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Electrons in Atoms
Chapter 5
Chemistry 11
Early Models of the Atom
Ancient Greeks were the first to come
up with the idea of atoms.
Democritus
suggested that all
matter was made of
tiny indivisible
particles called
atoms. (Greek
“atoma”)
Democritus
In the early 1800’s,
John Dalton came
up with the ATOMIC
THEORY.
His main points were:
Dalton’s Atomic Theory
1) All matter is made of atoms. Atoms
are indivisible and indestructible.
2) All atoms of a given element are
identical in mass and properties
3) Compounds are formed by a
combination of two or more different
kinds of atoms.
4) A chemical reaction is a
rearrangement of atoms.
crookes tube
J.J. Thomson's Experiments
Using Crooke’s tubes and other
equipment, J.J. Thomson discovered
the electron and measured its e/m
(charge to mass) ratio.
Later, “e” was found and the mass of an
electron was found to be 9.10938188 ×
10-28 grams (much lighter than H)
Thomson’s Plum
Pudding Model
Ernest Rutherford
Rutherford’s Scattering Experiment
Rutherford's Experiment
Rutherford could not explain why the
electron didn’t fall into the nucleus and
destroy the atom.
Neils Bohr
spectra and bohr
Bohr pictured the hydrogen atom as
having discrete energy “levels” which
the electron could “inhabit”. In it’s
ground state, the electron would be in
the lowest level (n=1)
When the atom was “excited” the
electron could “jump” to a higher level.
When the electron came back down, it
released energy in the form of light.
Each “jump” would give off light of a
particular wavelength or colour. This
gave rise to hydrogen’s spectrum.
According to Bohr, each energy “level”
corresponded to a different “orbit” of an
electron around the atom. (Like planets
around the sun.)
Bohr even calculated what the radii of
these orbits would be.
In the 1920’s things changed!
Although Bohr’s idea of energy levels
was still accepted, his idea of planetary
orbits for electrons was rejected!
REJECT ! !
The Quantum Mechanical Model (QMM)
• 1926- Austrian physicist Erwin Schrodinger used
the results of Rutherford and Bohr to devise
and solve a mathematical equation describing
the behavior of the electron in a hydrogen atom
• Unlike the Bohr model, the quantum mechanical
model does not involve an exact path the electron
takes around the nucleus
• The quantum mechanical model determines the
allowed energies an electron can have and how
likely (probability) it is to find the electron in various
locations around the nucleus
• The cloud is more dense where the probability of
finding an electron is high
• Atomic Orbital- a 3D region around the
nucleus describing the electron’s probable
location
Atomic Orbitals
• Energy Levels- are labeled by
• Principal Quantum Number (n)- 1, 2, 3, 4…
• Within each there are
• Energy Sublevels- the energy levels contained
with the principal energy level
Atomic Orbitals
Hydrogen’s First 4 Principal Energy Levels
Principal
Quantum
Number (n)
Sublevels
Number of
orbital's related
to sublevel
Max Number of
Electrons
1
s
1
2
2
s
p
1
3
8
3
s
p
d
1
3
5
18
4
s
p
d
f
1
3
5
7
32
Practice read pg 127-132
Qs 132 on page 1-7
Electron Arrangement in Atoms
• Electrons and the nucleus interact to make the
most stable arrangement possible.
• Electron Configurations- the ways in which
electrons are arranged in various orbitals
around the nuclei of atoms
3 Rules for electron configurations of atoms
1. Aufbau Principle : Electrons occupy the
orbitals of lowest energy first
2. Pauli Exclusion Principle: an electron orbital
may describe at most two electrons
– To occupy the same orbital, two electrons must
have opposite spins (↓or ↑)
3. Hund’s Rule- electrons occupy orbitals of the
same energy in a way that makes the number
of electrons with the same spin direction as
large as possible.
Noble Gas Configuration
• Method of representing electron configurations
of noble gases using bracketed symbols.
