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Chapter 5
Electrons in Atoms
 Light and quantized energy
Chapter 5
Arrangement of Electrons
I. Electromagnetic Waves
II. Dual Nature of Light
III. Bohr Model of the Atom
IV. Quantum Model
V. Quantum Numbers
VI. Determining Number of Orbital
Types and Electrons
VII. Electron Configurations
I. Electromagnetic Waves
A. Definition of a Wave
1. method by which energy is transferred
from one point to another
B. Definition of Electromagnetic Wave
1. a form of energy that exhibits wavelike behavior as it travels through
space
I. Wave Particles Nature of
Light
A. Electromagnetic radiation- form of energy
that exhibits wavelike particles.
Includes many kinds of waves
All move at 3.00 x 108 m/s ( c) speed of light
1. Light is a kind of electromagnetic radiation.
a. The study of light led to the
development of the quantum mechanical
model.
B. Parts of Wave
1.Origin - the base line of the energy
2.Crest - high point on a wave
3. Trough - Low point on a wave
4. Amplitude - distance from origin to crest
5. Wavelength - distance from crest to crest
6. Wavelength - is abbreviated l Greek letter
lambda.
Parts of a wave
Crest
Wavelength
Amplitude
Origin
Trough
C.Properties of Electromagnetic Waves
1. can travel in a vacuum
2. travel at 3 x 1010 cm per second
(this is the speed of light)
3. vary in wavelength and frequency
a. wavelength – distance between
corresponding points on waves
b. frequency – the number of waves
that pass a point in a given amount
of time (usually one second)
Electromagnetic Wave
Disturbance in a magnetic field is perpendicular to a
disturbance in an electric field
C. Frequency
1. The number of waves that pass a given
point per second.
2. SI units are hertz (hz) or cycles/sec
3. Abbreviated n the Greek letter nu
c = ln
C. Frequency and wavelength
watch slinky
1. Are inversely related
2. As one goes up the other goes down.
3. Different frequencies of light are different
colors of light.
4. There is a wide variety of frequencies
5. The whole range is called a spectrum
EMS
High
Low
energy
energy
Radio Micro Infrared
Ultra- XGamma
waves waves .
violet Rays Rays
Low
High
Frequency
Frequency
Long
Short
Wavelength
Wavelength
Visible Light
D. Examples of Electromagnetic Waves
1. radio waves
2. microwaves
3. infrared
4. white light (visible spectrum)
5. ultraviolet light
6. X-rays
7. gamma radiation
Continuous Electromagnetic Spectrum
---------> increasing wavelength ----------->
E. Wavelength, Frequency and Energy
1. wavelength and frequency
a. the frequency increases as the
wavelength decreases
b. inverse relationship
2. frequency and energy
a. as the frequency increases the
energy increases
b. direct relationship
Electromagnetic Spectrum
----------------> increasing frequency
---------------> decreasing wavelength
---------------->
---------------->
Diagram Showing Wavelength and Frequency
F. Types of Spectra
1. continuous – all wavelength within a
given range are included
2. electromagnetic – all electromagnetic
radiation arranged according to
increasing wavelength
a. unit for wavelength ranges from
meters to nanometers
b. unit for frequency is hertz (Hz)
3. Visible spectrum - light you can see
(ROY-G-BIV)
a. red has the longest wavelength and
the smallest frequency
b. violet has the shortest wavelength
and the greatest frequency
4. Bright Line spectrum (emission spectrum)
a. bands of colored light emitted by
excited electrons when they return to
the ground state
G. Spectroscopy
1. emission spectra of a substance is
studied to determine its identity
2. spectroscope – instrument that
separates white light into a spectrum
3. spectral lines – represent wavelength
of light emitted when excited electrons
fall back to the ground state
Emission Spectrum (Line Spectrum)
Picture of a Spectroscope
Emission Spectrum
Spectral line activity
 Put on spectrum glasses
 View Hydrogen argon, helium,etc.
II. Light Has a Dual Nature!!!
A. Light can act like a particle or a wave
1. emission and absorption of light by
matter can not be explained by wave
theory
2. only certain frequencies of light
produce the photoelectric effect
a. emission of electrons by some
metals when they are exposed to
light
II. 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.
Light is a particle

