History of the Atomic Model

1
History of the Atomic Model

• Chapter 4

2
Sir William Crookes (1879)

• Invented the cathode ray tube and investigated
  electrical charges in gases.
• Postulated a negatively charged particle that was
  eventually named the electron.

3
Schematic of a Cathode Ray Tube
4
Henri Becquerel (1896)

• Noticed that certain substances emitted radiation
  without any external energy source.
• Two types of radiation called alpha (a) and beta
  (b) due to their penetrating power.
• Third was later added on gamma (g)

5
John Joseph (JJ) Thomson (1903)

• Using the cathode ray tube invented by Crookes,
  discovered the electron, e.
• Also calculated the charge to mass ratio of an
  electron.
• Qe/me -1.7588x1011 C/kg
• Proposed the plum pudding model.
• Postulated positive particles.

0 -1
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Robert Millikan (1909)

• Working with Harvey Fletcher,
• Determined the charge on a single electron, Q
  -1.602x10-19 C
• Since Thomson found Q/m, they were able to
  calculate an electrons mass
• Which is 1,836 times lighter than a hydrogen
  atom.

What is an Electrons Mass?
9.11x10-31 kg
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Millikans Oil Drop Experiment
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Ernest Rutherford (1909)

• Discovered the positively charged dense central
  portion of the atom using his gold foil
  experiment
• ? Nucleus
• Atoms are mostly empty space!
• Discovered the proton ( p) in 1918

1 1
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Rutherfords experiment
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Results of foil experiment if Plum Pudding model
had been correct
Failure!
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Actual Results
12
An atom viewed in cross section

• The nucleus is 100 thousand times smaller than
  the atom!
• If the nucleus were the size of a marble, the
  atom would be as big as a football stadium!

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Timeline So Far

• 1803 Daltons Atomic Theory
• 1879 Crookes postulates a negatively charged
  particle ? electron.
• 1896 Bequerel discovers radioactive materials.
• 1903 Thompson defines the charge to mass ratio
  of the electron (thereby discovering it).
• 1909 Millikan discovers the charge (and
  therefore mass) of the electron
• 1909 1918 Rutherford discovers that the atom
  is mostly empty space and eventually the proton.

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Problem to Solve

• If electrons are negatively charged and protons
  are positively charged, what keeps them
  separated?
• Why dont they simply smash together like
  magnets?
• Hang on to your hats, its about to get weird!

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Updated Atomic Model (post 1909)

• Discarded Plum Pudding
• Electrons must orbit around central nucleus ?
  Planetary Model
• Nearly all of the mass is located in the dense,
  central, positively charged nucleus.
• Still, why dont the electrons fall into the
  nucleus like opposite magnetic poles?

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Enter the Dane

• Niels Bohr (1885 1962)
• Danish physicist
• Went to Cambridge with the complete works of
  Dickens
• (to learn the language, duh!).
• Worked on the electron/nucleus problem
• Often thought of as the Father of Quantum
  Mechanics

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Bohrs Solution (1913)

• Electrons are located in specific energy levels
• Electrons move in a definite orbit around the
  nucleus
• Areas between orbits are not allowed!

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More Problems to Solve

• Why are electrons limited to specific orbits?
• Quantum mechanics has all the answers uhsort
  ofmaybeprobablykind of
• With exception of hydrogen, atoms weigh more than
  sum of protons and electrons there must be
  another particle

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Walther Bothe Herbert Becker (1928)

• Aimed alpha radiation at light elements like
  boron.
• Found it gave off an extremely penetrating
  radiation.
• Thought it produced high energy gamma rays.

Bothe
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Radiation Comparison

• Alpha Radiation positively charged, easily
  stopped (paper, skin, etc)
• Beta Radiation negatively charged, stopped by a
  sheet of aluminum
• Gamma radiation high energy light (no charge),
  penetrates a lot of material
• New Radiation - ??? penetrates even more

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Irene Joliot-Curie (1897-1956)

• Daughter of Marie Pierre Curie
• Continued work of Bothe and Becker (1932)
• Aimed Bothes new beam at paraffin
• Ejected high energy protons
• Misinterpreted results cost her a Nobel Prize
• Also found a way to transmute elements! (1934)
• Inexpensive way to create radioisotopes for
  medicine
• Just like her mother, died from radiation exposure

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James Chadwick (1891-1974)

• Did not believe the work of Joliot-Curie
• She said beam was light waves
• He thought not enough mass
• Discovered neutron in 1932.
• No charge.
• Mass slightly more than a proton.

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Subatomic Particles
Name Abbrev Location Mass (kg) Charge (C) Effective Charge
Electron e- or e Orbit 9.019x10-31 -1.602x10-19 -1
Proton p or p Nucleus 1.673x10-27 1.602x10-19 1
Neutron n or n Nucleus 1.675x10-27 0 0
0 -1 1 1 1 0
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Comparison (contd)

• Protons Heavy, positive charge ? repelled by the
  nucleus.
• Electrons Nearly massless, negative charge ?
  repelled by other electrons surrounding an atom.
• Neutrons Heavy, no charge ? no interactions with
  nucleus or electrons.
• Can pass through a lot of material.
• If fast enough, can break nucleus apart!