– Neon= [Ne]
• Also used to shorten electron configurations
– Sodium: #11- instead of 1s22s22p63s1 can be
shortened to [Ne] 3s1
Exceptions to predicted configurations
• Chromium- [Ar] 4s13d5
• Copper - [Ar] 4s13d10
• Illustrates the increased stability of half-filled
and filled sets of s and d orbital's
Valence Electrons (V.E.)
• Electrons in the atom’s outermost orbital's
• Determine the chemical properties of an
element
• V.E. are used in forming chemical bonds
Electron Dot Structures
• Consists of the element’s symbol and inner-level
electrons surrounded by dots representing the
atom’s valence electrons
• V.E. are placed one at a time on the four sides of the
symbol and then paired up until all are used
• Examples
• Chemists found Rutherford’s nuclear model to be
lacking because it did not begin to account for the
differences in chemical behavior among various
elements
• Early 1900’s- scientists observed that certain elements
emit visible light when heated in a flame chemical
behavior
Wave Nature of Light
• Electromagnetic Radiation- form of energy
that exhibits wavelike behavior as it travels
through space
Vocabulary to know..
• Wavelength- shortest distance between
equivalent points on a continuous wave
– Symbol- λ (lambda)
– Unit- meters, centimeters, or nanometers (1
nm= 1x10-9m)
• Frequency- the number of waves that pass a give
point per second
– Symbol- ν (nu)
– Unit- Hertz (SI Unit)= (1/s)= (s-1)
• Amplitude- the wave’s height from the origin to a
crest, or from the origin to a trough
How are they related?
•
ALL electromagnetic waves,
including visible light,
travel at a speed of
c= 3.00x108 m/s
(MEMORIZE)
C= λν
• Speed of light= wavelength x frequency
Electromagnetic Spectrum
• Aka EM Spectrum
• Encompasses all forms of electromagnetic
radiation
– The only differences in the types of radiation
being their wavelengths and frequencies
ROYGBIV
Calculations 
• Microwaves are used to transmit information.
What is the wavelength of a microwave having
a frequency of 3.44 x109 Hz?
– Know: C= λν
• C= 3.00x108 m/s
• ν = 3.44 x109 Hz
• λ = ???
λ = c/ ν
 
8
3.00 x10 m/s
9
3.44 x10 s
-1
λ= 8.72 x10-2 m
DON’T FORGET YOUR SIG FIG RULES!!!
Particle Nature of Light
• Quantum Concept
– Explained why colors of heated matter
correspond to different frequencies and
wavelengths
– Max Plank- “matter can gain or lose only in
small, specific amounts called quanta”
• Quantum- the minimum amount of energy
that can be gained or lost by an atom
• Energy of a quantum is related to the
frequency of the emitted radiation by the
equation
Equantum= hv
– E= energy
– h = Plank’s Constant (6.626x10-34J)
– v= frequency
– Joule (J)= SI unit for energy
Photoelectric Effect
• Electrons, called photoelectrons, are emitted
from a metal’s surface when light of a certain
frequency shines on the surface
– Photon- a particle of EM radiation with no mass
that carries a quantum of energy
Ephoton= hv
Atomic Emission Spectra
• Set of frequencies of the
electromagnetic waves emitted
by atoms of the element
– Example- The light of neon sign is
produced by passing electricity
through a tube filled with neon gas.
Neon atoms release energy by
emitting light.
An atomic emission spectrum is characteristic of the element being
examined and can be used to identify that element
Section 5.2
Bohr Model of the Atom
• Proposed that the hydrogen atom has only certain
allowable energy states
• Ground State- lowest allowable energy state of an
atom
• Bohr’s model worked well to explain
Hydrogen- however it did not explain other
elements
• Substantial evidence indicates that electrons
do not move around the nucleus in circular
orbits
• The de Broglie equation predicts that all moving
particles have wave like characteristics
λ= h/mv
• The Heisenburg uncertainty principal - states that it is
fundamentally impossible to know precisely both the
velocity and position of a particle at the same time