 Missing the video
Light Has a Dual Nature (Particle + Wave)
Light Interference Pattern (Wave Nature)
3. In 1900 Max Planck observed that a hot
object loses energy in packets called
quanta
a. this energy is directly related to the
wave frequency ( E = hv)
b. in 1905 Einstein said this relationship
held for all electromagnetic radiation
Photoelectric Effect – Particle Nature
Light hits a metal and electrons are released and an
electric current may be produced
II. Light is a particle
A. photoelectric effect –
emission of electrons
by metals when light
shines on them
( must be a specific
frequency)
Which has more energy a
marble or a bowling ball?
A marble can’t knock down a
block no matter how
many times you throw it.
Photoelectric Effect – Particle Nature of Light
Only light of a certain frequency or higher
will cause the photoelectric effect
4. Vocabulary
a. quantum – quantity of energy gained
or lost by an atom when electrons are
excited
b. photon – a quantum of light
c. ground state – lowest energy level of
an atom
d. excited state – a heightened state of
energy in an atom
Energy and frequency






E=hxn
E is the energy of the photon
n is the frequency
h is Planck’s constant
h = 6.6262 x 10 -34 Joules sec.
joule is the metric unit of Energy
The Math in Chapter 5




Only 2 equations
c = ln
E = hn
Plug and chug.
Examples
 What is the wavelength of blue light with
a frequency of 8.3 x 1015 hz?
 What is the frequency of red light with a
wavelength of 4.2 x 10-5 m?
 What is the energy of a photon of each of
the above?
Atomic Spectrum
How color tells us about atoms
The Flame Test
A basic form of Emission
Spectroscopy
Prism
 White light is made
up of all the colors of
the visible spectrum.
 Passing it through a
prism separates it.
If the light is not white
 By heating a gas with
electricity we can get
it to give off colors.
 Passing this light
through a prism does
something different.
Atomic Spectrum
 Each element gives
off its own
characteristic colors.
 Can be used to
identify the atom.
 How we know what
stars are made of.
• These are called
discontinuous
spectra
• Or line spectra
• unique to each
element.
• These are
emission spectra
• The light is
emitted given off.
You and Your Partner
 Label each splint, by metal, take two splints
for each metal
 Dip the wet splint in the salt solutions
 Insert wooden splint at tip of inner
cone…do not let it burn
 Record color of flame
The Metals





Sodium - NaCl
Potassium - KCl
Strontium - Sr(NO3)2
Lithium - LiNO3
Calcium - CaCl2
 Unknown: A, B, C or D
Emission Spectroscopy
 Technique used to identify unknown
elements in a sample
Basis of Test
 Electrons in the ground state get excited
when energized
 Excited electrons are unstable
 Electrons fall back down to the ground state
by releasing energy
 Energy takes the form of visible light
Line Spectrum
 Characteristic wavelengths (colors) of light are
given off by elements
 These wavelengths are an elements line spectrum
 Hydrogen
434
410 nm
nm
656 nm
486
nm
Flame Test
 Used to identify metals in solution
 Electrons absorb energy from the flame to
enter the excited state
Safety
 Goggles and aprons
 Double cone flame
 Garbage in tin can
partially filled with
water
 Wash hands and lab
station
Bohr’s Model
 Electrons move like planets around the
sun.
 In circular orbits at different levels.
 Amounts of energy separate one level
from another.
III. The Bohr Model of the Atom
A. Electrons of hydrogen circle the
nucleus in orbits
1. orbits have a fixed amount of energy
in the ground state
2. orbits are a fixed distance from the
nucleus
3. orbits furthest from the nucleus have
the greatest energy
4. Electrons in the ground state can absorb
quanta of energy – become excited- and
move to a higher orbit
5. Electrons emit quanta of energy when
they return to the ground state
6. Model applies only to hydrogen atoms
Niels Bohr
(1885 – 1962)
Bohr Model of the Atom
Bohr
c.The electron must be in one orbit or
another – it cannot be in between- the
energy is quantized
d. Line spectrum- produced when an
electron drops from a higher energy orbit
to a lower energy orbit
i. A photon is emitted with energy E=hv
equals difference in energy between
the initial higher level and the final
lower orbit
I. BOHR

