Download Atomic Theory

You too can be as smart as Einstein
(almost)
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
The elements
 Earth – dry, heavy
 Water – wet, heavy
 Air – cool, light
 Fire – warm, light
The composition of a substance could be
estimated from its properties.
These ideas were based on observation,
logic and reason, but not
experimentation.
 Democritus
(460 B.C. - 370 B.C.)

Matter is made of small, hard indivisible
particles called atoms, which exist in the
void.
 These atoms differ in size and shape, but
not in any other way.
Quantitative differences (how much) vs.
Qualitative differences (what kind)

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Antoine Lavoisier
(1743-1794)
Discoverer of Oxygen
(disputed)
His work refuted the
phlogiston theory
Responsible for the
law of conservation
of matter.

Claude Louis Berthollet
Joseph Louis Proust

Berthollet – “compounds do not have a
fixed composition”.
Cu +

S

CuxSy
Every time he tried the experiment he got
a different result.

Proust - compounds have a fixed
composition.
2H2 + O2  2H2O

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He always got the same result.
Proust’s argument is called The Law of
Definite Proportions. He was proved to
be right.
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John Dalton (ca.
1804)
The father of
modern atomic
theory
Schoolteacher
Colorblind –
studied
colorblindness

The points of Dalton’s theory
All matter is made of atoms
 Atoms are indivisible and indestructible
 All atoms of one element are exactly alike,
and atoms of different elements are different.
 Atoms combine in small whole number
ratios to form compounds.



The Law of Multiple Proportions:
If two elements combine to make two
different compounds, the ratios of the
elements involved are small whole
numbers.
Examples: CO and CO2
CuS and Cu2S
H2O and H2O2

J. J. Thomson and the Electron (1897)
Thomson discovered the electron - he
called it a “corpuscle”.
 He used an instrument called a Crookes
tube.
Cathode (-)
Evacuated tube Anode (+)



He noticed a stream of charged particles
coming from the cathode, called cathode
rays.
Thomson proposed the "plum pudding"
atomic model - negatively charged
corpuscles swarm inside a cloud of
massless positive charge.

The gold foil experiment (1909)


Most of the alpha particles went straight
through, and a few were bounced
straight back.
Rutherford’s interpretation: The atom
has a small, hard, dense and positively
charged nucleus. The electrons are
outside the nucleus.

Discovery of the proton: Henry Moseley
(1913)
Moseley bombarded metals with x-rays
 Each successive element had one more
positive charge – called “atomic number”
 Rutherford proved that the nucleus of
nitrogen contains hydrogen nuclei – a
“proton” (1918-19)

Discovery of the neutron – James Chadwick
(1932)

Name
Charge
Mass (amu)
Location
Discoverer
Electron
-1
1/2000
outside
nucleus
Thomson
Proton
+1
1
nucleus
Moseley/
Rutherford
Neutron
0
1
nucleus
Chadwick

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Atomic number = number of protons in the
nucleus
Atomic number determines the identity of the
element
Mass number = protons + neutrons
Number of electrons = number of protons
Isotopes: two atoms of the same element with
different numbers of neutrons
C-12 and C-13 are isotopes of carbon

Nuclear symbols
13C
6

Write the nuclear symbol for lead-206.
206Pb
82
Atomic number
20
Symbol
Name
Ca
Calcium
40.078
Average atomic mass
Average mass of all the isotopes of an element
 Average is weighted
 Example: Boron has two isotopes, B-10 and
B-11
B-10: 19.9%
B-11: 80.1%
 Average atomic mass of boron:
10x0.199 = 1.99
11x0.801 = 8.811
Average atomic mass = 1.99 + 8.811 = 10.8amu


Niels Bohr and the stepwise atom (ca.
1918)

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
Rutherford suggested that electrons orbit
around the nucleus like planets around
the sun.
This did not explain emission spectra,
which gave sharp lines.
He theorized that electrons could only
travel in certain sized orbits, and not
anywhere in between.


The orbits were called energy levels.
Each orbit has a specific energy.
Electrons can jump from one level to
another; as they do, they absorb or emit
energy.

Erwin Schrödinger and probable cause
(ca. 1935)



Schrödinger’s work showed that electrons do
not move in actual “orbits”.
Electrons move randomly and form
“probability clouds”. The shape of these
clouds is similar to the shape of Bohr’s orbits.
The position and momentum of an electron
cannot be determined simultaneously
(Heisenberg Uncertainty Principle)

Schrödinger’s “electron cloud”

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Bohr suggested that electrons inhabit energy
levels around the nucleus.
Each level has a specific energy associated with
it.
The outermost (highest energy) level is called
the “valence shell”.
The electrons in the valence shell are called the
“valence electrons”.
The valence electrons are the most important
electrons in the chemistry of the atom.

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The number of levels depends on the number
of electrons.
The first level (K) holds two electrons.
The second level holds eight electrons.
The third level holds 18, and the fourth 32.
No atom can have more than eight electrons in
its valence shell.
When the valence shell reaches eight electrons,
the next two electrons are put in a higher level.
Then the lower level can be filled.



Lewis dot structures show how many electrons
are in the valence shell of an atom.
Lewis dot structure for sodium
The first electron always goes to the right of the
symbol.
The second is paired with the first.

Lewis dot structure of magnesium
The third goes on top.
Lewis dot structure of aluminum

The fourth goes on the left, and is not paired.
The fifth goes on the bottom, and successive
electrons are paired until a total of eight is
reached.
Lewis dot structure of silicon
Lewis dot structure of oxygen


Bohr’s model based on atomic spectra
Obtaining emission atomic spectra

Energy is applied to a gas or liquid sample.
 Flame test (for samples in solution)
 Gas discharge tube


The energy makes an electron or two jump to a
higher energy level.
The electrons fall back down to a lower level, and
give off energy in the form of light – bright lines
against a dark background.


Absorption spectra – light is passed
through a sample and analyzed – looks
like a rainbow with dark lines
Interpreting atomic spectra
The light given off is viewed through a
spectroscope.
 The spectroscope has either a prism or a
grating, which splits the light into its
component colors.

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Only a few sharp lines appear in the
spectrum.
Each line corresponds to a specific
electron transition.
Transition = jump from one energy level
to another

Light energy travels in the form of
waves.

Color depends on frequency.
High frequency = violet end of spectrum
 Low frequency = red end of spectrum


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Energy also depends on frequency, so
each color has its own energy. Blue or
violet is higher energy than red or green.
When a specific color line is seen in a
spectrum, the energy of the electron
transition responsible can be calculated.

Bohr reasoned that since only certain
lines are seen in atomic spectra, only
certain energies must be allowed in
electron orbits.