Download PPT #2

Development of
Atomic Models
Democritus

Greek philosopher

400 BC

“Atomos” concept

Matter can’t be divided
forever

Eventually, a piece
would be “indivisible”

“Atomos,” meaning “not
to be cut”
John Dalton (early 1800’s)

Coined the term “atom”.
Dalton’s Atomic Theory

Matter made of tiny
indivisible particles
called “atoms”.

Atoms of one element
different from atoms of
other elements.
Page from Dalton’s Journal
Dalton’s Theory Continued…

Compounds form when atoms
combine in fixed proportions.

Chemical reactions involve
rearrangement of atoms.

Atoms are conserved in
chemical reactions.
Dalton’s Atomic Theory called
“Hard Spheres Model”
JJ Thomson (1897)
Thomson’ Experiments

Studied “cathode
rays” (electric current)
in a “Crooke’s Tube”.

Fluorescent screen,
shows how cathode
ray behaved in a
magnetic field.
Lets draw a typical Crooke’s
Tube in our notes.
Cathode Rays were
negatively charged
They bent toward (+) plate
Cathode Ray Tube and Magnet
http://youtu.be/XU8nMKkzbT8
Cathode Rays
were particles
They couldn’t pass through matter.
Subatomic Particles Exist!

Concluded
“cathode ray”
particles came from
within atoms.

Discovered first
subatomic particle

negative electron
What about the Positive?

But…matter is neutral.

Therefore:
 A positive charge
must exist to balance
the negative.
Plum Pudding Model (Thomson)
Atoms are positively charged spheres with
negatively charged particles scattered throughout.
Yummy…
Brian Cox:
Thompson and Discovery of Electron
http://youtu.be/IdTxGJjA4Jw
Ernest Rutherford (1908)

Physicist who
worked in new field
of radioactivity.
Found 3 Different Types of
Radiation

Used magnetic field to
isolate 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 thin piece of gold foil.

Alpha particles have
positive charge, and
mass of 4 amu

Fluorescent screen
shows where particles
went.
Rutherford Gold Foil (45 sec)
http://www.youtube.com/watch?v=5pZj0u_XMbc
Observation:
Most alpha particles passed straight
through gold foil.
Conclusion:
Atom’s volume is mostly empty space.
Observation:
Some alpha particles
deflected at an angle
or bounced back.
Conclusion:
Atoms have a very
small, dense positively
charged nucleus.
http://www.kentchemistry.com/moviesfiles/Units/AtomicStructure/Ruthe
rford3.htm
Nucleus is extremely
small compared to the
size of the atom as a
whole.
Deflections happened
rarely (1/8000).
Modern Example of Gold Foil Experiment in Action
http://youtu.be/XBqHkraf8iE
Nuclear Model
Rutherford’s Model is called the “Nuclear Model”
Brian Cox: Rutherford and the Nucleus
http://youtu.be/wzALbzTdnc8
Comparison to Thomson

Positive charge only
contained in nucleus.

Negative particles
scattered outside
nucleus.

Charge is not
disbursed evenly.
Simulator:
https://phet.colorado.edu/sims/html/rutherfordscattering/latest/rutherford-scattering_en.html
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 Electrons Per Level?

n = the number of the energy level
2
2n
= maximum 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?

Human Bohr Model
http://www.youtube.com/watch?v=PLpZfJ4rGts
Draw a Bohr Ion

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

Questions on Bohr Model

What is similar about the electron configuration of atoms
in same group? In same period?

What is the general trend in atomic radius as you go
down a group? Why?

What happens to the size of an atom when it becomes a
negative ion? A positive ion?

Why might the outermost electrons be of most interest to
chemists?
(+) Ions (cations)
are smaller
Lose electron(s)
(-) Ions (anions)
are larger
Gain electron(s)
How Did Bohr
Come Up With His Model?

Studied spectral lines emitted by various
elements (especially Hydrogen)
What are Spectral Lines?

Energy absorbed by an atom causes it to emit a
unique set of colored lines.

