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Modern Atomic
Model
Electron modeling…
• To understand electrons,
scientists began comparing them
to light.
Behavior of light
•Light is a wave – similar to
water waves
•Visible light belongs to
electromagnetic spectrum
Spectrum
High
Low
energy
energy
Radio Micro Infrared
Ultra- XGamma
waves waves .
violet Rays Rays
Low
High
Frequency
Frequency
Long
Short
Wavelength
Wavelength
Visible Light
Behavior of light
• All forms of EMR travel at a
constant speed of 3.0 x 108 m/s,
when in a vacuum. This value
also works for the speed of light
in air (because air is mostly a
vacuum).
Behavior of light
Waves can be described in terms of
the following:
•Wavelength
•Frequency
•Amplitude
•Speed
Behavior of light
Waves can be described in terms of the following:
•Wavelength - , distance
between successive crests (or
any 2 corresponding points),
for visible light  = 400-750
nm
Behavior of light
• Frequency - the number of waves that
pass a given point per second, units
are cycles/sec or hertz (Hz),
abbreviated n - the Greek letter nu
c = n
Behavior of light
•Amplitude – height from
origin to crest
Behavior of light
•Speed – measured in m/s,
light moves @ constant speed
of 3.0 x 108 m/s, abbreviated
as c
Parts of a wave
Crest
Wavelength
Amplitude
Origin
Trough
Parts of Wave
•
•
•
•
•
Origin - the base line of the energy.
Crest - high point on a wave
Trough - Low point on a wave
Amplitude - distance from origin to crest
Wavelength - distance from crest to crest,
abbreviated
 (Greek letter lambda)
Behavior of light
• Because light has a constant
speed, we get a relationship
between n &  (c = n)
Frequency and wavelength
• Are inversely related
– As one goes up, the other goes down.
• Different frequencies of light go with
different colors of light.
• There is a wide variety of frequencies
– The whole range is called a continuous
spectrum
Behavior of light
Question 1: If light has  = 633 nm,
what is n?
Question
2:
Red
light
travels
at
3.0
x
108 m/s and has a  of 700 nm.
What is n?
Question
3:
Violet
light
has
a
n
of
7.5
x 1014 Hz. What is ?
Wave model problems
The wave model of light worked
well until the beginning of the
20th century. This is because
some scientists were observing
light and found that what they
saw did not fit the wave model.
Wave model problems
Black body radiation
• In 1900, Max Planck was studying
radiation given off when matter
was heated. The physics he knew
said that matter could absorb or
emit any quantity of energy. The
results of his experiments did not
fit with that idea.
Light is a Particle
• Energy is quantized.
– Light is energy
– Light must be quantized
– These smallest pieces of light are called
photons.
• Energy and frequency are directly related.
Wave model problems
•A quantum of light was
later called a photon.
Radiation is emitted or
absorbed in whole
numbers of photons.
Wave model problems
• To relate the quantum of
energy and the frequency
of the radiation, he created
the relationship E = hn.
Energy and frequency
• E=hxn
– E is the energy of the photon
– n is the frequency
– h is Planck’s constant
• h = 6.626 x 10 -34 Joules × seconds
Wave model problems
• What energy is given off when
your stove coils turn red?
(Remember that red light has a
14
frequency of 4.29 x 10 Hz.)
• Which has greater energy – red
or violet light?
Wave model problems
• Planck’s ideas were not
immediately accepted. It was not
until some time later that Albert
Einstein used Planck’s equation to
work on solving the photoelectric
effect.
Wave model problems
 Photoelectric effect
• Light shining on certain metals
can eject electrons.
Wave model problems
 Photoelectric effect
• The fact that light was able to
knock electrons loose wasn’t a
problem. What wave theory
couldn’t explain was why only
certain frequencies of light (or
higher) could knock out
electrons.
• Photoelectric Effect Simulation
Wave model problems
 Photoelectric effect
• Einstein proposed that light
consisted of energy quanta that
behaved as particles – not waves.
The quanta were called photons.
