Transcript EMR and the Bohr Model of the Atom

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Atomic Structure
and Periodicity
AP Chemistry
Chapter 7
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7.1 EMR
• Electromagnetic radiation – a form of
energy that exhibits wavelike behavior
as it travels through space.
• Types include visible light, X rays,
ultraviolet light, infrared light,
microwaves, and radio waves.
• Electromagnetic spectrum – All the
forms of electromagnetic radiation
together
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Electromagnetic Radiation
• Waves have a wavelength –
distance between corresponding
points on adjacent waves
• Use the Greek letter “lambda”,
, for wavelength, and units are
length units (m, cm, nm)
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Electromagnetic Radiation
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Electromagnetic Radiation
• Waves have a frequency –
number of waves that pass a
given point in a specific time
• Use the Greek letter “nu”, , for
frequency, and units are “cycles
per sec” or Hertz (Hz)
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Electromagnetic Radiation
• All radiation travels at the same
speed of light.
• c = 3.00 x 108 m/s
• = c
• This means that  must be in meters
and  must be in Hertz (1/s) so that
units cancel.
Electromagnetic Spectrum
Long wavelength  small frequency
Short wavelength  high frequency
increasing
frequency
increasing
wavelength
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Electromagnetic Spectrum
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Radio
waves
Microwaves
Infra-Red
Low frequency, Long wavelength
1 m - 1km
1 cm
0.01 mm
Visible light
400-700 nm
Ultra-Violet
100 nm
X-Rays
Gamma
Rays
1 nm
0.01 nm
High frequency, Short wavelength
Problems with Wave Theory
of Light
• Scientific belief around the
1900’s was that there was
NO relationship between
matter and light
• Light given off by objects
that were heated to high
temperatures could not be
explained.
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Black Body Radiation
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7.2 Nature of Matter
Max Planck
• Stated that objects radiated energy
in small packets of energy called
quanta
quantum- a specific amount of
energy that can be gained or lost by
an atom
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Particle Behavior of Light
• Energy and frequency are
directly related
E=h
• E is energy (J)
• h is Planck’s constant
• h = 6.626 x 10-34 J s
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Photoelectric Effect
Thomson (1839)
• First to observe the photoelectric
effect
• photoelectric effect - the
emission of electrons from a
metal surface when exposed to
light of a specific energy.
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Photoelectric Effect
1905 Albert Einstein
• stated that EMR could be viewed
as a stream of particles “photons”
• photon- a quantum of light
• energy of these photons could be
calculated by Planck’s equation
• stated that the photons strike the
electrons therefore ejecting them
from the metal
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Photoelectric Effect
Dual Wave-Particle
Behavior Of Light
• The success of Einstein’s work in
explaining the photoelectric effect
was largely responsible for the
acceptance of the particle behavior
of light
• Ephoton = h
• E = mc2
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Can matter act as a wave?
• Using Einstein’s and Planck’s
equations, de Broglie derived:
h

mv
• The momentum, mv, is a particle
property, whereas  is a wave
property.
• In one equation de Broglie
summarized the concepts of waves
and particles as they apply to low
mass, high speed objects.
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Sample Problem
• Compare the wavelength for an electron
(mass = 9.11 x 10-31 kg) traveling at a speed of
1.0 x 107 m/s with that for a ball (mass = 0.10
kg) traveling at 30 m/s.
Dual Wave-Particle
Behavior Of Matter
• Energy is a form of matter.
• All matter exhibits both particle and wave
properties.
• Large pieces of matter (i.e. baseball)
exhibits mostly particle properties.
• Tiny pieces of matter (i.e. photons) exhibits
mostly wave properties.
• Pieces of matter somewhere in the middle
(i.e. electrons) clearly show both types of
properties!
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Kirchoff and Robert
Bunsen (1854)
• Observed that light was given
off when they heated different
chemicals in their designed
burner
• They passed the light through
a prism and saw separate
lines instead of a continuous
spectrum.
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Absorption and Emission
Spectra
Emission spectra- the colors
produced by an object
when burned or heated.
Absorption spectra- the
colors that are not shown,
rather absorbed in the
spectrum
http://chemistry.beloit.edu/BlueLight/moviepages/ab_em_el.htm
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7.3 Hydrogen Spectrum
• Only four lines are emitted:
– Red, green, blue, violet
• Only certain energies are allowed.
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Why do elements
produce these lines?
To understand emission
spectrum, we must
understand these two terms:
Ground state: the lowest
energy state for the electron
Excited state: state where
electron has higher energy
than ground state.
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Why do elements
produce these lines?
Atoms are heated, which adds
energy. The electron become
excited (thus unstable). They
want to return to their normal,
or ground state. To do so,
they give off energy in the
form of EMR.
