Atomic Theory and Structure

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Atomic Theory and Structure
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John Dalton
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Leading to Discovery

• Series of meteorological observations 1787 in
  Manchester area
• Proved the validity of concept that rain is
  precipitated by a decrease in temperature
• Interest in meteorology led to study a variety of
  phenomena and instruments

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Discoveries leading up to model

• Careful quantitative measurements-allowed precise
  measurements of compounds
• Law of Definite ProportionsIn the early 1800s
  Dalton noted that oxygen and carbon combine to
  make two compounds. Each compound had twice as
  much oxygen as carbon
• This led him to propose the Law of Simple
  Multiple Proportions

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Daltons Atomic Theory

• In an attempt to explain how and why elements
  would combine with one another in fixed ratios
  and sometimes also in multiples of those ratios,
  Dalton formulated his atomic theory.
• Elements consisted of tiny particles called atoms
• Atoms of the same element had the same mass
  atoms of different elements had different masses
• Compounds consisted of atoms of different
  elements combined together
• He also stated that chemical reactions involved
  the rearrangement of combinations of those atoms

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Summary of Atomic Theory-1803

• All matter is made of indivisible and
  indestructible atoms
• All atoms of a given element are identical in
  their physical and chemical properties
• Atoms of different elements differ in their
  physical and chemical properties
• Atoms of different elements combine in simple
  whole-number ratios to form compounds
• Chemical reactions consist of the combination,
  separation, or rearrangement of atoms

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Daltons Model

• Dalton's model was that the atoms were tiny,
  indivisible, indestructible particles and that
  each one had a certain mass, size, and chemical
  behavior that was determined by what kind of
  element they were

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Daltons Model
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Problems with Theory

• assumed all gases were monoatomic, e.g. oxygen is
  O
• assumed simplest compounds were binary, e.g.
  water is HO
• his atomic weights were approximate
• no knowledge of isotopes

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The Thomson Model
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Discoveries Leading to the Model

• Thomson Followed William Crookes experiments
  that proved that the rays in the magnetic field
  bent

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Procedure

• voltage applied to metal at each end of tube
  (electrodes)
• Anode
• Cathode
• Magnet deflected
• Beammust contain
• Something negative
  

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What is it?
Discovered by J.J. Thomson
In the year 1897
An Electron!
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Theory and Model
Electron
Nicknamed the Plum Pudding model, it shows
electrons embedded in a positively charged ball
of matter.
Sphere of Positive Charge
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Picture of Thomsons Model
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Rutherford-Bohr Model
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Discoveries Leading up to the Model

• The scientists knew that the electron had a
  negative electrical charge. They also knew that
  atoms had no overall charge.

• Scientists guessed that since electrons are
  extremely small, whatever this positive something
  was, it must be make up most of the mass of
  atoms, and be much larger

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Rutherford

• Since the Thomson model, Ernest Rutherford had
  discovered the nucleus of an atom in his gold
  foil experiment
• He conducted experiments in which he shot large,
  charged particles (alpha particles) at a thin
  gold foil
• Most particles passed through the foil, but some
  came off at odd angles (concluded an atom must
  contain a central mass that would deflect these
  particles)

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The Model Itself

• A New Zealand scientist, Ernest Rutherford, and a
  Danish scientist, Niels Bohr, developed a way of
  thinking about the structure of an atom in which
  an atom looks very much like our solar system.
  (electrons revolve around nucleus like planets
  around the Sun)
• But Rutherford could not explain why the
  negatively charged electrons were not pulled into
  the positively charged nucleus as Coulombs law
  would suggest.

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Bohr

• Bohr had corrected a serious flaw by recognizing
  that electrons had to be in orbits (energy
  states).
• But his analysis of the energy given off when an
  electron dropped from a higher energy orbit to a
  lower energy orbit didnt hold up for atoms
  bigger than hydrogen (the simplest atom, with
  only one proton and no neutrons)
• More work needed to be done with the model

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Rutherford-Bohr Model

• -Ernest Rutherford-

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Despite its technical flaws, however, the
Rutherford-Bohr model is still useful because it
is simple and helps people understand atomic
structure
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The Quantum Theory
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Discoveries leading up to

• The Rutherford model describes electrons in terms
  of there energy state
• Bohr postulated that electrons did not radiate
  energy while in orbit around the nucleus, but
  Bohrs model could not explain the spectra of
  larger atoms
• In 1924, a French physicist named Louis de
  Broglie suggested that electrons could act as
  both particles and waves.
• An Austrian physicist named Erwin Schrodinger
  derived a set of equations or wave functions in
  1926 for electrons
• According to Schrodinger, electrons confined in
  their orbits would set up standing waves and you
  could describe only the probability of where an
  electron could be. The distributions of these
  probabilities formed regions of space about the
  nucleus were called orbitals.

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The Quantum Theory

• The present day quantum model postulates that
  electrons have the properties of both particles
  and waves.

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Electrons n More

• The wave-particle duality of quantum theory
  allows electrons to be described as waves, using
  the electrons de Broglie wavelength.

Louis de Broglie
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Quantum Atom (Cont.)

• Although the position of an electron is
  uncertain, quantum theory does not allow the
  electron from being at some places. The easiest
  way to describe the differences between the
  allowed and prohibited positions of electrons in
  an atom is to think of the electron as a wave.

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Electron Cloud

• In the quantum theory, electrons are located in
  orbitals, which are regions of space in which you
  can expect to find electrons of specific energy.
• Despite the similar name, and orbital is
  different from and orbit. An orbital is a region
  of high probability for finding a particular
  electron. It is as if the electron were smeared
  into a cloud.
• We use the quantum numbers to find where each
  electron is, but certain combinations of n, l,
  and m are not always permitted.

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Quantum model
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Additional concepts

• Pauli exclusion principle- no more than two
  electron can occupy an orbital, and they would
  spin in opposite directions
• Aufbau principle- electrons will fill lower
  energy positions first (build up from there

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• Hunds rule-orbitals of the same energy will half
  fill each of those orbitals first, and then go
  back and fill those orbitals

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Quantum Numbers

• Set of 4 numbers that describe the address of a
  particular electron
• First is the Principle Quantum- states the main
  energy level n 1,2,3,4,,,
• Second is the Azimuthal (or you can think shape
  l (0n-1)

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• 3rd is the magnetic quantum
• m (-l0l)
• 4th is the magnetic spin
• ms 1/2 or -1/2

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5f
7s
6p
5d
4f
6s
5p
4d
5s
4p
3d
4s
3p
3s
2p
2s
1s
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