Atomic Structure and Periodicity

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Atomic Structure and Periodicity

• By Ms. Buroker

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Electromagnetic radiation

• Energy that moves through space in a wavelike
  manner.
• Electromagnetic radiation is self-propagating
  (i.e. it doesnt require a medium to travel
  through. It can travel through the vacuum of
  space.)

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Electromagnetic Radiation
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Electromagnetic Spectrum
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Example Problems
1.) Which color in the visible spectrum has the
highest frequency? Which has the lowest
frequency? 2.) Is the frequency of the
radiation used in a microwave oven higher or
lower than that from a FM radio station
broadcasting at 91.7MHz? 3.) Is the wavelength
of x-rays longer or shorter than that of
ultraviolet light?
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Example Problem

• The brilliant red colors seen in fireworks are
  due to the emission of light with wavelengths
  around 650nm when strontium salts such as
  Sr(NO3)2 and SrCO3 are heated. Calculate the
  frequency of red light of wavelength
• 6.50 x 102nm.

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Traditional Theories Intensity of
radiation should increase continuously
with decreasing wavelength
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A Quandary

• Are matter and energy different?
• Matter was thought to consist of particles,
  whereas energy in the form of light
  (electromagnetic radiation) was described as a
  wave. Particles were things that had mass and
  whose position in space could be specified.
  Waves were described as massless and delocalized
  that is, their position in space could not be
  specified, It also was assumed that there was no
  intermingling of the matter and light.

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Max Plank (1858- 1947)

Lead to the discovery of a quantum which is
like a packet of energy released when there is a
transfer.
Vibrations in atoms are quantized (i.e. Only
certain vibrations with certain frequencies are
allowed)
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The Photoelectric Effect
Assumption Energy is carried on the amplitude of
a wave (as in classical waves). Prediction
Since the amplitude of an EM wave correlates with
the brightness of the light, light of high enough
intensity irradiating a metal for a long enough
period of time should be able to eject electrons
from the surface of a metal. Reality It was the
frequency of the light that determined whether or
not electrons were ejected regardless of the time
of irradiation. The brightness (amplitude) only
determined how many were ejected per unit time
once ejection started occurring.
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Example Problem

• The blue color in fireworks is often achieved
  by heating copper(I) chloride (CuCl) to about
  1200?C. Then the compound emits blue light
  having a wavelength of 450nm. What is the
  increment of energy (the quantum) that is emitted
  at 4.50 x 102nm by CuCl?

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Einstein and the photon
Einstein (1905) incorporated Plancks equation
with the idea that light had (mass-less) particle
properties (photon). These photons are packets
of energy where E depends on the frequency of the
photon (E h?).
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Einstein and the Photon

• He proposed that electromagnetic radiation is
  itself is quantized and can be viewed as a stream
  of particles called photons.
• Ephoton hn hc
• l

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Example problem
A 60W, monochromatic laser beam gives off photons
of wavelength 650nm. How long does it take for
2.5moles of these photons to be given
off? Answer E hc/? (6.626x10-34Js)(3.0x108m
/s)/6.50x10-7m 3.06x10-19J/photon (3.06x10-19J/
photon)(6.022x1023photons/mol)(2.5mol)
4.60x105J 4.60x105J / (60J/s) 7673.4s
/(3600s/1hr) 2.13h
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Example Problem
Compare the energy of a mole of photons of orange
light (625nm) with the energy of a mole of
photons of microwave radiation having a frequency
of 2.45GHz (1GHz 109s-1). Which has the greater
energy? By what factor is one greater than the
other? Answer E (6.626x10-34Js)(3.00x108m/s)/(
6.25x10-7m) 3.18x10-19J 3.18x10-19J(6.022x1023)
1.92x105J/mol E (6.626x10-34Js)(2.45x109Hz)
1.623x10-24J 1.623x10-24(6.022x1023)
9.78x10-1J/mol 1.9x105 / 9.7x10-1 1.96x105
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Atomic Line Spectra
Atomic Emission Spectra
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Atomic emission spectra
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Bohr Model of the atom
First connection between line spectra and quantum
ideas of Planck and Einstein. Bohr Model
Electrons orbit the nucleus of the atom like
planets going around the sun. Only certain
stable orbits are allowed (quantized) to keep
the electron (a charged, accelerating particle)
from crashing into the nucleus.
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n is an integer equal to or greater than 1
and Rhc 2.179x10-18J/atom or 1312kJ/mol
Principle quantum number Note Potential
energy 0 at infinity
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Electron transition diagram
1) Electron in ground state 2) Electron jumps to
excited state from absorption of outside energy
(Energy absorbed positive) 3) Electron
transitions back down giving off photon that is
equal in energy to the transition down (energy
emitted negative)
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Lyman (UV), Balmer (Visible), and Paschen
(IR), series of the hydrogen atom.
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Example Problem
Calculate the energy of the n3 state of the H
atom in a) joules per atom and b) kilojoules per
mole. Rhc 2.179x10-18J/atom or
1312kJ/mol Answer E -Rhc/n2
-2.179x10-18J/atom / 32 -2.421x10-19J E
-1312kJ/mol / 32 -145.8kJ/mol
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Quantum Numbers
Each electron in an atom has a unique set of 4
quantum numbers which describe it.