A. Niels Bohr(1885-1962) – Danish
physicist- worked with Rutherford
1. Electron circles the nucleus in orbits
a. The closer the orbit to the nucleus the lower the energy
level.
b. The total energy of the electron increases as the distance
from the nucleus increases
Lyman, Balmer, Paschen Series for Hydrogen
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 Levels
Bohr
 Make a model of Bohr’s Hydrogen. Bohr
was only correct about Hydrogen.
 Draw a nucleus
 Use the radius in book move decimal one
space to right
 Measure in centimeters
 Use markers to illustrate the excited vs
ground state electrons.
Where the electron starts
 The energy level and electron starts from is
called its ground state.
 As it absorbs energy it goes up to an excited
state energy level.
 Was each level equally distant from the
other?
 Then what happens?
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
Ultraviolet
Visible
Infrared
 Further they fall, more energy, higher
frequency.
 This is simplified
 the orbitals also have different energies inside
energy levels
 All the electrons can move around.
Complete Bohr model from the previous lesson
IV. Quantum Model of the Atom
A. Problem With the Bohr Model – Why
could the electron in hydrogen orbit in
only a small number of allowed paths?
B. Solving the Problem
1. Louis de Broglie – electrons have a
dual nature - they can act like
particles or waves !!!
Diffraction Patterns
x-rays through Al
electrons through Al
2. Schrodinger – developed equations
that treat electrons in atoms like waves
a. describe the shapes of the orbitals
in which electrons have a high
probability of being found
b. quantum theory – mathematical
explanation for the wave properties
of electrons that apply to all atoms
Louis de Broglie
(1892-1987)
Electrons have a dual
nature (particle + wave)
Erwin Schrodinger
(1887-1961)
Schrodinger equation
describes wave
properties of electrons
mathematically
I. Quantum model of the atom
A. Louis DeBroglie- (1892-1987) French physicist
1. Electrons have a wave/ particle nature –so if light is
passed through a slit – wave interference occurs- proved
by the equation (1924)
Wavelength = Planck’s constant/mass times velocity
If all moving objects have wave characteristics why don’t we
see ourselves waving?
Everybody – stadium wave
Matter is a Wave
 Does not apply to large objects
 Things bigger then an atom
 A baseball has a wavelength of about 1032 m when moving 30 m/s
 An electron at the same speed has a
wavelength of 10-3 cm
 Big enough to measure.
The physics of the very small
 Quantum mechanics explains how the very
small behaves.
 Classic physics is what you get when you
add up the effects of millions of packages.
 Quantum mechanics is based on probability
because
II.Heisenberg Uncertainty
Principle
 It is impossible to know exactly the position
and velocity of a particle at the same time.
 The better we know one, the less we know
the other.
 The act of measuring changes the
properties.
 Look at the fan
More obvious with the very
small
 To measure where a electron is, we use light.
 But the light moves the electron
 And hitting the electron changes the frequency of
the light.
 Watch the balloon
Before
Photon
Moving
Electron
After
Photon
changes
wavelength
Electron
Changes
velocity
C. Principles of the Quantum Model
1. electrons act like waves
2. probability of an electron being found
at various distances from the nucleus
3. orbitals – a 3-D region about the
nucleus where a specific electron may
be found
4. electrons have greater energy as their
distance from the nucleus increases
5. energies of orbitals are quantized within
main energy levels
6. the exact location of electrons can not
be pinpointed – they are found in regions
of high probability called orbitals or
electron clouds
Quantum Atomic Model
 Similarities between Bohr and Schrodinger
– 1. the closer the orbital to the nucleus the lower
the energy
– 2. to move from a lower to a higher level the
energy absorbed must be equal to the
difference between the levels
Quantum Model
 3.When e- drops from a a higher to lower
level electromagnetic radiation is emitted
=difference in energy levels
 4. the most probable location of the (H) e- is
a distance equal to Bohr’s lowest energy
level.
Atomic Orbitals
 Principal Quantum Number (n) = the
energy level of the electron.
 Within each energy level the complex
math of Schrodinger’s equation describes
several shapes.
 These are called atomic orbitals
 Regions where there is a high probability
of finding an electron. 90%
Orbitals (s, p, d, f)
Orbitals (s, p, d types)
s orbitals
(one type)
p orbitals
(3 types)
d orbitals
( 5 types)
Orbitals in Sodium (Na)
S orbitals
 1 s orbital for every
energy level
 Spherical
shaped
 Each s orbital can hold 2 electrons
 Called the 1s, 2s, 3s, etc.. orbitals.
P orbitals




