The History of Our Model of the Atom

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The History of Our Model of the Atom
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The History of the Atomic Model

• Original idea Ancient Greece (400 B.C..)
• Democritus and Leucippus Greek philosophers
• Greeks did not experiment working with ones
  hands was for slaves
• Much of their ideas based upon logic and
  reasoning alone

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So why does Democritus think all things are made
of atoms?

• A thought game
• If you have a stick, you can break it in half.
• Each half can then be broken in two halves.
• This process can continue . . . Indefinitely?
• Democritus said no!!
• Eventually you reach something that cant be
  divided he called these atoms.

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Other thoughts in Greece

• Aristotle - Famous philosopher
• All substances are made of 4 elements
• Fire - Hot
• Air - light
• Earth - cool, heavy
• Water - wet
• Blend these in different proportions to get all
  substances

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Democritus vs. AristotleWho Wins?

• Aristotle was more famous
• He won
• His ideas carried through middle ages.

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Whos Next?

• Late 1700s - John Dalton - England
• Teacher- summarized results of his experiments
  and those of others in Daltons Atomic Theory
• Combined ideas of elements with that of atoms

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Daltons Atomic Theory

• All matter is made of tiny indivisible particles
  called atoms.
• Atoms of the same element are identical, those of
  different atoms are different.
• Atoms of different elements combine in whole
  number ratios to form compounds
• Chemical reactions involve the rearrangement of
  atoms. No new atoms are created or destroyed.

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Law of Definite Proportions

• Each compound has a specific ratio of elements
• It is a ratio by mass
• Water is always 8 grams of oxygen for each gram
  of hydrogen

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Law of Multiple Proportions

• If two elements form more that one compound, the
  ratio of the second element that combines with 1
  gram of the first element in each is a simple
  whole number.

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What the . . . .?

• Water is 8 grams of oxygen per gram of hydrogen.
• Hydrogen Peroxide is 16 grams of oxygen per gram
  of hydrogen.
• 16 to 8 is a 2 to 1 ratio
• True because you have to add a whole atom -- you
  cant add a piece of an atom.

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Next? Giving the atom parts.

• J. J. Thomson - English physicist. 1897
• Made a piece of equipment called a cathode ray
  tube.
• It is a vacuum tube - all the air has been
  removed.

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Thomsons Evidence
-

Voltage Source
Cathode Ray Tube
Gas at very low pressure
Cathode Ray
Cathode
Anode
Passing an electric current makes a beam appear
to move from the negative to the positive end.
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Thomsons Evidence
-

Voltage Source

Cathode Ray
- - - - -
Cathode
Anode
The addition of an electric field identified the
presence of negatively charged particles.
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Thomsons Evidence

• Concludes that cathode rays are streams of
  charged particles that are found in ALL kinds of
  atoms because
• The surrounding gas doesnt matter
• The metal of the cathode/anode doesnt matter
• Always the same results

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Thomsons Plum Pudding Model

• Found the electron
• Couldnt find positive (for a while)
• Said the atom was like plum pudding
• A bunch of positive stuff, with the electrons
  able to be removed

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Describing the electron

• Robert Millican -- American scientist
• 1916 Oil Drop Experiment
• Mass of the electron 1/1840 of the Hydrogen
  atom
• Charge of the electron 1.6 x 10-19 C
• Later determined mass of proton

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Looking for pudding and finding nothing

• Ernest Rutherford English physicist (1910)
• Believed in the plum pudding model of the atom.
• Wanted to see how big atoms are
• Used radioactivity -- Alpha particles -
  positively charged pieces given off by uranium
• Shot them at gold foil which can be
• made a few atoms thick

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Rutherfords Evidence

• When the alpha particles hit a florescent
• screen, it glows.
• A picture of his experimental set-up is on (page
  90)

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Rutherfords Evidence

• When the alpha particles hit a florescent
• screen, it glows.
• A picture of his experimental set-up is on (page
  90)

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Rutherfords Expectations

• The alpha particles would pass through
• without changing direction very much
• Because . . . the positive charges were spread
  out evenly. Alone they were not enough to stop
  the alpha particles

Path of alpha particle
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• Because he thought the mass of the atom was
  uniformly spread throughout its volume

Path of alpha particle
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Rutherfords Explanation

• Atom is mostly empty
• Small dense, positive piece at center
• Alpha particles are deflected by it if they get
  close enough

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Rutherfords Explaination


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But what keeps electrons (-) from falling into
the nucleus ()?

• Niels Bohr (1885-1962) -- Danish physicist
• New theory in 1913 -- electrons have orbits about
  the nucleus -- like the planets around the sun.
  This became the planetary theory.
• Proposed that electrons could only exist at given
  energy levels.
• He proposed this after looking at the light
  spectrum given off by hydrogen.
• Like a ladder, an electron can only be at any
  given rung of the ladder at any given time.
• To move between energy levels, the electron must
  gain or lose an exact amount of energy called a
  quantum of energy.
• Unlike a ladder, the energy levels are not evenly
  spaced.

