Index
Atoms and Elements
Chemistry Honors Unit 02
Based on the PowerPoints by Kevin Boudreux.
Honors Unit 02 Module 02: Atomic Theory
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Unit 02 Module 02
ATOMIC THEORY
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The Discovery of the Nucleus
• In 1910, Ernest Rutherford tested the “plum pudding”
model of the atom by firing a stream of alpha (α) particles at
a thin sheet of gold foil (about 2000 atoms thick).
• Alpha particles are a type of radiation given off by many
naturally-occurring radioactive elements (Ra, Rn, Po, U,
etc.). They are 7000 times more massive than electrons and
have a positive charge that is twice the magnitude of the
electron’s charge.
• It the “plum-pudding” model were correct, the mass of the
atom would be spread out evenly through the entire volume
of the atom, and all the alpha particles should have sailed
right through the foil – but that’s not what happened…
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The Discovery of the Nucleus
• …instead, while most of the alpha-particles
sailed through the gold foil, some were
deflected at large angles, as if they had hit
something massive, and some even bounced
back toward the emitter.
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The Discovery of the Nucleus
Rutherford’s Experiment – Nuclear Atom
Movie remove to decrease file size. To view the movie follow the link listed below.
If you are on the school network, you will need to go to my YouTube channel
(William Habiger) and select the video from the Unit 02 play list.
My YouTube channel:
http://www.youtube.com/user/habiger?edufilter=TC5tGIQ9RIvel5OGeEG9ew&safe=a
ctive
Source:http://www.youtube.com/watch?v=5pZj0u_XMbc
Other Resources:
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/ruther14.swf
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Rutherford’s Gold Foil Experiment
Video: Rutherford Gold Foil Experiment – Backstage Science
Movie remove to decrease file size. To view the movie follow the link listed below.
Source:http://www.youtube.com/watch?v=XBqHkraf8iE
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The Nuclear Atom Model
• Rutherford concluded that all of the positive charge and
almost all of the mass (≈99.9%) of the atom was
concentrated in the center, called the nucleus. Most of
the volume of the atom was empty space, through
which the electrons were dispersed in some fashion.
• The positively charged particles within the nucleus are
called protons; there must be one electron for each
proton for an atom to be electrically neutral.
• This did not account for all of the mass of the atom, or
the existence of isotopes (more later); the inventory of
subatomic particles was “completed” (for the moment)
by James Chadwick in 1932, who discovered the
neutron, and uncharged particle with about the same
mass as the proton, which also resides in the nucleus.
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Rutherford’s Model of the Nuclear Atom
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Spectrometer
Spectrometers use diffraction (diffraction grating) or refraction (prism) to separate light into
its different components.
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Emission Line Spectra
The lines (by wavelength or frequency) show the energy differences between energy levels
in an atom.
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Atomic Fingerprints
• Each element has it own set of wavelengths of radiation
(visible light, infrared light, ultraviolet light, etc…) that
it emits when energetically excited. These spectral
lines can be used as “fingerprints” to identify which
elements are present in a sample, the sun, or a star.
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Emission (Bright Line) Spectra
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The Bohr Model of the Hydrogen Atom
• In 1913, Neils Bohr (Nobel Prize, 1922) suggested a model for the
H atom that explained line spectra. In this model, the energy levels
in atoms are quantized, having only certain allowed energy levels
associated with fixed electron orbits called stationary states.
– When an atom absorbs an amount of energy equal to the energy
difference between two fixed orbits (E1), an electron jumps from the
low energy orbit (the ground state) to a higher-energy orbit (an excited
state)
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The Bohr Model of the Hydrogen Atom
– When an atom releases energy, the electron falls back down to
the ground state, releasing a photon of light which corresponds
exactly to the energy difference (E1) between orbits.
• In this model, line spectra arise because the atom’s
energy has only certain discrete levels, and atoms
can absorb or emit energy only in these “chunks.”
