Download History of Atomic Theory

Development of
Atomic Models
Democritus

Greek philosopher

400 BC
Question

Is there a limit to
the number of
times matter could
be divided?
Democritus Theory

Eventually, you would
reach a piece that was
“indivisible”

Named this smallest
piece of matter “atomos,”
meaning “not to be cut.”
Atomos

Small, hard particles.

Differ in shape and size
for each substance
Aristotle and Plato

All matter made up of
combination of earth,
fire, air and water.
Aristotle
The Four Elements??

This concept influenced
early chemists called
alchemists.
Buried in History
“Atomos” theory was ignored
and forgotten for more than
2000 years!
John Dalton (early 1800’s)

Performed careful scientific
experiments.

Coined the term “atom”.
Dalton’s Atomic Theory

Matter is made of tiny
indivisible particles
called atoms.

Atoms of an element
are alike, and different
from atoms of other
elements.
Dalton’s Atomic Theory

Compounds are atoms of
different elements combined
in fixed proportions.

Chemical reactions involve
rearrangement of atoms.

Atoms cannot be created or
destroyed, but are conserved.
Pages from Dalton’s Journal
Hard Spheres
Dalton’s model is called the
“Hard Spheres Model”
JJ Thomson (1897)
Thomson’ Experiments

Studied “cathode
rays” (electric current)
in a “Crooke’s Tube”.

Fluorescent screen,
shows how ray
behaved in a
magnetic field.
Cathode Rays were
negatively charged
Cathode Rays were
particles

http://youtu.be/XU8nMKkzbT8
http://youtu.be/Z61zCaAFky4
 http://youtu.be/IdTxGJjA4Jw

JJ is Awesome

Concluded the
negative “cathode
ray” particles came
from within atoms.

Discovered the first
subatomic particle
(electron).
What about the Positive?

But…matter is neutral.

Must be a positive
charge in the atom to
balance the negative.
Plum Pudding Model

Positively charged
sphere with with
negatively charged
particles scattered
throughout.
Yummy…
Ernest Rutherford (1908)

Physicist who
worked with the
new field of
radioactive
emissions.
Different Types of Radiation

Used a magnetic field to
determine there were
three types of radiation.

Alpha (α)
Beta (β)
Gamma (γ)


Charges of Radiation

The radiation had different charges.
Identify the charge each type of radiation has.
Gold Foil Experiment

Shot alpha particles,
at a very thin piece of
gold foil.

These particles have
a positive charge

Fluorescent screen
shows where the
particles went.
 Observation:
Almost all alpha particles passed straight
through the gold foil.
 Conclusion:
Most of the atom’s volume is empty space.
 Observation:
A few alpha particles were deflected at an
angle or bounced back.
 Conclusion:
Atoms have a very small, dense positively
charged nucleus.
Nucleus is extremely small compared to
the size of the atom as a whole.
 Deflections happened rarely (1/8000).

The Nuclear Model
Rutherford’s Model is
called the
“Nuclear Model”
Comparison to Thomson

Positively charge
contained in nucleus.

Negatively particles
scattered outside
nucleus.

Not dispursed evenly.

http://chemmovies.unl.edu/ChemAnime/R
UTHERFD/RUTHERFD.html

http://youtu.be/wzALbzTdnc8

http://youtu.be/XBqHkraf8iE
Niels Bohr (1913)
 Came
up with the
“Planetary Model”
Bohr’s Theory

Electrons circle
nucleus in specific
energy levels or
“shells”.

The higher the
“energy level” the
higher the electron’s
energy.
Energy Levels

Different energy levels can contain
different numbers of electrons.
How many per level?

n = the number of the energy level
2
2n
= the total number of electrons
an energy level can
hold.
Ex: Level 3 can hold 2(3)2 = 18 electrons
Draw a Bohr Atom

Ex: The Fluorine Atom (F)
Protons = 9
 Neutrons = 10
 Electrons = 9

How many energy levels do you draw?
 How many electrons in each level?

Draw a Bohr Ion

They only difference is that one or more
electrons gets added or taken out of the
outer energy level.

Ex: The Magnesium Ion (Mg+2)
Protons = 12
 Neutrons = 12
 Electrons = 10

(+) Ions (cations)
(+) ions are smaller
Lost electron(s)
(-) Ions (anions)
(-) ions are larger
Gained electron(s)
How Did Bohr Come Up With His
Model?

