasing order of
energy.
• Different blocks on the
periodic table (shaded in
different colors in this
chart) correspond to
different types of
orbitals.
© 2009, Prentice-Hall, Inc.
Some Anomalies
Some irregularities
occur when there
are enough
electrons to half-fill
s and d orbitals on
a given row.
© 2009, Prentice-Hall, Inc.
Some Anomalies
For instance, the
electron
configuration for
copper is
[Ar] 4s1 3d5
rather than the
expected
[Ar] 4s2 3d4.
© 2009, Prentice-Hall, Inc.
Some Anomalies
• This occurs
because the 4s and
3d orbitals are very
close in energy.
• These anomalies
occur in f-block
atoms, as well.
© 2009, Prentice-Hall, Inc.
Quantum Number Terms
• Ground State: Lowest energy
state of an atom
• Excited State: A state in which
an atom has a higher potential
energy then it has in its ground
state
• Orbital: A 3D region around the
nucleus that indicates the
probable location of an
electron
• Quantum Numbers: Specify
the properties of atomic
orbitals and the properties of
electrons in orbitals
• Principle Quantum Number:
(n) Indicates the main energy
level occupied by the electron.
• Angular Momentum Quantum
Number: (l) Indicates the shape
of the orbital.
• Magnetic Quantum Number:
(m) Indicates the
orientation of an orbital
around the nucleus.
• Spin Quantum Number: (+1/2,
-1/2) Indicates the two
fundamental spin states of an
electron in an orbital
http://www.tannerm.com/Quick_atom/hund.gif
Electron Configuration Theories
• Aufbau’s Principle: An
electron occupies the
lowest-energy orbital that
can receive it.
• Pauli’s Exclusion Principle:
No two electrons in the
same atom can have the
same set of four quantum
numbers.
• Hund’s Rule: Orbitals of
equal energy are each
occupied by one electron
before any orbital is
occupied by a second
electron, and all electrons
in singly occupied orbitals
must have the same spin.
Ways to Represent Electron Configuration
1.Expanded Electron Configuration
2.Condensed Electron Configurations
3.Orbital Notation
4.Electron Dot Structure
Write the above four electron configurations for Zinc, Zinc
ion and Cu ion.