ATOMIC STRUCTURE

1
ATOMIC STRUCTURE

• A quick history of the development of atomic
  theory

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Democritus (460-370 B.C.)

• 1st to propose that all matter was composed of
  tiny individual particles he called atomos.
• Atomos (atoms) could not be created, destroyed or
  futher divided
• Not widely accepted due to more influential
  philosophers of his day (Aristotle) who
  challenged him to prove it.

3
John Dalton (1766-1844)

• Daltons Atomic Theory was proposed in 1803
• All matter is composed of extremely small
  particles called atoms
• All atoms of an element are identical having the
  same size, mass, and chemical properties. Atoms
  of one element are different from atoms of any
  other element.

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Daltons Atomic Theory (cont)

• Atoms cannot be created, divided into smaller
  particles, or destroyed.
• Different atoms combine in simple whole number
  ratios to form compounds.
• In chemical reactions, atoms are separated,
  combined, or rearranged.

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The Electron

• Study developed from the discovery of negatively
  charged particles with the use of a cathode ray
  tube.
• All matter showed evidence of having these
  negative particles
• J.J. Thomson (1856-1940) is credited with the
  discovery because his experiments showed that
  this mysterious negative particle was MUCH, MUCH
  smaller than an atom

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The Electron

• Robert Millikan (1868-1953) provided important
  details such as the exact charge of the electron
  (Millikans Oil Drop Experiment).
• This helped to calculate the exact mass of the
  electron
• Masselectron 9.1 x 10-28 g
• Thomson proposes the plum pudding model of the
  atom. (i.e. negatively charged electrons were
  distributed throughout a uniform positive charge.

7
The Nuclear Atom

• Ernest Rutherford (1871-1937) discovers a highly
  dense positive center of the atom using his
  famous experiment using gold foil and alpha
  particles in 1911 (diagram on p.95).
• But studies showed the atom was too heavy.

8
The Neutron

• James Chadwick discovered the neutron in 1932
• The neutron has a mass very similar to that of
  the proton, but carries no electrical charge.

Chart found on p.97 of text
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The Elements

• Matter as we know it is made up of atoms of
  different elements
• There are over 110 known elements
• Each element is defined by how many protons are
  found in its atoms nucleus. This is called the
  elements atomic number
• In a neutrally charged atom
• of protons of electrons
• The mass number of an atom is
• mass number protons neutrons

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Atomic Mass

• Atomic masses of elements are measured in Atomic
  Mass Units (amu)
• The amu is defined as exactly as the mass of the
  carbon-12 atom
• What we see in the periodic table is the average
  atomic mass of an element

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Isotopes

• Not all atoms of an element are identical!
• An isotope is an atom of an element that has the
  same number of protons but a different number of
  neutrons
• The average atomic mass of an element is found
  from adding up the mass numbers of each isotope
  with regards to their natural abundance.

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Atomic Mass Calculation

• Mass contribution of the isotope
• mass contribution
• (atomic mass of isotope)(natural abundance)
• Add up the mass contributions of each isotope
• avg atomic mass
• mass contribution 1 mass contribution 2 ..