Atomic Structure

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Chapter 5
Atomic Structure Atoms and Ions
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Michael Faraday

• Faraday discovered that ionic substances, such as
  NaCl would not conduct electricity in the solid
  state, but would conduct electricity when
  dissolved in water or molten (melted).
• He proposed that atoms could become charged and
  move toward oppositely charged electrodes
• He, and others, proposed that there may be a
  fundamental particle of electricity. Stoney
  called it the electron.

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Crookes Tube

• William Crookes, in the late 1800s, discovered
  that if electrodes, connected to a high voltage
  power supply, were separated and placed in a
  vacuum tube, the tube would glow with a
  yellow-green light.
• It was discovered that the light was a ray that
  came from the cathode (negative electrode) and
  traveled to the anode (positive electrode).
• The back end of the tube could be painted with a
  luminous pigment and the path of the ray could be
  studied.
• The important work done with Crookes tube was
  performed later by J.J. Thomson.

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Crookes Tube(Cathode Ray Tube)
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Crookes Tube (Cont.)
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Crookes Tube (Cont.)(Detecting positive
particles)
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Properties of Electrons

• The ray that Thomson studied was a beam of
  electrons. He was able to determine that the
  electron
• Was negatively charged
• Was affected by a magnetic field
• Had a mass/charge ratio that he measured and
  calculated
• Robert Millikan (an American) used an oil-drop
  experiment to measure the charge on oil droplets
  falling between charged plates. The charge on
  the droplets was a whole number multiple of the
  charge on the electron.

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Millikans Oil Drop Experiment for Determining
the Charge on the Electron
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Calculation of the Properties of the Electron

• Thomson found that the mass to charge ratio of
  the electron was -5.686x10-12 kg/C
• Millikan found the charge on the electron was
  -1.62 x 10-19 C. (C is the abbreviation for the
  coulomb, a unit of charge.)
• Therefore the mass of the electron can be
  calculated mass charge x mass/charge
• mass (-1.602x10-19C)(-5.686x10-12 kg/C)
• 9.1 x 10-31 kg 9.1 x 10-28 g

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How are negative and positive charges arranged in
the atom?
Thomson proposed a model of a spherical atom
composed of diffuse, positively charged matter or
field, in which electrons were embedded like
plum pudding.
Thomsons Plum Pudding Model of the Atom
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X-Rays

• Roentgen discovered X-rays which could penetrate
  opaque materials.
• He discovered X-rays while working with a cathode
  ray tube.
• These rays seemed to be emitted from the Cathode
  Ray tube itself and could be detected in another
  room.

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Radioactivity

• Becquerel discovered radioactivity.
• He left some uranium crystals on top of some
  photographic film which was covered with opaque,
  black paper. When he developed the film, he
  observed spots due to rays coming from the
  uranium crystals. The rays had penetrated the
  paper!

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Radioactivity and the Curies

• Pierre and Marie Curie
• Marie named the effect observed by Becquerel
  radioactivity.
• Radioactivity is the spontaneous emission of
  radiation from certain unstable elements.
• Pierre and Marie discovered radium and polonium,
  radioactive elements.
• They discovered much about radioactivity and
  radioactive elements

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Radioactivity (continued)

• Marie and Pierre Curie and Becquerel shared the
  Nobel Prize in Physics in 1903.
• Marie won the Nobel Prize in Chemistry in 1911.
• It is generally believed that both Curies died as
  a consequence of radiation poisoning
• Their daughter, Irene, also won a Nobel prize in
  1935

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Properties of alpha (a), beta (b), and gamma (g)
rays
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Method for Studying Emissions
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Rutherford set out to test Thomsons hypothesis
He bombarded gold foil with a particles. If
Thomsons plum pudding hypothesis were correct,
the a particles would be expected to be deflected
only to a small extent, if at all, because they
should act as dense, positively charged bullets
and go right through the gold atoms. The
embedded electrons could not deflect the a
particles any more than a bowling ball would be
deflected by ping pong balls.
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Rutherfords Alpha Scattering Experiment used to
Discover the Nucleus
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A Surprising Result Was Observed

• Although most of the particles went straight
  through the gold foil, a few of them were
  deflected by the foil at various angles. In fact
  some of the a particles bounced right back at the
  source.
• Thomsons plum pudding model did not explain
  this. What model would explain it?

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Rutherfords nuclear model of the atom.

• Atoms consisted of a central nucleus which had a
  positive charge and which had a very small
  volume, but it also contained most of the mass of
  the atom. Surrounding the nucleus were
  electrons, which had very little mass, but which
  occupied most of the volume of the atom.
• What was in the nucleus?