Wave model problems
 Photoelectric effect
• The photoelectric effect problem
was then solved by the idea that
radiation is emitted or absorbed
in whole numbers of photons or
radiation particles.
Wave model problems
 Photoelectric effect
• It was later proven that light
could definitely act as a particle.
So, we now have light acting as
both a wave and as particles.
(This will be the basis for
understanding how e- behave.)
Atomic Spectrum
What color tells us about atoms
Prism
• White light is made up
of all the colors of the
visible spectrum.
• Passing it through a
prism separates it.
If the light entering the prism is
not white…
• By heating a gas or
using electricity, we
can get the gas to give
off colors
• Passing this light
through a prism does
something different
than white light
Atomic Spectrum
• Each element gives off
its own characteristic
colors
• Can be used to
identify the element
• How we know what
stars are made of
Wave model problems
 Bright line spectrum
• Scientists noticed that you could
vaporize an element in a flame to
produce different flame colors.
You can then use a prism to sort
the colors to produce a line
spectrum (only certain colors are
produced).
Wave model problems
 Bright line spectrum
• Problem: Each element produced
a different line spectrum.
• These are called
line spectra
• They are
unique to each
element.
• These are
emission
spectra (the
light is emitted
or given off)
An explanation of Atomic Spectra
Rutherford’s Model
• Discovered the
nucleus
• Small dense and
positive
• Electrons around
nucleus in electron
cloud
Bohr’s Model
• Why don’t the electrons fall into the
nucleus?
• Move like planets around the sun.
• In circular orbits at different levels.
• Energy separates one level from another.
Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
Bohr’s Model
Increasing energy
Fifth
Fourth
Third
Second
First
Nucleus
Further away
from the nucleus
means more
energy.
There is no “in
between”
energy
Energy is in
Levels
Bohr Model
• Niels Bohr was able to explain the
bright line spectrum. To do so, he
created a model with the
following parts:
•The e- could orbit the nucleus
only in allowed paths or
orbits.
Bohr Model
•The H atom has set energy
possibilities that depend on
which orbit the e- occupies.
•The ground state occurs when
the e- is in the orbit closest to
the nucleus.
Bohr Model
•The orbit where the e- is
determines the outer
dimensions of the atom.
•The energy of the e- increases
as it moves into orbits that
are farther and farther from
the nucleus (excited atom).
Where the electron starts
• The energy level an electron starts from
is called its ground state.
Changing the energy
• Let’s look at a hydrogen atom
Changing the energy
• Heat or electricity or light can move the
electron up energy levels
Changing the energy
• As the electron falls back to ground state it
gives the energy back as light
Changing the energy
• May fall down in steps
– Each with a different energy, frequency, and
wavelength
The Bohr Ring Atom
n=4
n=3
n=2
n=1
The Bohr Ring Atom
• The farther the electrons fall, the more
energy released and higher frequency
produced
• All the electrons can move
Bohr Model
• Bohr said that the energy of an
electron is quantized (like light)
so there have to be energy levels
(orbits) where the e- can be. The
e- must be given a certain
amount of energy to “jump”
orbits.
Bohr Model
• If the electron is in the lowest
energy level possible (closest to
nucleus), that is called the
ground state.
Bohr Model
•If energy is put into the e-, it
goes to a higher level or an
excited state. The spectral
lines produced were due to
the energy given off when the
e- fell.
Bohr Model Problems
• Unfortunately, Bohr’s model only
worked for H. What it did do was
to get other scientists thinking.
De Broglie
•Determined that particles of
matter could act as waves.
•Described the wavelength of
moving particles.
Conclusion:
Matter exhibits both wave and
particle properties!
Where are the electrons?
• We know they are outside of the nucleus.
• We say they are in an electron cloud.
• But where?
Quantum-mechanical model
Takes into account the following:
• Treats e- as waves within the
atom.
Quantum-mechanical model
• e- inside of atoms have specific E
and occupy 3-D regions about the
nucleus called orbitals. [Orbitals
are different than orbits.]