Start at the second energy level
3 different directions
3 different shapes
Each can hold 2 electrons
P Orbitals
D orbitals
 Start at the third energy level
 5 different shapes
 Each can hold 2
electronshttp://www.falstad.com/qmatom/
F orbitals
 Start at the fourth energy level
 Have seven different shapes
 2 electrons per shape
F orbitals
Summary
# of
Max
shapes electrons
Starts at
energy level
s
1
2
1
p
3
6
2
d
5
10
3
f
7
14
4
V. Quantum Numbers
A. Principal Quantum Number
1. main energy level
B. Orbital Quantum Number
1. shape of orbital (s,p,d,f)
C. Magnetic Quantum Number
1. orientation of orbital about the nucleus
D. Spin Quantum Number
1.indicates clockwise or counterclockwise spin of the electron (+ or – ½)
Create a model of sub-atomic levels
By Energy Level




First Energy Level
only s orbital
only 2 electrons
1s2
 Second Energy Level
 s and p orbitals are
available
 2 in s, 6 in p
 2s22p6
 8 total electrons
By Energy Level
 Third energy level
 s, p, and d orbitals
 2 in s, 6 in p, and 10
in d
 3s23p63d10
 18 total electrons
 Fourth energy level
 s,p,d, and f orbitals
 2 in s, 6 in p, 10 in d,
ahd 14 in f
 4s24p64d104f14
 32 total electrons
By Energy Level
 Any more than the
fourth and not all the
orbitals will fill up.
 You simply run out of
electrons
 The orbitals do not fill
up in a neat order.
 The energy levels
overlap
 Lowest energy fill
first.
Increasing energy
7s
6s
5s
7p
6p
5p
4p
4s
3p
3s
2p
2s
1s
6d
5d
4d
3d
5f
4f
I.Electron Configurations
 A. The way electrons are arranged in
atoms.
 1..Aufbau principle- electrons enter the
lowest energy first.
 2.This causes difficulties because of the
overlap of orbitals of different energies.
 B.Pauli Exclusion Principle- at most 2
electrons per orbital - different spins
Electron Configuration
 C. Hund’s Rule- When electrons occupy
orbitals of equal energy they don’t pair up
until they have to .
 Let’s determine the electron configuration
for Phosporus
 Need to account for 15 electrons
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p  The first two electrons
go into the 1s orbital
2p
 Notice the opposite
spins
 only 13 more
5f
4f
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
5f
4f
3d
3p  The next electrons go
into the 2s orbital
2p
 only 11 more
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p • The next electrons go
into the 2p orbital
2p
• only 5 more
5f
4f
Increasing energy
7s
6s
5s
4s
3s
2s
1s
7p
6p
5p
4p
6d
5d
4d
3d
3p • The next electrons go
into the 3s orbital
2p
• only 3 more
5f
4f
Increasing energy
7s
6s
5s
4s
7p
6p
6d
5d
5p
4d
4p
3p •
3s
2s
1s
2p •
•
•
5f
4f
3d
The last three electrons
go into the 3p orbitals.
They each go into
separate shapes
3 unpaired electrons
1s22s22p63s23p3
The easy way to remember
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
1s
• 2 electrons
Fill from the bottom up
following the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
1s 2s
• 4 electrons
Fill from the bottom up
following the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
• 12 electrons
Fill from the bottom up
following the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
6
2
3p 4s
• 20 electrons
Fill from the bottom up
following the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
6
2
10
6
3p 4s 3d 4p
5s2
• 38 electrons
Fill from the bottom up
following the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
6
2
10
6
3p 4s 3d 4p
5s2 4d10 5p6 6s2
• 56 electrons
Fill from the bottom up
following the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
6
2
10
6
3p 4s 3d 4p
5s2 4d10 5p6 6s2
4f14 5d10 6p6 7s2
• 88 electrons
Fill from the bottom up
following the arrows
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p
1s
•
2
2
6
2
1s 2s 2p 3s
6
2
10
6
3p 4s 3d 4p
5s2 4d10 5p6 6s2
4f14 5d10 6p6 7s2
5f14 6d10 7p6
• 108 electrons
Write these electron
configurations
 Titanium - 22 electrons
 1s22s22p63s23p64s23d2
 Vanadium - 23 electrons
1s22s22p63s23p64s23d3
 Chromium - 24 electrons
 1s22s22p63s23p64s23d4 is expected
 But this is wrong!!
Exceptions to Electron
Configuration
Orbitals fill in order
 Lowest energy to higher energy.
 Adding electrons can change the energy of
the orbital.
 Half filled orbitals have a lower energy.
 Makes them more stable.
 Changes the filling order
Chromium is actually