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The Bohr Model
• The Bohr model rationalized the existence of line
spectra, and mathematically predicted the
wavelengths of radiation emitted by the hydrogen
atom in its line spectrum.
• Unfortunately, the Bohr model couldn’t predict
where these fixed electron orbits came from and
why the electrons didn’t lose energy over time. It
also failed to predict the location of line spectra
for elements with more than one electron.
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Summary
of
Bohr
’
s
Model
of
the
hydrogen
atom.
The Model:
Index
•
The energy levels or orbits in an atom are quantized (i.e. allowed to have
only certain energies.)
•
To be in that energy level, an electron must have that exact amount of
energy.
•
The electron stays in it’s ground state unless excited.
•
If a photon of light with the amount of energy equal to the difference in
energy levels strikes the electron, the electron will move to that energy
level.
•
When an electron drops from an excited energy level to the ground state, a
photon of light is released. Ephoton = E
•
Model only works perfectly for hydrogen.
•
Exactly explains the line spectrum for hydrogen.
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Bohr’s Model of the hydrogen
atom.
The Model:
• The nucleus is surrounded by energy levels (orbits).
• The electron stays in the ground state unless excited.
• If a photon of light with the amount of energy equal
to the difference in energy levels strikes the
electron, the electron will move to that energy level.
• When an electron drops from an excited energy
level to the ground state, a photon of light is
released. Ephoton = E
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Bohr’s Model w/ energy levels
http://www.colorado.edu/physics/2000/quantumzone/bohr2.html
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Bohr’s Model
http://www.colorado.edu/physics/2000/quantumzone/bohr.html
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“Not Really” Bohr Model
Oxygen
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The Atomic Theory Today
• An atom is an electrically neutral, spherical entity
composed of a positively charged central nucleus
surrounded by one or more negatively charged electrons.
• The nucleus contains about 99.7% of the atom’s mass, but
occupies 1 ten-trillionth of its volume.
• The nucleus contains the protons, which have positive
charges, and neutrons, which are neutral. Neutrons are
very slightly heavier than protons; protons are 1835 times
heavier than electrons.
• The electrons (e-) surrounding the nucleus have negative
charges. The number of protons in the nucleus equals the
number of electrons surrounding the nucleus in a neutral
atom.
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Modern Atomic Theory
1. All matter is composed of atoms.
2. Atoms of one element differ in properties from
atoms of another element.
3. Atoms are divisible into smaller particles called
subatomic particles
4. A given element is made up of atoms with
different masses. (Isotopes)
5. Chemical reaction involves only the combination,
separation, or rearrangement of atoms.
6. Compounds are formed by the combination of
atoms in smallHonors
whole
number ratios.
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Quantum Model of Atom
• Energy levels consist of main energy levels
(shells) which are made up of subshells.
Subshells are made up of orbitals.
• An orbital is a three dimensional space in which
there is a 90% probability of containing an
electron.
• Electrons are particles that behave like a standing
wave.
• Electrons appear to move randomly within their
orbitals however they follow a certain wave
pattern according to the amount of energy they
have.
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Index
Structure of Atoms
Movie removed because you do not have access to the movie file.
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Electron Orbitals
Electron Density
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Electron Orbitals
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Electron Orbitals
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Examples
Hydrogen
e-
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Examples
Helium
ee-
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Examples
Lithium
e-
e-
e-
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Examples
Beryllium
e-
e-
e-
eHonors Unit 02 Module 02: Atomic Theory
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Examples
Beryllium
e-
e-
e-
eHonors Unit 02 Module 02: Atomic Theory
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Examples
Boron
eee-
e-
e-
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Examples
Carbon
e-
eee-
e-
e-
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Examples
Nitrogen
e-
eee-
e-
e-
e-
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Examples
Oxygen
e-
eee-
e-
e-
e-
e-
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Examples
Fluorine
e-
eee-
e-
e-
e-
e-
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Examples
Neon
e-
e-
eee-
e-
e-
e-
e-
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Orbitals
S type orbitals
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Orbitals
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Orbitals
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Orbitals
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