Chapter 4
Electron Structure of the Atom
Review of Atomic Structure
• The center of the atom is called the nucleus.
• In it are the particles with mass: the
protons and neutrons
• Protons determine the identity of an atom
• Electrons determine the properties of an
atom
• Where are the electrons?
Light
•
•
What is light, and what does it have to do
with the electrons in an atom?
Light is electromagnetic radiation.
Electromagnetic Radiation (EMR)
• Form of energy with wavelike behavior as it
travels through space at the speed of light
• Seven major types:
• Gamma () Rays
• X rays
• Ultraviolet (UV)
• Visible Light (ROYGBIV)
• Infrared (IR)
• Microwaves
• Radio waves
Electromagnetic Spectrum
(p. 98)
5
Wavelength (λ)
• The distance between two corresponding
points on a wave
• Units are same as length - m, or commonly
nm (10-9 m)
Figure 7.5
Know the wavelengths for the
visible spectrum
700
650
600
550
500
450
400 nm
•
Frequency ( )
• number of wave cycles that move through a
point in space in 1 s
• hertz (Hz)
• same as inverse seconds (1/s) or (s-1)
( )
• Which has the greater wavelength the red light or the
green light?
• Which has a greater frequency?
Figure 7.7
Frequency & Wavelength
• inversely proportional
• i.e. as one increases the other
decreases
c = λν
c = speed of light (3.00 x 108 m/s)
λ = wavelength (in meters)
ν = frequency (in Hz)
Practice
i.
What is the frequency and wave
type of a wave with a wavelength of
5.2 x 10-7 m (5.2 x 102 nm)?
wave type: green visible light
Practice
i.
What is the frequency and wave
type of a wave with a wavelength of
5.2 x 10-7 m (5.2 x 102 nm)?
wave type: green visible light
c  

c

3.00 10 m / s
14
14
v

5.8

10
(1/
s
)

5.8

10
Hz
7
5.2 10 m
8
•
What is the approximate
frequency of blue light?
R O Y G B IV
Wavelength = 475nm
400
700
•
Convert nm to m
R O Y G B IV
400
700
475nm 1109m

 4.75 107 m
1nm
ii. What is the approximate
frequency of blue light?
3.00 108 m
s
7
4.75 10 m
  6.32 1014 Hz

400

700

c
R O Y G B IV
Homework
1. Calculate the frequency of light with a
wavelength of 3.41 x 103 cm?
2. Calculate the wavelength of light with a
frequency of 3.21 x 1016 Hz? What type
of EMR is it?
3. What is the frequency of orange light?
Line Spectra
•
Continuous spectrum
•
•
•
all the wavelength of in the visible spectrum
Produced by white light
Line Spectrum
•
•
•
•
distinct colored lines
each a single wavelength of light
Visible when an element has been heated
“atomic fingerprint”
Energy is Quantized!
Quantized = quantity = specific measured amount
•
•
•
•
Max Planck
energy produced by atoms can only have certain values
only distinct lines are seen in element line spectra
can only exist at certain wavelengths.
Max Planck
• Proposed that objects emit energy in small
packets called “quanta”
• Quantum: min. quantity of energy lost or
gained by an atom
• This quantized energy is related to the
frequency of the energy
Photons – waves and particles
Energy of a photon
− directly proportional to the frequency
Ephoton = hν
−
inversely proportional to the wavelength
Ephoton = hc/
Ephoton = (in Joules)
h = Planck’s constant (6.626 x 10-34 Js)
ν = frequency (in Hz)
  wavelength in meters
Low 
High 
Low E

High 
Low 
High E

Practice
i.
What is the energy of a photon with a
frequency of 1.000x1017Hz?
E  h
1
E  (6.626 10 Js)(1.000 10 )
s
E  6.626 1017 J
34
17
Practice
ii.
What is the energy of a photon
with a wavelength of 8.1m?
8.1 m110 m
 8.1106 m
.......1 m
6
E
hc

