UNIT 3 ATOMIC STRUCTURE

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UNIT 3 ATOMIC STRUCTURE

• History of Atomic Theory Is matter continuous
  or discontinuous?
• Ancient Greeks most thought matter continuous
• Democritus said matter was made of parts
  atomos
• The Greeks did not believe in experimentation
  no way to prove!

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B. Daltons Theory of the Atom (1803)

• Based on experimentation
• Law of Conservation of Matter (Mass)
• Matter can neither be created nor destroyed
• During a Chemical change, Mass does not change
• Law of Definite Proportions A compound must
  have a fixed ratio of elements
• ex. H2O, CO2, NH3

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1. All matter is made up of parts. These parts,
called atoms, are indivisible (can not be broken
into smaller parts) and indestructible spheres.
2. Elements are one type of atom. All atoms of
the same element are exactly the same.
3. Compounds are formed by joining 2 or more
elements in fixed ratios.

• 4. In chemical reactions atoms are merely
  rearranged.

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C. Thompsons Model 1898

• Hypothesis atoms have parts and these parts are
  related to electricity.
• Experiment use a high electric voltage to
  remove a part of the atom (Cathode Ray Tube
  Experiment).
• Results the high voltage caused a negatively
  charged ray to be discharged from a piece of
  metal.

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• Thompsons Theory
• Atom has a negative part electron
• Since matter is neutrally charged, there must be
  a yet undiscovered positive part to balance the
  charge proton
• These parts must be spread out in the atom so
  that like charges dont repel each other
• The Plum
• Pudding
• Model

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D. Rutherford's Theory of the Atom (1908)

• Based on evidence from the alpha
• particle, gold foil experiment.

• Since alpha particles bounced back, there must
  be something large and positively charged in the
  atom

• All the positively charged protons are grouped
  together in the middle of the atom in a nucleus.

• Electrons are scattered about in the large
  empty space around the nucleus.

• Called the NUCLEAR MODEL of the atom.

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E. Problems with the Rutherford Model

• What holds the nucleus together?

• Neutrons James Chadwick (1932) discovered
  that a neutral particle could be knocked out of
  Beryllium metal foil. This neutral particle,
  called a neutron, acted as a buffer between
  protons.

• Nuclear energy Einstein (1920) theorized
  that an unknown energy existed that acted as a
  glue to hold the nucleus together (nuclear
  energy).

• Where are the electrons exactly?

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• II. Modern description of the atom

• Parts of the atom subatomic particles

amu atomic mass unit u 1
amu 1/12th the mass of a C-12 atom Note
actual mass of electron is 1/1836 amu
or 0.00053amu which is nearly zero
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B. Structure of the Atom

• Protons and Neutrons are found together in the
  nucleus called nucleons.
• Electrons found outside of nucleus.
• Atomic number the number of protons all atoms
  of the same element must have the same number of
  protons!
• Mass number protons neutrons (notice
  electrons not included, mass 0)
• Charge of atom protons electrons (notice
  neutrons not included, charge 0)

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C. Symbols of atoms of elements

• To identify a specific atom of an element, the
  atomic and mass number can be given
• ex. mass no. 12 or just mass
• atomic no. 6 number
• The atomic number for any element can be found on
  the periodic table so does not need to be
  specified.
• The mass number can vary so it must be specified.
  

C
C 12
pn
p
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D. Isotopes

• A specific mass version of an element
• Isotopes of the same element have different
  masses because they differ in the number of
  neutrons.
• Ex. Carbon has 3 naturally occurring isotopes
• C C C
• p 6 6 6
  atomic no.
• e 6 6 6 no.
  of protons
• n 6 7 8 mass
  no. protons
• For a neutral atom, no. of electrons no. of
  protons

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14
12
6
6
6
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E. Ions

• A neutral atom has equal number of protons and
  electrons.

• An ion is a charged atom, not neutral.
  Electrons are either gained or lost!

• A negative charge means electrons are gained, a
  positive charge means electrons are lost.