Quantum-mechanical model
• The size and shape of the orbitals
depends on the E of the e- that
occupy them.
• All orbitals in an atom make up the
e- cloud around the nucleus. The ecloud gives the atom a size and
shape.
• The e- can’t be located exactly in the
atom. There are areas of probability
to find an e-.
The Quantum Mechanical Model
• Has energy levels
for electrons.
• Contains orbitals of varying shapes and
sizes
• It can only tell us the probability of finding
an electron a certain distance from the
nucleus.
The Quantum Mechanical Model
• The electron is found
inside a blurry
“electron cloud”
– An area where there
is a chance of finding
an electron.
Atomic Orbitals
• Principal Quantum Number (n) = the
energy level of the electron.
• Within each energy level complex math
describes several shapes.
– These are called atomic orbitals
Summary of atomic orbitals
# of
Max
shapes electrons
Starts at
energy level
s
1
2
1
p
3
6
2
d
5
10
3
f
7
14
4
By Energy Level
• First Energy Level
• only s orbital
• only 2 electrons
• Second Energy Level
• s and p orbitals are
available
• 2 in s, 6 in p
• 8 total electrons
By Energy Level
• Third energy level
• s, p, and d orbitals
• 2 in s, 6 in p, and 10
in d
• 18 total electrons
• Fourth energy level
• s,p,d, and f orbitals
• 2 in s, 6 in p, 10 in d,
and 14 in f
• 32 total electrons
By Energy Level
• Any thing past the fourth level - not all
the orbitals will fill up.
–You simply run out of electrons
• The orbitals do not fill up in a neat
order.
Filling order
• Lowest energy fill first.
• The energy levels overlap
• The orbitals do not always fill up order of
energy level.
Increasing energy
7s
6s
5s
7p
6p
6d
5p
5d
4s
4p
4d
3s
3p
3d
2s
2p
1s
5f
4f
Electron Configurations
• The way electrons are arranged in atoms.
• Aufbau principle - electrons enter the lowest
energy first
– This causes difficulties because of the overlap of
orbitals of different energies.
• Pauli Exclusion Principle - at most 2 electrons
per orbital with different spins
• Hund’s Rule - When electrons occupy orbitals
of equal energy they don’t pair up until they
have to
Electron Configurations
Notations of e- in atoms:
•
Orbital diagram - unpaired e- represented
by  or  , paired e- shown as .
– Write orbital diagrams for elements 1-10.
•
e- configuration notation - no more lines
& arrows, uses # of e- in a sublevel as a
superscript over the sublevel designation
– Write electron configuration notation
for elements 1-10.
Electron Configurations
Notations of e- in atoms:
• Shortcut for orbital diagrams &
e- configuration notation
– Uses noble gas (group 18
elements) “core”
Noble Gas Shortcuts…
• 1. Find element on periodic table.
• 2. Move up 1 row on table and go to Noble
Gas at end of that row
• 3. Put this noble gas symbol inside [ ].
• 4. Now, write out what is left over after the [ ].
Examples
• Li =1s2 2s1
– Noble gas in row above is He
– [He] 2s1 is the same as 1s2 2s1
• Be = 1s2 2s2
– Noble gas in row above is He
– [He] 2s2 is the same as 1s2 2s2
• Na = 1s2 2s2 2p6 3s1
– Noble gas in row above is Ne
– [Ne] 3s1 is the same as 1s2 2s2 2p6 3s1
Electron Configurations
Questions:
1)
2)
3)
4)
5)
6)
7)
8)
9)
10)
11)
How many sublevels are found in the 3rd energy level?
If there were an 8th energy level, how many sublevels would it have?
How many orbitals are in the s sublevel? p sublevel? d sublevel? f
sublevel?
If there were a sublevel past the f sublevel, how many orbitals would it
have?
How many orbitals are in the 2nd energy level?
How many orbitals are in the 4th energy level?
How many orbitals would be in the 6th energy level?
How many electrons are able to go into an orbital?
How many electrons would there be in the s sublevel? p sublevel? d
sublevel? f sublevel?