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Atomic Requirements

• All atoms of a given element must have the same
  number of protons!
• 1 Proton 1 Dalton (1 Da 1.6605x10-27 kg)
• In order to be neutral, atoms must have same of
  electrons as protons (charges must cancel out to
  zero)
• Electrons do not add weight (compared to p n)
• Total weight comes from additional neutrons.
• 1 Neutron 1 Da

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Solving the Mass Problem

• Total mass is sum of protons, neutrons, and
  electrons.
• 1 Atom of Helium has 2 protons, 2 electrons, and
  2 neutrons 4 Da
• But, it turns out that some atoms of elements
  weigh more (or less) than the others

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Elements and Neutrons

• Atoms can have more (or less) neutrons
• Isotopes are atoms with the same of protons,
  but different of neutrons
• Since protons determine chemical identity,
  neutrons just add mass.
• Isotopes occur in different ratios in nature.

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Two isotopes of Sodium.
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Specifying Isotopes

• Two ways to write an isotope
• C or simply C
• Carbon-12
• These isotopic symbols tell you how many of each
  particle is in the isotope.
• The Mass number (A) is the sum of p n
• The Z number is the number of p
• The Charge, Q, is p - e-

12 6
12
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Uranium-235

• Uranium has 92 protons
• Needs 92 electrons to cancel them out
• Its mass number is 235 (92 p 143 n)

235
92 -92 0, so no charge!
U
92
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Nitrogen-14 ion
Mass Number (A) p n
Charge p - e-
Nitrogen has 7 protons and 7 neutrons. It tends
to gain 3 extra electrons when it forms an ion.
-3
14
N
7
Atomic Symbol
Z Number p (nuclear charge)
Lets Practice!
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Mass of Atoms

• Mass of an atom is sum of its p, n, e-.
• Subatomic particles are really (really) light.
• Use the Dalton (Da) to measure the mass also
  called the unified mass unit (u) or atomic mass
  unit (amu)
• 1 Da 1/12th mass of a Carbon-12 atom
• 1 Da 1.6605x10-27 kg

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Particle Weights

• Electrons lightest of subatomic particles
• 0.0005 Da (usually ignored)
• Protons
• 1.0073 Da
• Neutrons heaviest of subatomic particles
• 1.0087 Da
• p n are usually rounded to 1 Da, while e- are
  assumed to be 0 Da

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Atomic Mass

• Mass of a specific isotope.
• Carbon-12 12.00 Da
• Carbon-13 13.00 Da
• Oxygen-16 15.9949 Da
• Must be measured experimentally due to Mass
  Deficit

6 p 6 n
6 p 7 n
8 p 8 n???
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Mass Deficit

• When subatomic particles combine, their masses
  change (dont ask why)
• Sometimes, the combination weighs less, other
  times, it weighs more (depends on numbers of each
  combining).
• Gain/lose mass according to Einsteins famous
  equation E mc2
• Atomic Hydrogen (Nuclear) bombs release this
  energy!

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Carbon-12

• Was chosen as the standard atom.
• All atomic masses are based on this specific
  isotope.
• Mass of C-12 defined as exactly 12.00 Da.
• Masses of other atoms are relative to C-12.

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Atomic Masses of Isotopes

• Uranium Isotopes Oxygen Isotopes

Isotope Mass (Da)
14 14.008596
15 15.003066
16 15.994915
17 16.999132
18 17.999161
19 19.003580
20 20.004077
Isotope Mass (Da)
232 232.037162
233 233.039635
234 234.040952
235 235.043930
236 236.045568
237 237.048730
238 238.050788
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Abundance of Isotopes

• Isotopes naturally occur in different
  proportions.
• Oxygen isotopes

Isotope Mass (Da) Abundance
O-16 15.994915 99.757
O-17 16.999132 0.038
O-18 17.999161 0.205
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Atomic Weight

• Different from Atomic Mass
• Also called Relative Atomic Mass
• Average weight of all naturally occurring
  isotopes.
• Provides more accurate weight of a typical sample
• A weighted average

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Calculating Atomic Weight

• Atomic Mass Abundance mass fraction
• 100
• Carbon has 2 stable isotopes
• Carbon-12 12.00Da is 98.93
• Carbon-13 13.00Da is 1.07
• 98.93/10012.00 Da 11.872 Da
• 1.07/10013.00 Da 0.139 Da
• 12.011 Da

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Oxygens Atomic Wt
Isotope Mass (Da) Abundance
16 15.994915 99.757
17 16.999132 0.038
18 17.999161 0.205

• 15.9949Da 99.757/100 15.9560Da
• 16.9991Da 0.038 / 100 0.0065Da
• 17.9992Da 0.205 / 100 0.0369Da
• Total 15.9994Da

Thats the on the Periodic Table!
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Summary

• 3 subatomic particles
• electron (-1, nearly massless)
• proton (1, 1 Da)
• neutron (0, 1 Da)
• Isotopes same protons, different neutrons
• Written two ways
• Atomic Mass mass of 1 atom
• Atomic Weight average weight of all isotopes