Download History of Atomic Theory

History of Atomic Theory
Alchemy ~ Before 400 B.C.
Experiment: Pseudoscience concerned with:
• Changing metal to gold
• Finding an eternal life elixir
Aristotle
Beliefs:
• All matter was made up of a combination of the four elements
• Four elements: fire, wind, earth, water
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History of Atomic Theory
Democritus ~ 400 B.C.
Experiment: None. He had beliefs that were disregarded.
Beliefs:
• Named the atom "Atomos." It means indivisible
• Matter is composed of atoms too small to be seen.
• Empty space between atoms
• Atoms are solid and homogeneous
• Different atoms have different sizes and shapes
Democritus
Model: Sphere
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History of Atomic Theory
Dalton ~ early 1803 English School Teacher
John Dalton
Experiment: Studied chemical reactions, making observations and measurements
Beliefs:
Five Principles-Dalton's Atomic Theory
1. All matter is made of indestructible atoms.
2. Atoms of the same element are identical in their physical and chemical properties.
3. Atoms of different elements have different physical and chemical properties.
4. Atoms of different elements combine in simple whole number ratios to form chemical
compounds.
5. In chemical reactions, atoms cannot be subdivided, created, or destroyed. They are
combined, separated, or rearranged.
Model: Sphere-Billiard balls
Which principles remain true today?
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History of Atomic Theory
Experiment: Cathode Ray Tube
Electricity is passed through a gas tube. The gas
beam can be bent with a magnet.
Joseph John Thomson
J.J. Thomson ~ 1897 English Physicist
Discoveries:
• Atoms consist of charged particles
• The negatively charged particles are called electrons (1897)
• The positively charged particles are called protons (1920)
Model: Plum Pudding or Chocolate Chip Cookies
The chocolate chips are electrons stuck in positive dough
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History of Atomic Theory
Rutherford ~ 1911 English Physicist
(a student of Thomson)
Experiment: Gold Foil Experiment
Positively charged alpha particles are shot
at a piece of thin gold foil. Most alpha
particles had little deflection. Some were
deflected at large angles.
Ernest Rutherford
Discoveries:
• A positively charged core of an atom called the nucleus
• Electrons surround the nucleus
• The rest of the atom is empty space
Model: Nuclear atom
Quote: "It was about as believable as if you had fired a 15 inch shell at a
piece of tissue paper, and it came back and hit you."- Rutherford
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History of Atomic Theory
James Chadwick ~ 1932 English Physicist
(Student of Rutherford)
Experiment:
• Beryllium foil was bombarded with alpha particles
Sir James Chadwick
• A neutral radiation was emitted
• Emitted radiation would then knock protons out of the nuclei of other substances
• The radiation was a stream of neutral particles having the same mass as a proton
Discoveries:
• A neutrally charged subatomic particle
• The particle was called a neutron
Model: Same as Rutherford
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Review History of Atom
http://glencoe.mcgraw­hill.com/sites/dl/free/0078759864/164155/00044672.html 7
History of Atomic Theory
Bohr ~ 1913 Danish Physicist
(a student of Rutherford)
Problem: According to the laws of physics, charged particles
will radiate energy when orbiting and spiral into the
nucleus. Rutherford's atoms would all collapse.
Niels Bohr
Experiment: Light Spectrums
Electrons give off energy in the form of colored
light by falling from an excited state to a
ground state.
Discoveries:
• Electrons can only be found at certain energy
levels
• Each energy level requires a certain amount of
energy
• Lower (closer) levels have lower energy
• Higher (farther) levels have higher energy
Model:
Photo courtesy NASA
Hydrogen Spectrum
Helium Spectrum
Ladder Rungs-Cannot stand in between rungs
Vocabulary:
ground state- all electrons in their lowest possible
energy levels
excited state-electrons absorbed energy & jump to
a higher energy level
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History of Atomic Theory
Quantum Mechanics
Erwin Schrodinger ~ 1926 Austrian Physicist
Problem: Bohr's model works well for Hydrogen, but fails for
every other element
Mathematical Equation:
Experiment:
Erwin Schrodinger
NONE! It is a mathematical model. It cannot be
represented by anything that exists in the real world.
Discoveries:
• Mathematical model that deals with the probability of finding an electron within a given
space
• The probability is 90%
• The given space are called orbitals (or electron clouds)
• There are four orbitals with different shapes s p d f
• These orbitals can be related to the periodic table
• Electrons have wave properties
Model:
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Electromagnetic Radiation: waves that are produced by electrically charged
particles
Ex: sunlight, X rays, microwaves
Electromagnetic Spectrum
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All electromagnetic radiation exhibits wave like behavior
(wavelength, frequency, and speed)
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Quantum Numbers: Specify the properties of atomic orbitals and the properties of electrons
in those orbitals. Each e- has a set of four numbers; no two electrons in
the same atom can have the same four numbers.
Four Quantum Numbers:
Principal
Angular
Magnetic
Spin
1. Principal Quantum Number (n)- Refers to the distance of the orbital from the
nucleus (says which of the main energy levels an
e- is in)
• When n=1 is closest to the nucleus and has the least energy
• n=1,2,3,4,5, etc.
2. Angular Quantum Number (l)- Refers to the shape of the orbital (also associated
a. Possible Shapes
with angular momentum)
• s, p, d, f (in order of increasing energy)
• l=integer values from 0 to n-1
b. Number of possible shapes is limited by the principal quantum number
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3. Magnetic Quantum Number (m)- Orientation of orbital(s) (direction of angular
momentum)
- m=any integer from -l to +l
4. Spin Quantum Number (s)- State of the electron that occupies an orbital
-Electrons are assigned one of the two possible directions it can be spinning
- + spin or - spin
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Heisenberg's Uncertainty Principle: you can
never know how fast an electron is moving and
where an electron is at the same time. In
other words, you can find out where the
electron started and you can see where the
electron ended up but how it got there WE
DON'T KNOW!
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Electron Configuration: The arrangement of electrons in an atom
is known as the atoms electron configuration.
Rules:
1. Aufbau Principle: An electron occupies the lowest energy level that can receive it.
2. Pauli Exculsion Principle: No two electrons in the same atom can have the same
set of four quantum numbers.
• If two electrons have the same n, m, and l values they have to have different spins
• Two electrons fit into each orbital
3. Hund's Rule: Orbitals of equal energy are each occupied by one electron before
any orbital is occupied by a second electron, and all electrons in singly occupied
orbitals must have the same spin
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4p
3d
4th energy level
4s
3p
3rd energy level
3s
2nd energy level
2p
2s
1st energy level
1s
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Attachments
Pictures of atom
Chadwick apparatus
Time line of scientists
Schrodinger equation
Conversation with science
orbitals