Used to identify elements present in a sample.
(elemental “Fingerprint”)
Spectral Lines are Different for
Each Element
Answer: 1
What Causes Spectral Lines?
Jumping Electrons!!
Video of Line Spectra of Hydrogen
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/linesp
16.swf
Jumping Electrons

Electrons normally exist in lowest energy level
possible called “ground state”. (stable)

“Ground state” e- configurations are written on
periodic table

Ex:
Aluminum is 2-8-3
Calcium is 2-8-8-2
An Electron Gets “Excited”
Electrons absorb a photon (or “quanta”) of energy
and “jump up” to a higher energy level farther
from nucleus.
This is called “excited state”. (unstable)
Jumping Electrons

Electrons quickly “fall back down” to
ground state. (stable)

They emit a photon (or “quanta”) of energy
that corresponds to how far they jumped.

Each spectral line corresponds to a
specific photon of energy that is released.
Model Of Hydrogen Atom and Electrons Jumping
http://www.upscale.utoronto.ca/PVB/Harrison/BohrModel/Flash/BohrModel.html
REMEMBER
Absorb Energy
Jump Up
Emit Energy
Fall Down
Excited vs Ground State

Periodic table lists ground state electron
configurations for neutral atoms.


To recognize an “excited state” configuration,
count the electrons and see if the
configuration matches the one on the table.
Ex: 2-8-7-3 = 20 electrons
Calcium (atomic # 20) is 2-8-8-2
 So this must be showing one of the ways
calcium could be in the excited state.

Valence Electrons

Electrons in highest occupied energy level.

Involved in forming bonds with other atoms.

Atoms are most stable when they have a
“stable octet” of 8 valence electrons

Noble Gases: (Group 18)

“Inert” and unreactive (have stable octet)
 Ex:
Argon 2-8-8, Neon 2-8
Valence Electrons

Look at last number in atom’s electron
configuration to determine number of
valence electrons.

Ex:
Al
 Ca
F

2-8-3
2-8-8-2
2-7
3 valence
2 valence
7 valence
Lewis Dot Diagrams

Shows number of valence electrons an
atom has as “dots” around atom’s symbol.
Phosphorus is 2-8-5
Kernel
Nucleus and non-valence electrons
 Inner part of atom not involved directly in
reactions


Ex:

Al
2-8-3
has 10 kernel electrons
and 3 valence electrons
The Nature of Light

Dual Nature of Light:
 behaves as both waves and as particles
(depending on what type of experiment is being
performed.)

Speed of Light: light waves travel at same velocity
 C = 3.0 x 108 meters/sec




What is Light?
https://www.youtube.com/watch?v=eCVPhjHh57E
Greatest Discovery in Physics: (Duality of Light)
https://www.youtube.com/watch?v=XB-iLRsq8A8
Electromagnetic Spectrum

Spectral lines can
come from all areas
of EM Spectrum.

Visible colors make
up only a small part

EM waves carry different amounts of energy
based upon their wavelength and frequency.
Wavelength (λ): distance
between two peaks of a wave
Frequency (γ): number of
peaks that pass per second.
(Hertz (Hz) or cycles/sec)
Which wave has higher energy?
Relationship of Frequency, Wavelength and Energy of colored line
http://employees.oneonta.edu/viningwj/sims/plancks_equation_s.html
Crash Course: Atomic Basics (history)
https://www.youtube.com/watch?v=FSyAehMdpyI
Calculating Energy of a Spectral Line
(HONORS)
STEP 1:
Given wavelength of spectral line find it’s frequency.
c=λxү
c = the speed of light = 3 x 108 meters/sec
λ = wavelength (in meters)
1 x 10-9 meter = 1 nm
1 x 10-10 meter = 1 Angstrom
ү = frequency of the wave (Hertz or waves/sec, s-1)
Calculating Energy of a Spectral Line
(HONORS)
STEP 2:
Using frequency find energy of the line (in Joules)
E=hxү
E = energy (Joules)
h = Planck's constant = 6.63 × 10-34 kg x m2 / sec
ү = frequency of the wave (Hertz or waves/sec, s-1)
Calculating Energy for Specific
Jumps Between Energy Levels
(For Hydrogen)
∆ E = Efinal - EInitial
-2.180 x 10-18 J/en2
Final Energy level
-
-2.180 x 10-18 J/en2
Initial Energy level
Honors Good Overview Video


What is Light: Lecture on Photon emission and absorbtion (2 hours)
https://www.youtube.com/watch?v=axUkUuj6aus