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Atomic Interpretation of the Alpha Particle
Scattering Experiment
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We already knew the atom contains electron(s)

• Goldstein discovered a positively charged
  particle that had a charge equal to the electron,
  but of opposite sign. It had a mass of 1 amu
  (1837 times the mass of the electron. This
  particle is called the proton.
• Rutherford concluded that the nucleus contained
  protons. He could account for the charge of the
  nucleus, but the mass of was too large for the
  number of protons.

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In 1932, Chadwick discovered a second nuclear
particle, the neutron

• Protons and neutrons make up most of the mass of
  the atom and are in the nucleus.
• Electrons are very light and are flying around
  outside the nucleus.

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Rutherfords Nuclear Model of the Atom

• Despite the success of Rutherfords model at
  explaining much of what was known about atomic
  structure, there were problems.
• The biggest problem was an apparent violation of
  the laws of physics. A charged particle, when
  accelerated, was known to emit electromagnetic
  radiation. However, electrons, according to
  Rutherford where orbiting in circular
  (accelerating) orbits around the nucleus and did
  not emit electromagnetic radiation.
• It was apparent that a more sophisticated model
  was needed.

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Electromagnetic Radiation

• Before we can explore our model of the atom
  further, we need to look more closely at energy
• Chemistry is the study of matter and energy. One
  type of energy is electromagnetic radiation. Let
  us look more closely at the properties of
  electromagnetic waves. Electromagnetic waves
  consist of oscillating, perpendicular electric
  and magnetic fields.
• The wavelength of radiation is the distance
  between peaks in a wave. (?)
• The frequency is the number of peaks that pass a
  point in a second. (? )

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Wavelength of Light
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A Simple Frequency and Wavelength Formula

• ln c
• l c/n
• n c/l
• l is wavelength measured in length units (m, cm,
  nm, etc.)
• n is frequency measured in Hz (s-1).
• c is the velocity of light in vacuum
• 3.0 x108 ms-1

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Electromagnetic Spectrum
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Electromagnetic Waves

• Describe electromagnetic radiation and give
  examples of it in relation to the electromagnetic
  spectrum.
• Type l (nm) n (Hz)
• radio (Rf) 108 - 1012 104-109
• microwave 106-108 109-1012
• infrared (IR) 750-106 1012-1014
• visible (vis) 400-750 1014-1015
• ultraviolet (UV) 10-400 1015-1016
• X-rays, g rays 10-4-1 1016-1022

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The Electromagnetic Spectrum
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Light Quanta and Photons

• Quantum- A packet of energy equal to hn. The
  smallest quantity of energy that can be emitted
  or absorbed.
• Photon- A quantum of electromagnetic radiation.
• Thus light can be described as a particle
  (photon) or as a wave with wavelength and
  frequency. This is called wave-particle duality
  (one of the most profound mysteries of science)

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Elemental Line Spectra
When certain elements are heated or
electronically excited, they emit light of
different colors. When the light is separated
into various colors by a spectroscope, a line
spectrum is observed.
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Bright Line Emission Spectrum from Excited Element
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Emission Spectrum

• Explain the emission line spectrum of light,
  based on the Bohr model of the hydrogen atom.
• Bohr explained the line spectrum by asserting
  that the electrons in the atoms could be in
  certain quantized energy levels.
• The spectrum arose due to transitions between
  quantized energy levels. The energy of the
  emitted light was equal to the difference in
  energies of the levels.

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Bohrs explanation of line spectra

• Bohr explained the line spectrum by asserting
  that the electrons in the atoms could be in
  certain quantized energy levels.
• The spectrum arose due to transitions between
  quantized energy levels. The energy of the
  emitted light was equal to the difference in
  energies of the levels.
• Electrons in atoms can not have any energy. They
  can only have certain amounts of energy. The
  electrons are said to be quantized. The emission
  (bright) line spectra are produced when electrons
  fall from a high energy level (excited state) to
  a lower energy level.

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Energy of emission lines
Ephoton DE Ehigh-Elow DE hn hc/l nlight
DE/h l hc/DE
Since Ehigh and Elow are discreet numbers, DE
must be a discreet number. Therefore, n and l
must be discreet numbers, giving rise to single
frequencies and wavelengths of light. Hence, the
line spectra.
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Emission Lines
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The Bohr Atom

• Bohr was able to accurately predict the energy
  levels of the one-electron atom, hydrogen.
• He suggested that multi-electron atoms would have
  electrons placed in the energy levels described
  by his theory.
• A certain maximum number of electrons could be in
  each level.