• Principal quantum number
• Angular momentum quantum number
• Magnetic quantum number
• Spin quantum number

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Pauli Exclusion Principle
No two electrons in an atom can have the same
four quantum numbers.
Wolfgang Pauli
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Principal Quantum Number
Generally symbolized by n, it denotes the shell
(energy level) in which the electron is located.
Number of electrons that can fit in a shell
2n2
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Angular Momentum Quantum Number
The angular momentum quantum number, generally
symbolized by l, denotes the orbital (subshell)
in which the electron is located.
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Magnetic Quantum Number
The magnetic quantum number, generally symbolized
by ml, denotes the orientation of the electrons
orbital with respect to the three axes in space.
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Assigning the Numbers

• The three quantum numbers (n, l, and ml) are
  integers.
• The principal quantum number (n) cannot be zero
  ? 1, 2, 3, etc.
• The angular momentum quantum number (l ) can be
  any integer between 0 and n - 1. For n 3, l
  can be either 0, 1, or 2.
• The magnetic quantum number (ml) can be any
  integer between -l and l.
• For l 2, m can be either -2, -1, 0, 1, 2.

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Spin Quantum Number
Spin quantum number (ms) denotes the behavior
(direction of spin) of an electron within a
magnetic field.
Possibilities for electron spin
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An orbital is a region within an atom where
thereis a probability of finding an electron.
This is a probability diagram for the s orbital
in the first energy level
Orbital shapes are defined as the surface that
contains 90 of the total electron probability.
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Schrodinger Wave Equation
Equation for probability of a single electron
being found along a single axis (x-axis)
Erwin Schrodinger
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Heisenberg Uncertainty Principle
One cannot simultaneously determine both the
position and momentum of an electron.
You can find out where the electron is, but not
where it is going.
OR
You can find out where the electron is going, but
not where it is!
Werner Heisenberg
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Orbitals of the same shape (s, for instance) grow
larger as n increases
Nodes are regions of low probability within an
orbital.
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The s orbital has a spherical shape centered
around the origin of the three axes in space.
s orbital shape
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P orbital shape
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d shaped orbitals
Things get a bit more complicated with the five d
orbitals that are found in the d sublevels
beginning with n 3. To remember the shapes,
think of
double dumbells
and a dumbell with a donut!
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Shape of f orbitals
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Orbital filling table
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Irregular confirmations of Cr and Cu
Chromium steals a 4s electron to half fill its
3d sublevel
Copper steals a 4s electron to FILL its 3d
sublevel
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Determination of Atomic Radius
Half of the distance between nuclei in
covalently bonded diatomic molecule
"covalent atomic radii"
Periodic Trends in Atomic Radius

• Radius decreases across a period

Increased effective nuclear charge due to
decreased shielding

• Radius increases down a group

Addition of principal quantum levels
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Table of Atomic Radii
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Ionization Energy
Ionization Energy the energy required to remove
an electron from an atom

• Tends to increase across a period

Electrons in the same quantum level do not
shield as effectively as electrons in inner
levels
Irregularities at half filled and filled
sublevels due to extra repulsion of electrons
paired in orbitals, making them easier to
remove

• Tends to decrease down a group

Outer electrons are farther from the nucleus
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Ionization of Magnesium
Mg 738 kJ ? Mg e-
Mg 1451 kJ ? Mg2 e-
Mg2 7733 kJ ? Mg3 e-
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Table of 1st Ionization Energies
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Electron Affinity
Electron Affinity is the energy change associated
with the addition of an electron. So think the
opposite of ionization energy. 1.) Electron
affinity tends to increase across a period. 2.)
Affinity tends to decrease as you go down in a
group. Electrons farther from the nucleus
experience less nuclear attraction Some
irregularities due to repulsive forces in the
relatively small p orbitals
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Table of Electron Affinities
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Electronegativity
Electronegativity is a measure of the ability of
an atom in a chemical compound to attract
electrons. Trend . 1.) Electronegativity tends
to increase as you go across a period. 2.)
Electronegativity tends to decrease as you go
down a group or remain the same.
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Factors Determining Electronegativity
When you subtract the electronegativity values of
two atoms bound together you use the value to
determine what kind of bond you have. Non-polar
covalent 0-0.3 Polar Covalent Bonds 0.3- 1.7
Ionic Bonds 1.7- 3.3
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Periodic Table of Electronegativities
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Ionic Radii
Cations

• Positively charged ions

•   Smaller than the corresponding
• atom

Anions

•   Negatively charged ions

•   Larger than the corresponding
• atom

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Table of Ion Sizes
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Summary of Periodic Trends