• No. of electrons atomic no. - charge

_
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Ex. 6 C2 charge ex. 9 F
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F. Atomic Mass

• The atomic mass is a weighted average mass of all
  naturally occurring isotopes of an element. This
  average is not a whole number!
• Used when dealing with a group of atoms occurring
  naturally, not just one atom! The mass of a
  single atom is its mass number and is a whole
  number.
• The weighted average takes into account an
  isotopes abundance in nature. The value of the
  atomic mass is closest to the mass number of the
  most abundant isotope.

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Ex. Hydrogen

• Hydrogen has three well known isotopes

Atomic mass Mass no. x abundance for each
isotope
100
Atomic mass H 1 x .99206 2 x .00794
.99206 .01588 1.00794 amu
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III. Electron Structure

• Bohr model of the atom Neils Bohr asked where
  are the electrons exactly and why dont they fall
  into the nucleus?
• Bohr knew Moving electrons give off light.
• Experiment Move an electron by jolting it
  with various amounts of electrical energy and
  measure the light energy given off.
• Expected result the electron can exist at any
  distance outside of nucleus so all values of
  light energy should be obtained a continuous
  spectrum.

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Electromagnetic Spectrum
Prentice Hall Chapter 5 simulation 3
Teacher Domain electromagnetic spectrum
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Atom jolted with varying amounts of energy should
give off various
wavelengths
of light
Expected Results
nucleus
a continuous spectrum!
The electron jumps from its normal low energy
position, (the ground state), to a higher energy
position, (the excited state), and gives off a
specific wave length of light (a color). If an
electron can exist anywhere outside the nucleus,
all colors of light should be seen.
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Actual results
Only certain amounts of energy absorbed, only
certain colors of light obtained. The electron

can exist only certain distances
nucleus
from nucleus
1 2 3 4
Bright line spectrum
Principle quantum numbers
Energy levels Places where electrons can be
found
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B. Making Bohr Models of an Atom

• Bohr drew rings around the nucleus to represent
  these places the electron can exist.
• The rings are called orbits, stationary states,
  shells or principle energy levels.
• The electron can posses only certain quantities
  (a quantum) of energy the energy is quantized.

1st energy level holds 2 electrons 2nd energy
level holds 8 electrons 3rd energy level holds
18 electrons 4th energy level holds 32 electrons
4 3 2 1

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C. Electron Configurations

• A list of the number of electrons on each energy
  level of a Bohr model.
• Found at the bottom of each element box on the
  Periodic Table.
• When drawing Bohr models of elements higher than
  18, use the electron configurations in the table
  to fill the higher energy levels.
• Ex. Germanium Ge
• 32 electrons
• 2-8-18-4

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D. Ground State vs. Excited State

• Ground state is the lowest energy configuration
  of electrons in an atom it is the normal
  configuration found in the table.
• ex. Oxygen 2-6
• Excited state is any higher energy state it is
  not the normal configuration.
• ex. 2-5-1 (one electron from the 2nd level
  moved to the 3rd level)
• What element is this? Add up electrons on each
  level (8) and look up atomic number (oxygen).
• Is this the ground or excited state? Compare to
  oxygens configuration in the table, if its the
  same its a ground state, if different its an
  excited state.

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E. Valence Electrons

• Electrons in the outer most (highest) energy
  level only.
• These electrons are involved in chemical
  reactions.
• There can be no more than 8 valence electrons.
• Kernel all inner electrons and nucleus.
• ex. Li 2-1
  Valence
• 1 valence e-
  Kernel

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F. Dot Structures

• Shows only the valence electrons.
• Use the letter symbol to represent the kernel and
  dots to represent the valence electrons.
• Use electron configuration in table to tell the
  number of valence electrons (dots) or
• Use column number (without teen).
• Ex. Beryllium Be Oxygen O

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IV. Quantum (Wave) Mechanical Model

• Heisenberg Uncertainty Principle
• You can not measure exactly where an electron is
  and where it is going because when you measure
  one value you change the other!
• If the error in a measurement is larger than the
  measurement, the measurement is useless!
• The Bohr model fails because it tries to tell
  exactly where an electron is and how it moves.

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B. Orbitals

• Electrons can not be placed in exact spots.
• Orbitals are probability areas in space where
  electrons may be found electron clouds.
• The simplest such area is the s or spherical
  orbital
• 
  

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Prentice Hall Chapter 5 Animation 5