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A New Physics develops a new atomic model

• Erwin Schrodinger (1887-1961) --Austrian
  physicist
• Albert Einstein
• Werner Heisenberg (1901 1976) German
  physicist

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Quantum Mechanical Model

• Electrons have only certain possible energy
  levels -- like the Bohr model.
• In a departure from the Bohr model (electrons in
  shells or orbits), the QM model estimates the
  probability of finding an electron in a given
  region.
• Electron, proton, and neutron now have sub-parts.
  
• We need to revise how we think about electrons
  they do not appear to be as simple as the point
  masses that we like to think of them as.

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Subatomic Particles
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Structure of the Atom

• The Nucleus
• With protons and neutrons
• Positive charge
• Almost all the mass
• Electron Cloud
• Most of the volume of
• an atom
• The region where the electron can be
• found

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Atoms are small . . .

• Measured in picometers, 10-12 meters
• Hydrogen atom, 32 pm radius
• Nucleus tiny compared to atom
• If the atom was the size of a stadium, the
  nucleus would be the size of a marble.
• Radius of the nucleus near 10-15m.
• Density near 1014 g/cm

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Counting the Pieces . . .

• Atomic Number of protons
• of protons determines kind of atom
• of protons of electrons for a neutral atom
• Mass Number protons neutrons
• All the things with mass

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Atoms are small . . .

• Measured in picometers, 10-12 meters
• Hydrogen atom, 32 pm radius
• Nucleus tiny compared to atom
• If the atom was the size of a stadium, the
  nucleus would be the size of a marble.
• Radius of the nucleus near 10-15m.
• Density near 1014 g/cm

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Elemental Symbols
X
Mass Number
Charge (as required)
Atomic Number
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Elemental Symbols
C
C
12
14
6
6
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Isotopes

• Dalton was wrong.
• Atoms of the same element can and do have
  different numbers of neutrons
• Different mass number for atoms of the same
  element due to additional neutrons
• Called isotopes
• Example C-12, C-13, C-14
• Example O-16, O-17, O-18

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Atomic Mass Units (AMUs)

• Mass of a proton 1.67267 x 10-27 kg
• Mass of a neutron 1.6750 x 10-27 kg
• Inconvenient numbers
• Instead, we define a new unit and reference it to
  the mass of a specific isotope of carbon
  carbon-12
• carbon-12 has 6 neutrons and 6 protons. By
  definition, we say carbon-12 as a mass of
  12.00000 amu -- atomic mass unit.
• 1 AMU 1.6606 x 10-27 kg.
• Then, 1 proton and 1 neutron have a mass of
  approximately 1 AMU.

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Why arent the atomic masses reported on the
Periodic Table close to whole numbers of AMUs?

• The reported numbers are average atomic mass
  units, reflecting the existence and relative
  abundance of isotopes for any given atom.
• IN NATURE, MOST ELEMENTS OCCUR AS A MIXTURE OF
  TWO OR MORE ISOTOPES.

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Atomic Masses

• Many atomic masses on PT are close to a whole
  number of AMUs sodium - 22.990 phosphorous -
  30.974 gold - 196.97
• But some are not chlorine 35.453.
• There are two naturally occuring isotopes of
  chlorine Chlorine-35 and Chlorine-37.
• If equal numbers were found in nature, we would
  expect chlorine to have an average atomic mass
  near 36 AMU. However, we do not get this because
  75 of chlorine is Cl-35. So the average is
  weighted toward the amu of Cl-35.

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ATOMIC MASS

• ATOMIC MASS the atomic mass of an element is a
  weighted average mass of the atoms in a naturally
  occurring sample of the element. It reflects
  both the mass and the relative abundance of the
  isotopes occurring in nature.

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Finding an Average Atomic Mass

• Naturally occurring Sulfur has four isotopes
• S-32 31.972 AMU 95.00
• S-33 32.971 AMU .7600
• S-34 33.967 AMU 4.220
• S-36 35.967 AMU .01400

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To determine the average atomic mass, do the
following

• Obtain the atomic mass and relative abundance of
  each isotope.
• Convert the abundance to a multiplication
  factor by dividing by 100.
• Multiply each isotopes atomic mass by its
  multiplication factor.
• Sum the products of Step 3.

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For Sulfur

• 31.972 AMU x .95
• 32.971 AMU x .0076
• 33.967 AMU x .0422
• 35.967 AMU x .00014
• 32.06242 AMU
• Or 32.06 AMU using the proper sig. figs.

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You try it

• Naturally occurring zinc is comprised of five
  isotopes. Determine its average atomic mass
  given the following

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Answer

• 65.38682 AMU, or 65.39 AMU when limited to the
  proper number of sig. figs.