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Electrons in Energy Levels

• The maximum number of electrons in any energy
  level is 2n2
• Level 2n2 maximum number of
  electrons
• 1 2(1)2 2
• 2 2(2)2 8
• 3 2(3)2 18
• 4 2(4)2 32

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Bohr Diagrams

• Illustrations of electrons in energy levels are
  called Bohr Diagrams.
• The electrons in the outer levels are called
  valence electrons.
• The valence electrons are those involved in
  chemical bonding.
• Examples of Bohr diagrams are shown on the next
  slide

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Bohr Diagrams
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Valence Electrons

• The outer electrons in an atom can be represented
  with dots in the Lewis electron dot symbol. Each
  outer electron is represented by a dot around the
  atomic symbol
• Sodium has one valence electron, hence one dot
• Na
• Sodium ion has lost its valence electron, no
  dots
• Na

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Lewis Symbols
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Lewis Dot Structures for Main Group Elements are
Determined form the Group Number

• Group IA elements (alkali metals) have 1 valence
  electron 1 dot Na
• Group IIA elements (alkaline earths) have 2
  valence electrons 2 dots Ca
• Group VIII elements (noble gases) have eight
  dots, an octet Ne
• He, has only 2 electrons, 2 dots He

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Bohrs Model was improved upon in the 1920s with
the Quantum Mechanical Model.

• Since Bohrs model only worked for the hydrogen
  atom, a more sophisticated model was needed.
• The next breakthrough was made by Louis de
  Broglie, who suggested that electrons, like
  photons have wave properties
• De Broglie thought that Bohrs energy levels were
  created by the wave properties of the electron

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De Broglie suggested an electron could only have
a path that allowed a whole number of wave
patterns
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Other Contributors to the Quantum Mechanical Atom

• Schrödinger used de Broglies ideas to create
  some powerful wave equations to describe the
  electron.
• Heisenberg used probability and matrices to
  describe the electron. He stated a controversial
  Uncertainty Principle. The path of an electron
  can not be determined. It is uncertain. Thus a
  specific orbit for an electron can not be known.

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Evolving Theories of the Atom
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In addition to the energy levels of Bohr, there
are sub-levels

• Bohrs energy levels were assigned a principal
  quantum number, n, which could values of 1, 2, 3
  This quantum number, n designates the energy
  level and size of the region in space the
  electrons might be found.
• Within an energy level there are sublevels or
  subshells, designated s, p, d, and f. These
  subshell designations tell the shape of the
  region in space the electrons might be found.

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Charge Cloud Representations of s Orbitals
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Shapes of p Orbitals
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px, py, and pz Orbitals
A p subshell contains 3 p orbitals, each lies
perpendicular to the others on the X-Y-Z axes
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s, p, d, and f Orbitals
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Orbitals

• Each orbital has its own set of quantum numbers.
  
• Each orbital can contain 2 electrons, one with
  spin 1/2 the other with spin -1/2
• The quantum number and energy levels can be
  described with an orbital diagram.
• A summary of an orbital diagram is called an
  electron configuration.

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Energies of Orbitals in Multi-Electron Atoms

• Several factors affect the energy of electrons in
  multi electron atoms
• Nuclear charge
• Electron repulsions
• Additional electrons in the same orbital
  (shielding)
• Additional electrons in inner orbitals
• Orbital shape (ml)
• spin (ms)
• Pauli Exclusion Principle No two electrons in
  the same atom can have the same set of four
  quantum numbers.

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Aufbau Principle

• Electrons arrangements are built-up by filling
  various energy levels, starting with lower
  energies, filling orbitals two at a time. They
  go into the orbitals one with spin 1/2, the
  other with spin - 1/2.
• An orbital diagram is useful in showing this
  arrangement.

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Orbital Diagrams
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Electronic Configurations
By adding electrons to the diagram, lowest energy
to highest, remembering Hunds rule and the
quantum rule that no orbital can hold more than
two electrons, an elcetronic configuration can be
created
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Electron Configurations can be Determined From
the Position in the Periodic Table

• Elements in group 1(1A) end in ns1.
• Elements in group 2 (2A) end in ns2
• Elements in group 13 (3A) end in ns2np1
• Elements in group 14 (4A) end in ns2np2
• Elements in group 15 (5A) end in ns2np3
• Elements in group 16 (6A) end in ns2np4
• Elements in group 17 (7A) end in ns2np5
• Elements in group 18 (8A) end in ns2np6

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Periodic Table Family Filling Diagram
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Periodic Table Orbital Block Diagram