Periodic Table

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Periodic Table Chemistry
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ATOMIC STRUCTURE
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ELECTROMAGNETIC RADIATION
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Electromagnetic Radiation

• Most subatomic particles behave as PARTICLES and
  obey the physics of waves.

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Electromagnetic Radiation
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Electromagnetic Radiation

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Electromagnetic Radiation

• Waves have a frequency
• Use f for frequency, and units are cycles per
  sec
• All radiation f x ? c where c velocity
  of light 3.00 x 108 m/s
• Long wavelength --gt small frequency
• Short wavelength --gt high frequency

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Electromagnetic Spectrum

• Long wavelength --gt small frequency
• Short wavelength --gt high frequency

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Electromagnetic Radiation

• Red light has ? 700 nm. Calculate the
  frequency.

http//colossus.chem.umass.edu/bvining/downloads/c
hemland2/ElectroMagneticSpectrum.htm
http//science.widener.edu/svb/tutorial/wavefreqcs
n7.html
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Electromagnetic Radiation
Short wavelength --gt high frequency high
energy

• Long wavelength --gt
• small frequency
• low energy

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Electromagnetic Spectrum
http//colossus.chem.umass.edu/bvining/downloads/c
hemland2/ElectroMagneticSpectrum.htm
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Quantization of Energy
Max Planck (1858-1947) Solved the ultraviolet
catastrophe

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Quantization of Energy

• An object can gain or lose energy by absorbing or
  emitting radiant energy in QUANTA.

• Energy of radiation is proportional to frequency

E h f
h Plancks constant 6.6262 x 10-34 Js
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Quantization of Energy
E h f
Light with large ? (small f) has a small E.
Light with a short ? (large f) has a large E.
http//colossus.chem.umass.edu/bvining/downloads/c
hemland2/ElectroMagneticSpectrum.htm
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Photoelectric Effect
Experiment demonstrates the particle nature of
light.

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Photoelectric Effect

• Classical theory said that E of ejected electron
  should increase with increase in light
  intensitynot observed!
• No e- observed until light of a certain minimum E
  is used.
• Number of e- ejected depends on light intensity.

A. Einstein (1879-1955)
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Photoelectric Effect

• Understand experimental observations if light
  consists of particles called PHOTONS of discrete
  energy.

PROBLEM Calculate the energy of 1.00 mol of
photons of red light. ? 700. nm 4.29 x
1014 sec-1

shockwave
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Energy of Radiation

• Energy of 1.00 mol of photons of red light.
• E h?
• (6.63 x 10-34 Js)(4.29 x 1014 sec-1)
• 2.85 x 10-19 J per photon
• E per mol
• (2.85 x 10-19 J/ph)(6.02 x 1023 ph/mol)
• 171.6 kJ/mol
• This is in the range of energies that can break
  bonds.

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Excited Gases Atomic Structure
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Atomic Line Emission Spectra and Niels Bohr

• Bohrs greatest contribution to science was in
  building a simple model of the atom. It was based
  on an understanding of the SHARP LINE EMISSION
  SPECTRA of excited atoms.

Niels Bohr (1885-1962)
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Spectrum of White Light
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Line Emission Spectra of Excited Atoms

• Excited atoms emit light of only certain
  wavelengths
• The wavelengths of emitted light depend on the
  element.

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Spectrum of Excited Hydrogen Gas

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Line Emission Spectra of Excited Atoms
High E Short ? High f
Low E Long ? Low f

• Visible lines in H atom spectrum are called the
  BALMER series.

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Line Spectra of Other Elements
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The Electric Pickle

• Excited atoms can emit light.
• Here the solution in a pickle is excited
  electrically. The Na ions in the pickle juice
  give off light characteristic of that element.

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Atomic Spectra and Bohr
One view of atomic structure in early 20th
century was that an electron (e-) traveled about
the nucleus in an orbit.

• 1. Any orbit should be possible and so is any
  energy.
• 2. But a charged particle moving in an electric
  field should emit energy.
• End result should be destruction!

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Atomic Spectra and Bohr

• Bohr said classical view is wrong.
• Need a new theory now called QUANTUM or WAVE
  MECHANICS.
• e- can only exist in certain discrete orbits
  called stationary states.
• e- is restricted to QUANTIZED energy states
  each energy level is called a QUANTA.

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Atomic Spectra and Bohr

• Only orbits where n integral number of orbits
  are permitted.
• Each energy level can be mathematically
  represented.
• Results can be used to explain atomic spectra.
• However, these equations only worked well with
  the first few elements on the periodic table.

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Origin of Line Spectra
Simulation The hydrogen atom
Balmers equation
E -R(1)2 (1)2 ny nx
Balmer series
R is the Rydberg constant
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Atomic Line Spectra and Niels Bohr

• Bohrs theory was a great accomplishment.
• Recd Nobel Prize, 1922
• Problems with theory
• theory only successful for H.
• introduced quantum idea artificially.
• So, we go on to QUANTUM or WAVE MECHANICS

Niels Bohr (1885-1962)
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Quantum or Wave Mechanics

• de Broglie (1924) proposed that all moving
  objects have wave properties.
• For light E mc2 E h? hc / ?
• Therefore, mc h / ?
• and for particles
• (mass)(velocity) h / ?

L. de Broglie (1892-1987)
l h / m v

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Quantum or Wave Mechanics

• Schrodinger applied idea of e- behaving as a wave
  to the problem of electrons in atoms.
• He developed the WAVE EQUATION
• Solution gives set of math expressions called
  WAVE FUNCTIONS, ?
• Each describes an allowed energy state of an e-

E. Schrodinger 1887-1961
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WAVE FUNCTIONS, ?

• ??is a function of distance and two angles.
• Each ? corresponds to an ORBITAL the region
  of space within which an electron is found.
• ? does NOT describe the exact location of the
  electron.
• ?2 is proportional to the probability of finding
  an e- at a given point.
• KEY POINT Schroedingers equations are
  mathematical equations that describe the
  probability of finding an electron. These
  equations can be graphically represented on x-y-z
  axes.

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Uncertainty Principle

• Problem of defining nature of electrons in atoms
  solved by W. Heisenberg.
• Cannot simultaneously define the position and
  momentum ( mv) of an electron.
• We define e- energy exactly but accept limitation
  that we do not know exact position.

W. Heisenberg 1901-1976
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Types of Orbitals
s orbital
p orbital
d orbital
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s Orbitals
All s orbitals are spherical in shape.
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1s Orbital
1 s
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2s Orbital
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3s Orbital
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p Orbitals
There is a PLANAR NODE thru the nucleus.
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p Orbitals

• The three p orbitals lie 90o apart in space

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2px Orbital
3px Orbital
2 px
2 py
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d Orbitals

• When n 3, there are 3 subshells in the shell.
• s subshell with single orbital
• p subshell with 3 orbitals
• d subshell with 5 orbitals

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d Orbitals

• s orbitals have no planar node and so are
  spherical.
• p orbitals have 1 planar node, and so are
  dumbbell shaped.
• This means d orbitals have
• 2 planar nodes

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3dxy Orbital
3 dxy
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3dxz Orbital
3 dxz
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3dyz Orbital
3 dyz
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3dx2- y2 Orbital
3 d x2 - y2
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3dz2 Orbital
3 dz2
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f Orbitals

• When n 4,
• ---gt s subshell with single orbital
• ---gt p subshell with 3 orbitals
• ---gt d subshell with 5 orbitals
• ---gt f subshell with 7 orbitals