History of the Atom

1
History of the Atom
2
The idea of an Atom

• Democritus and Leucippus Greek philosophers who
  thought maybe small particles like grains of sand
  could be made of something smallerbut they were
  over-ruled.
• The Early Greeks thought of everything as being
  made up of four basic elements because that is
  what Aristotle told them
• Earth.
• Air.
• Fire.
• Water.

3
John Dalton

• Dalton thought an atom was a tiny indivisible
  sphere of matter.
• 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.

4
The problems with Daltons theory

• following scientists did find that the atom
  contained smaller pieces
• and weve learned that an atom can be
  created and destroyed

5
J.J. Thomson
6
Thomsons cathode ray tube

• .

-

Vacuum tube
Metal Disks
7
Thomsons Experiment

-
8
Thomsons Experiment

-
9
Thomsons Experiment

-
10
Thomsons Experiment

-

• Passing an electric current makes a beam appear
  to move from the negative to the positive end

11
Thomsons Experiment

-

• Passing an electric current makes a beam appear
  to move from the negative to the positive end

12
Thomsons Experiment

-

• Passing an electric current makes a beam appear
  to move from the negative to the positive end

13
Thomsons Experiment

• By adding an electric field

14
Thomsons Experiment

-

• By adding an electric field

15
Thomsons Experiment

-

• By adding an electric field

16
Thomsons Experiment

-

• By adding an electric field

17
Thomsons Experiment

-

• By adding an electric field

18
Thomsons Experiment

-

• By adding an electric field he found that the
  moving pieces were negative

19
JJ Thompson

• Thompson concluded that the atom a divisible
  sphere that was positively charged and embedded
  with negative particles.
• The problem was not supported by Rutherfords
  experiment

20
Thomsons 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

21

• Other pieces
• Proton was found to be a positively charged
  particle 1840 times heavier than the electron
• Neutron - no charge but the same mass as a
  proton.

22
Ernest Rutherford

• Rutherford, a New Zealand scientist, had studied
  radiation and wanted to work with Thompson in
  investigating the make-up of the atomic
  structure.
• Rutherford had discovered three types of
  radiation and named them alpha particles, beta
  particles and gamma particles.
• Rutherford designed an experiment to probe the
  atom using radiation (alpha particles) as
  bullets.

23
Florescent Screen
Lead block
Uranium
Gold Foil
24
He Expected

• The alpha particles to 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

25
What he expected
26
Because, he thought the mass was evenly
distributed in the atom
27
What he got
28
How he explained it

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

29
(No Transcript)
30
(No Transcript)
31
Rutherfords gold foil experiment

• Conclusion every atom must have a nucleus
  containing a dense positive mass which has a lot
  of empty space around it in which tiny negative
  particles are found.

32

• As technology, chemistry and physics advanced the
  model of the atom was again questioned and
  altered by a Danish scientist named Neils Bohr.
• The problem with the Rutherford model was if
  the nucleus is positive and opposite charges
  attract, why didnt the negative electrons just
  crash in to the nucleus?

33
Modern View

• The atom is mostly empty space
• Two regions
• Nucleus- protons and neutrons
• Electron cloud- region where you might find an
  electron

34
Density and the Atom

• Since most of the particles went through, it was
  mostly empty.
• Because the pieces turned so much, the positive
  pieces were heavy.
• Small volume, big mass, big density
• This small dense positive area is the nucleus

35
(No Transcript)
36
Neils Bohr

• Bohrs model was compared to our solar system.
  While the Sun attracts the planets with great
  gravitational force, the plants move in an orbit
  due to the speed of their motion.
• The Bohr model describes the electrons in rapid
  motion with set amounts of energy around the
  nucleus. Where the electron is depends on how
  much energy it has. There are energy levels
  around the nucleus, like orbits, and due to their
  energy the electrons move around the nucleus in a
  set path.

37

• Bohrs Theory
• Allowed Orbitals
• An electron can only orbit around an atom in
  specific orbits
• Radiationless Orbits
• An electron in an allowed orbit does not emit
  radiant energy as long as it remains in the
  orbit.
• Quantum Leaps
• An electron gains or loses energy only by moving
  from one allowed orbit to another.
• The lowest energy state is known as the ground
  state
• Higher states are known as excited states

38

• Each time an electron males a "quantum leap,"
  moving from a higher energy orbit to a lower
  energy orbit, it emits a photon of a specific
  frequency and energy value.

39
Counting the pieces

• The number of protons in an atom determines what
  atom it isthis is referred to as the atoms
  atomic number.
• For neutral atoms, the number of protons
  (positive charges) and the number of electrons
  (negative charges) are equal.
• The mass number is the mass of the nucleus (the
  number of protons plus the number of neutrons).

40
Atomic Mass

• The mass of an atoms comes from the particles in
  the nucleus.
• Each proton and neutron has a mass of 1.66 x
  10-24 g. Comparably the mass of an electron is
  negligible.
• The mass of an individual atom is too small to
  measure, therefore we base the mass on ratios
  comparing elements to each other.

41

• Carbon-12 is used as the reference element for
  atomic mass.
• Carbon-12 has 12 protons and 12 neutrons
• therefore, defining one atomic mass unit as
  1/12 of a carbon-12 atom, means that
• 1 proton has a mass of 1 amu and
• 1 neutron has a mass of 1 amu.

42
Isotopes

• We now know that a number of the postulates of
  Daltons atomic theory were wrong.
• We know that particles smaller than the atom
  exist.
• We also know that atoms of the same element can
  have different masses, because they can have
  different numbers of neutrons. These are called
  isotopes.

43
Isotope Symbols
18 O
8 O

• To symbolize the composition of an isotope, the
  superscript to the left is used to represent the
  mass number, and the subscript to the left is
  used to represent the atomic number.
• Some common isotopes can be found on page 78 in
  your textbook.

44
Average atomic mass

• The average atomic mass of an element is based on
  the weighted average of its isotopes and their
  percent abundance

45
(No Transcript)
46
Quantum Mechanics

• The Quantum Concept.
• In 1900 Max Plank introduced the idea that matter
  emits and absorbs energy in discrete units called
  quanta.
• In 1905 Albert Einstein extended the quantum
  concept to include light and that light consist
  of discrete units called photons.
• The energy of a photon is directly proportional
  to the frequency of vibration.
• Ehf
• where E energy
• h Planks constant 6.63 X 10-34 J?s
• f frequency

47
Quantum Mechanics

• Quantum mechanics states that light and matter,
  including electrons, have a dual nature of both
  particles and waves.

48

• (A) Light from incandescent solids, liquids, or
  dense gases, produces a continuous spectrum as
  atoms interact to emit all frequencies of visible
  light (B) Light from an incandescent gas produces
  a line spectrum as atom emit certain frequencies
  that are characteristic of each element.

49

• Atomic hydrogen produces a series of
  characteristic line spectra in the ultraviolet,
  visible, and infrared parts of the total
  spectrum. The visible light spectra always
  consist of two violet lines, a blue-green line,
  and a bright red one.

50

• These fluorescent lights emit light as electrons
  of mercury atoms inside the tube gain energy from
  the electric current. As soon as they can, the
  electrons drop back to their lower-energy orbit,
  emitting photons with ultraviolet frequencies.
  Ultraviolet radiation strikes the fluorescent
  chemical coating inside the tube, stimulating the
  emission of visible light.