Studied the spectral lines emitted by
various elements (especially Hydrogen)
What are Spectral Lines?

Energy gets absorbed by an atom causing it to
emit a unique set of colored lines.

Used to identify what elements are present in a
sample. (elemental “Fingerprint”)
Spectral Lines are Different for
Each Element
http://www.mhhe.com/physsci/chemistry/
essentialchemistry/flash/linesp16.swf
What Causes Spectral Lines?
Jumping Electrons!!
Jumping Electrons

Electrons normally exist in the lowest energy
level possible called the “ground state”. (stable)

“Ground state” e- configurations are written on the
periodic table for each element.

Ex:
Aluminum is 2-8-3
Calcium is 2-8-8-2
An Electron Gets “Excited”
Electrons can absorb a photon of energy and
“jump up” to a higher energy level farther from the
nucleus.
This is called the “excited state”. (unstable)
Jumping Electrons

They quickly “fall back down” to the
ground state. (stable)

They emit a photon of energy that
corresponds to how far they jumped.

This photon of energy is seen as a
spectral line!

Each spectral line corresponds to a
specific photon of energy that is released.
REMEMBER
Absorb Energy
Jump Up
Emit Energy
Fall Down
Electromagnetic Spectrum

Spectral lines can
come from all areas
of the EM Spectrum.

Lines of visible colors
make up only a small
part of the spectrum.

EM waves carry different amounts of energy
based upon their wavelength and frequency.
Which wave has higher energy?

http://www.upscale.utoronto.ca/PVB/Harris
on/BohrModel/Flash/BohrModel.html
Calculating the Energy of a
Spectral Line
STEP 1:
If you know the wavelength of the spectral line you can
find it’s frequency.
c=λxү
c = the speed of light = 3 x 108 meters/sec
λ = wavelength (in meters)
ү = frequency of the wave
Calculating the Energy of a
Spectral Line
STEP 2:
Using the frequency find the energy of the line (in Joules)
E=hxү
E = energy in Joules
h = Planck's constant = 6.63 × 10-34 kg x m2 / sec
ү = frequency of the wave
Electron Cloud Model
Electron Cloud Model

Sometimes called:




Wave Mechanical Model
Quantum Mechanical Model
Orbital Model
Charge Cloud Model
How is it Different from the
Planetary Model?
Heisenberg’s Uncertainty
Principle:
It is impossible to know the
exact location and momentum of
an electron at the same time.

We can’t tell exactly where an
electron is!!

Electrons exist in
“orbital clouds”

The denser the
region of the cloud
the higher the
probability of
finding an electron
there.

http://youtu.be/45KGS1Ro-sc
How are Electrons Organized?
Electron
Hotel
Energy Levels (1-7)

Electrons can be at different distances
from the nucleus.
 Energy
1
Levels
2 3
Lowest energy
Closest to nucleus
4
5
6
7
Highest energy
Farthest from Nucleus
Sublevels (s, p, d, f)

Each energy level can have a certain number of
sublevels.
Energy Level Sublevels Possible
1
s
2
s, p
3
s, p, d
4
s, p, d, f
5
s, p, d, f, (g)
6
s, p, d, f, (g, h)
7
s, p, d, f, (g, h, i)
Energy of Sublevels

Sublevels have different levels of energy.
s
Lowest energy
p
d
f
Highest energy
Orbitals in Sublevels

Each sublevel contains a different number of orbitals.

A maximum of two electrons can exist in an orbital.
Sublevel
s
p
d
f
# of Orbitals
1
3
5
7
Max e- in Sublevel
2 e6 e10 e14 e-
Electron Spin

Pauli Exclusion Principle:
In order for two electrons to
occupy the same orbital, they
must have opposite spins.

Electrons in an orbital spin
in opposite directions
Shapes of Orbitals

Orbitals come in different shapes and
sizes.

They are the region of highest probability
of finding an electron.
s Orbital
Probability cloud has a spherical shape
p Orbitals (px, py, pz)
“Dumbell”
shape
Three p orbitals can
exist, on the x, y, z
axis in space
d Orbitals

Five possible d orbitals exist
f Orbitals

Seven possible f orbitals exist