1s22s22p63s23p64s13d5
Why?
This gives us two half filled orbitals.
Slightly lower in energy.
The same principal applies to copper.
Copper’s electron
configuration





Copper has 29 electrons so we expect
1s22s22p63s23p64s23d9
But the actual configuration is
1s22s22p63s23p64s13d10
This gives one filled orbital and one half
filled orbital.
 Remember these exceptions
Valence Electrons
 I. Valence electrons are defined as electrons
located in the highest occupied energy level of an
atom.
A. Inner electrons are not shown.
1. inner electrons are not part of the bonding relationship
between elements
B.Electron dot structure- show the symbol of an element with dots
to represent the electrons on the highest energy level
1. G.N. Lewis- American chemist (1875-1946) – devised the
method
The Lewis dot structure for Oxygen
O
Oxygen is in group VIA so it has 6 valence electrons
The Lewis dot structure for Chlorine
Cl
chlorine is in group VIIA so it has 7 valence electrons
The Lewis dot structure for calcium
Ca
calcium is in group IIA so it has 2 valence electrons
Lewis dot structure of a compound
NH3
1) How many valence electrons does N have?
N is in group VA so it has 5 valence electrons
2) How many valence electrons does H have?
H is in group IA so each H has one valence electron
3) How many valence electrons does Neon have.
Making calcium chloride
Ca + Cl
Cl
Ca( Cl )2
Lewis dot structure of a compound
NH3
H
N H
H
Lewis dot structure and making ammonium ion
NH4+
H
+
H
H N H
H
+
Orbitals (s, p, d, f)
Orbitals (s, p, d types)
s orbitals
(one type)
p orbitals
(3 types)
d orbitals
( 5 types)
Orbitals in Sodium (Na)
C. Principles of the Quantum Model
1. electrons act like waves
2. probability of an electron being found
at various distances from the nucleus
3. orbitals – a 3-D region about the
nucleus where a specific electron may
be found
4. electrons have greater energy as their
distance from the nucleus increases
5. energies of orbitals are quantized within
main energy levels
6. the exact location of electrons can not
be pinpointed – they are found in regions
of high probability called orbitals or
electron clouds
VI. Determining Number of Orbital
Types and Electrons
A. If n = the number of the principal energy
level or shell ( 1-7) and there is a maximum
of 2 electrons per orbital then:
1. n = the possible number of orbital types
for that shell
2. n2 = total number or orbitals possible
3. 2n2 = total number of electrons possible
4. Heisenberg Uncertainty
Principle
Both the velocity and
position of a particle
(electron) can not be
measured at the
same time
Werner Heisenberg
(1901-1976)
B. Examples
If n = 3 then in energy level 3:
 3 orbital types possible (s,p,d)
 9 orbitals are possible
 18 electrons are possible
(n)
(n2)
(2n2)
If n = 4 then in energy level 4:
 4 orbital types possible (s,p,d,f) (n)
 16 orbitals are possible
(n2)
 32 electrons are possible
(2n2)
VII. Electron Configuration
A. Rules and Principles
1. Aufbau Principle – an electron
occupies the lowest energy orbital that
can receive it
2. Hund’s Rule – orbitals of equal energy are
each occupied by one electron before
any orbital is occupied by a second
electron
3. Pauli Exclusion Principle – no two
electrons in the same atom can have the
same set of four quantum numbers
B. Types of Electron Configurations
1. Electron –configuration notation
a. indicates number of the principal
energy level, the orbitals, and
the number of electrons possible
2. Orbital Notation – arrows indicate location
and spin of electrons
3. Electron-dot structure – indicates valence
shell electrons