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Scientists associated
with the H spectrum
• Balmer: developed a numerical
relationship between the
wavelength of the lines in the
spectrum and the amount of
energy
• Lyman: discovered lines
produced in the UV range.
• Paschen: discovered lines
produced in the IR range
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7.4 The Bohr Model
1913 Neils Bohr
-worked with Rutherford to
study the H spectrum.
-Bohr’s model is sometimes
referred to as the “Planetary
model” based upon his
postulates.
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Bohr
Model
of
the
Atom
Bohr Model of the Atom
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Postulates of Bohr’s model:
1. The single electron of hydrogen can circle
the nucleus in fixed paths called orbits or
stationary states.
2. The electron can jump to higher orbits
when energy is added.
3. The angular momentum of the electron is
quantized.
-The electron’s energy can be calculated in
the different orbits.
Bohr
Model
of
the
Atom
Bohr Model of the Atom
How does this relate to the
Hydrogen spectrum?
• Bohr calculated the
energy that the electron
would lose as it fell from
higher orbits to lower
orbits.
• Bohr’s calculations
agreed exactly with
Lyman, Balmer and
Paschen’s observations.
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Bohr Model of the Atom
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Sample Problem
• Calculate the energy required to
excite the hydrogen electron from
n=1 to n=2.
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Sample Problem
• Calculate the energy required to completely
remove the electron from a hydrogen atom in
its ground state.
ninitial = 1 to nfinal = ∞
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Bohr
Model
of
the
Atom
Bohr Model of the Atom
• Bohr’s model worked very
well for the Hydrogen atom.
• Through Bohr’s work, as well
as the other scientists
mentioned, a very good
understanding of the electron
within the atom was now in
place.
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Downfalls to Bohr’s
Downfalls
to Bohr’s
Model
Model
of the
Atom
1. Bohr’s model of the atom worked very
well for the hydrogen atom and the
He+, but failed when applied to
multielectron atoms.
2. Bohr’s model could not explain why
the electron could not exist between
orbits.
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Now What?
• We need a new approach to
the atom!
• Big Three: de Broglie,
Heisenberg & Schrodinger
• Developed wave
mechanics AKA quantum
mechanics (7.5)
Heisenberg’s Uncertainty
Principle
With respect to atomic particles,
we cannot determine exactly
1. the position
2. direction of motion
AND
3. speed
simultaneously.
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Schrodinger’s Wave
Equation
• Schrödinger proposed an equation that
contains both wave and particle terms.
• Solving the equation leads to wave
functions .
• The wave function gives the probability
distribution of an electron.
• We call wave
functions orbitals.
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7.6 Quantum Numbers
• When we solve the Schrödinger
equation for the hydrogen atom, we
find many wave functions (orbitals)
that satisfy it.
• Each orbital is characterized by a
series of numbers called quantum
numbers that describe various
properties of the orbital.
Principal Quantum
Number, n
• Related to the size and energy of the
orbital – think energy level
• n has integer values: 1,2,3…
• As n becomes larger, the atom
becomes larger and the electron is
further from the nucleus.
• A larger n value also corresponds to
higher energy because the electron is
less tightly bound to the nucleus.
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Angular Momentum
Quantum Number, l
• Related to the shape of the atomic orbitals
• This quantum number depends on the
value of n. The values of l begin at 0 and
increase to (n - 1).
• Because we use numbers to describe the
first quantum number, we usually use
letters for l (s for l =0, p for l = 1, d for l =2
and f for l = 3).
• Usually we refer to the s, p, d and forbitals.
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Magnetic Quantum
number, ml
• Provides the 3D orientation of the
orbital in space
• Value depends on l. The magnetic
quantum number has integer
values between -l and +l.
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Quantum Numbers
7.7 Orbital Shapes &
Energies
• Each orbital has a unique probability
distribution.
• Nodes = areas of zero probability
• To simplify, we think of orbitals in terms of
their overall shapes, which becomes larger as
n increases.
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p orbitals
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d orbitals
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f orbitals
Energies of orbitals in
Hydrogen
• For hydrogen, energy is
determined by value of n
• All orbitals with the same value of
n have the same energy – they are
degenerate.
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7.8 Electron Spin & Pauli
Exclusion Principle
• Developed by Samuel Goudsmit & George
Uhlenbeck (University of Leyden in the
Netherlands)
• 4th quantum number necessary to account for
the details of emission spectra of atoms
• Electron has a magnetic moment with two
possible orientations when placed in an
external magnetic field.
• Magnetic spin quantum number ms can only
have two possible values +½ and -½
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7.8 Electron Spin & Pauli
Exclusion Principle
• Wolfgang Pauli developed Pauli
exclusion principle
• In a given atom, no two electrons
can have the same set of four
quantum numbers
• An orbital can hold only 2
electrons, and they must have
opposite spins
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7.9 Polyelectronic
Atoms
Effective Nuclear Charge
• Electrons are attracted to the nucleus, but repelled
by the electrons that screen it from the nuclear
charge.