How many electrons would you find in the 1st energy level? 4th energy
level?
If we had a 5th energy level, how many electrons would it have?
Periodic Table
History
History of the Periodic Table
Dobereiner
• Triads of elements with shared properties
Cannizzaro
• method for measuring atomic masses and
interpreting the results of measurements
Newlands
• arranged elements by atomic masses,
properties repeated after every 8 elements
→ law of octaves
History of the Periodic Table
Mendeleev & Meyer
• arranged elements according to
the increase in atomic mass
History of the Periodic Table
Mendeleev
• left spaces for undiscovered
elements & predicted properties
of those elements
• credited with discovering
periodicity
History of the Periodic Table
2 Questions:
1. Why could most elements be
arranged by increasing atomic
mass, but a few could not?
2. What was the reason for
chemical periodicity?
History of the Periodic Table
Mosely
• shooting electrons at various metals to
produce X-rays
• frequencies of the X-rays were unique to
the metals
• assigned a whole number to each
element → atomic numbers
» arrange elements by atomic
numbers to get families with similar
properties
History of the Periodic Table
Periodic Table
• an arrangement of the elements
in order of their atomic numbers
so that elements with similar
properties fall in the same column
(group/family)
History of the Periodic Table
Periodic Law
• the physical and chemical
properties of elements are
periodic functions of their atomic
numbers
Periodic Table
Information
Types of Elements
Periodic Table Information
Types of elements:
• Metal – lustrous, good conductors, most are
solids, malleable, ductile
• Nonmetal - poor conductors, no luster,
neither malleable nor ductile, most are
gases, wide variety of other physical
properties
• Metalloid - properties of metals & nonmetals
Parts of the
Periodic Table
Periodic Table Information
Period
• horizontal rows of elements
Group (family)
• vertical columns of elements
Periodic Table Information
Group (family)
•
•
•
•
•
•
•
alkali metals – group 1 (except hydrogen)
alkaline earth metals – group 2
transition metals – all of the d-block elements
inner transition metals – all of the f-block elements
metalloids – elements that touch the “staircase”
halogens – group 17
noble gases – group 18 (elements with filled outer
shells of electrons)
Back to Electrons…
(electron configurations)
Electron
Configuration
Notation
(using shortcut)
More Electron Configurations
Practice writing shortcut econfigurations.
• Element # 15
• Element # 8
• Element # 34
Group Electron
Configuration
More Electron Configurations
Group e- configurations – all elements in
family have similar “endings”
Write the shortcut e- configuration for:
• Li, Na, K
»all end with __, so group
configuration is __
• C, Si, Ge
»all end with __, so group
configuration is __
More Electron Configurations
Practice identifying elements using group
configurations:
• Element in period 3 with group
configuration p4
• Element in period 6 with group
configuration s2
• Element in period 2 with group
configuration p6
• Element in period 7 with group
configuration s1
Periodic Trends
Periodic Trends
Electron configurations are able to
cause periodic variations in
elemental properties...
Periodic Trends
Valence electrons - electrons that may
be lost/gained/shared when
chemical compounds are formed
• Period  As we go across a period,
the number of valence electrons
increases.
• Group  As we go down a group, we
find that the number of valence
electrons stays constant.
Periodic Trends
Size of the atomic radius – one-half the
distance between the nuclei of identical
atoms joined in a molecule
• Period  As we go across a period,
the general trend is for the atomic radii
to decrease.
• Group  As we go down a group,
there is a general increase in atomic
radii. This happens because we are
seeing more and more energy levels
being added.
Periodic Trends
Ionization energy – energy required to
overcome nuclear attraction and remove an
electron from a gaseous element
• Period  As we go across a period,
the general trend is for the ionization
energy to increase.
• Group  As we go down a group,
there is a general decrease in
ionization energy.
Periodic Trends
Electronegativity - measure of the power
of an atom in a chemical compound to
attract electrons
• Period  As we go across a period,
the general trend is for the
electronegativity to increase.
• Group  As we go down a group,
there is a general decrease in
electronegativity.