• The nuclear charge experienced by an electron
depends on its distance from the nucleus and the
number of core electrons.
• As the average number of screening electrons (S)
increases, the effective nuclear charge (Zeff)
decreases.
• As the distance from the nucleus increases, S
increases and Zeff decreases.
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7.10 History of the
Periodic Table
• Dobereiner – triads (groups of 3 elements
share similar properties)
• Newlands – octaves (certain properties repeat
for every eighth element)
• Meyer & Mendeleev – present form of periodic
table
• Mendeleev – considered father of periodic table
because he predicted the existence and
properties of still unknown elements and left
space for them in his periodic table
• Fundamental difference – modern periodic
table organized by atomic number not mass
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7.11 Aufbau Principle
& the Periodic Table
• As protons are added one by one to the
nucleus to build up the elements, electrons
are similarly added
• Electron configurations tells us in which
orbitals the electrons for an element are
located.
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Periods 1 - 3
Three rules:
– electrons fill orbitals starting with lowest n
and moving upwards;
– no two electrons can fill one orbital with the
same spin (Pauli);
– for degenerate orbitals, electrons fill each
orbital singly before any orbital gets a
second electron (Hund’s rule).
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Period 4 and Beyond
• After Ca the d orbitals begin to fill.
• After the 3d orbitals are full the 4p orbitals
being to fill.
• From Ce onwards the 4f orbitals begin to fill.
• Note: La: [Xe]6s25d14f0
• Elements Ce - Lu have the 4f orbitals filled
and are called lanthanides.
• Elements Th - Lr have the 5f orbitals filled and
are called actinides.
• Most actinides are not found in nature.
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Electron Configurations
and the Periodic Table
• The periodic table can be used as a guide for
electron configurations.
• The period number is the value of n.
• Groups 1 and 2 have the s-orbital filled.
• Groups 13 - 18 have the p-orbital filled.
• Groups 3 - 12 have the d-orbital filled.
• The lanthanides and actinides have the f-orbital
filled.
• Note that the 3d orbital fills after the 4s orbital.
Similarly, the 4f orbital fills after the 6s orbital.
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Noble Gas Notation
• There is a shorthand way of writing
electron configurations
• Write the core electrons corresponding
to the filled Noble gas in square
brackets.
• Write the valence electrons explicitly.
• Example, P: 1s22s22p63s23p3
• but Ne is 1s22s22p6
• Therefore, P: [Ne]3s23p3.
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Practice Problem
Determine the expected electron
configurations for each of the
following:
• S
• Ba
• Ni2+
• Eu
• Ti+
Effective Nuclear
Charge - revisited
• Many properties of atoms depend on electron
configurations and how strongly valence
electrons are attracted to the nucleus.
• Coulomb’s Law – strength of the interaction
between 2 electrical charges depends on the
size of the charges and the distance between
them.
• Zeff = Z – S where Z is # protons in nucleus
and S is number of core electrons
• Explains differences in sublevel energies but
also describes periodic trends.
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Effective nuclear
charge
• The effective nuclear charge increases as we
move across any row (period) of the periodic
table (Z gets larger while S stays the same)
• The effective nuclear charge also increases
as we go down a column (group) of the
periodic table, but the effect is far less than
going across a row.
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Atomic Radius
Simple diatomic molecule
• The distance between
the two nuclei is called
the bond distance.
• If the two atoms which
make up the molecule
are the same, then half
the bond distance is
called the covalent
radius of the atom.
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Atomic Radius
• Atomic size varies consistently through
the periodic table.
• As we move down a group, the atoms
become larger.
• As we move across a period, atoms
become smaller.
• There are two factors at work:
– principal quantum number, n (down a group)
– the effective nuclear charge, Zeff (across a
period)
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Atomic Radius
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Ionization Energy
• Ionization energy – minimum amount of
energy required to remove an electron from
the ground state of an isolated gas atom or
ion.
• Na(g)  Na+(g) + e- First ionization energy
• Na+(g)  Na2+(g) + e- Second ionization
energy
• The greater ionization energy, the more
difficult it is to remove the electron.
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Ionization Energy
• Ionization energy increases for each
additional electron removed from an atom.
• There is a sharp increase in ionization energy
when a core (non-valence) electron is
removed.
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Ionization Energy Trend
• Same factors influence ionization energy –
effective nuclear charge & distance of electron
from nucleus.
• Increasing effective charge or decreasing
distance from nucleus increases attraction
between electron & nucleus – more difficult to
remove an electron so ionization energy
increases. (Both happen when move across row)
• As we move down group, the atomic radius
increases (due to larger n) while effective nuclear
charge only increases slightly. Attraction
between nucleus & electron decreases, so
ionization energy decreases.