34
(6.626 10 Js )(3.00 10 m / s)
E
6
8.110 m
E  2.5 1020 J
8
Atomic Spectra
• When visible light passes through a prism,
its components separate into a spectrum.
• White light, such as sun light or light from
a regular light bulb, gives a continuous
spectrum:
Photoelectric Effect
• Emission of electrons from a metal
when light shines on it
• This was explained by Einstein
• Based on Planck's work, Einstein proposed
that light also delivers its energy in
chunks; light would then consist of little
particles, or quanta, called photons, each
with an energy of Planck's constant times
its frequency
• Electrons are only emitted if the photons
have a high enough energy (high enough
frequency)
Atomic Spectra
• Colored light gives only specific colors in a
line spectrum:
27
What does this have to do with
electrons?
• Hydrogen atom: 1 proton, 1 electron
• Passing electricity through a tube
containing hydrogen gas gives off a pink
light
• Light is given off as electrons fall from an
excited state to a ground state
• The light can be separated into four
frequencies
Spectra of other elements
(spectral signatures applet)
Energy Level Transition
• Excited State: higher energy state
Ephoton
• Ground State: lowest energy state
Ephoton= Eexcited- Eground=h
Bohr Model links electron and emissions
•
•
•
•
Niels Bohr (Danish physicist)
• planetary model
electron are in orbits
orbit 1 closest to the nucleus
increasing numbers as the
orbits get further away from
the nucleus.
•
•
•
•
•
Bohr Model
Orbits have a fixed
radius.
Electrons cannot exist
between orbits
lowest energy is closest
to the nucleus
increases as the orbits
get further away
Electrons absorb or emit
energy when they change
orbitals
∆E = Ef – Ei
Deficiencies of Bohr Model
• Did not work for other atoms
• Doesn’t explain chemical behavior of atoms
• More complex model needed
• However, the Bohr model did give insight
into the quantized behavior of the atom
that was better understood in later days
Modern Models of the Atom
• Bohr
• deBroglie
• Schrödinger (Wave Mechanical Model)
• To view at home, click here. Click on Run
Now!
Modern Model of the Atom
(Quantum Mechanical Model)
• electrons exist in orbitals.
• Orbitals are 3-dimensional
regions in space where an
electron is likely to be found
• not a circular pathway
• The electron is thought of
having wave properties
35
• The exact location of the
electron cannot be known.
(Heisenberg Uncertainty
Principle)
• The location is described as a
probability
• This is a Probability Map for
lowest-energy state of the
electron in an H atom
36
A typical representation of an
s-orbital
Principal Energy Levels
• Orbitals of similar size exist in the same
principal energy level (n=1, 2, 3…)
• The principal energy levels correspond to
Bohr’s energy levels and represent a
distance from the nucleus
38
Energy levels contain orbitals
• Lower energy orbitals are smaller.
• Higher energy orbitals are larger;
further away from the nucleus.
• An orbital can hold at most 2 electrons.
39
Sublevels
•
Orbitals are arranged in sublevels, which have
specific shapes
•
s, p, d, and f are sublevels
s
p
d
f
s Orbitals
Figure 7.14
p Orbitals
d Orbitals
f Orbitals
44
Memorize terms!
• Principal Energy level: distance from
nucleus
• Sublevel: series of orbitals having equal
energy
• Orbital: particular region where an
electron exists
• Orbitals make up sublevels, and sublevels
make up energy levels
What each energy level holds
(Table 2, p. 110)
• The first energy level (n=1)
• A single s orbital
• The second energy level (n=2)
• 2s orbital and 2p sublevel (three 2p orbitals)
• The third energy level (n=3)
• 3s orbital, 3p sublevel (three 3p orbitals) and
3d sublevel (five 3d orbitals)
• The fourth energy level (n=4)
• 4s orbital, 4p sublevel, 4d sublevel and 4f
sublevel (seven 4f orbitals)
Hydrogen Orbital Diagram
• Aufbau Orbital diagrams
• Show the sublevels and orbitals that can exist
at each principal energy level
• Each box represents an orbital
• Groups of boxes represent sublevels
• In the hydrogen atom only, the sublevels within a principal
energy level all have the same energy.
Multielectron Orbital Diagram
• In the multielectron atoms, the sublevels within a principal
energy level have different energy levels.
Orbital Diagram Rules
•
Aufbau principle
•
•
Pauli exclusion principle
•
•
A maximum of two electrons can occupy each orbital, and
they must have opposite spins.
Hund’s rule
•
•
Electrons fill orbitals starting with the lowest-energy
orbitals.
Electrons are distributed into orbitals of identical energy
(same sublevel) in such a way as to give the maximum
number of unpaired electrons.
Electrons are always filled in their ground state, or
lowest energy state.
Hydrogen’s Orbital Diagram
1 electron
Helium’s Orbital Diagram
Lithium’s Orbital Diagram
Boron’s Orbital Diagram
Carbon’s Orbital Diagram
Pg. 251 (Carbon’s orbital diagram)
Filling Orbital Diagrams
• Why do the electrons in the p sublevel
occupy separate orbitals?
• It takes a little bit of
energy to pair up
electrons, so single
electrons occupy
different orbitals with
the same energy
(Hund’s Rule)
Orbital Diagrams for the 1st Ten Elements
Electron Configurations
An electron configuration is a shorthand notation for
representing the number of electrons in each sublevel
Carbon has 6 electrons.
Therefore, using the orbital diagram we obtain:
1s22s22p2
60
Order of Fill
Periodicity of Electron Configurations
• Can you tell the patterns among the following groups of
elements?
• Alkali Metals (Group IA)
Li
Na
1s22s1
1s22s22p63s1
Mg
Ca
1s22s22p63s2
1s22s22p63s23p64s2
Cl
Br
1s22s22p63s23p5
1s22s22p63s23p64s23d104p5
Ne
Ar
1s22s22p6
1s22s22p63s23p5
• Alkali Earth Metals (Group IIA)
• Halogens (Group VIIA)
• Noble Gases (Group VIIIA)
Using the Periodic Table for Electron Configurations
• Blocks contain elements with the same highest-energy sublevel.
Periodicity of Electron Configurations
• Notice the number of columns in the s, p,
d and f blocks is the same as the number
of electrons allowed in each sublevel
Periodicity of Electron Configurations
• The principal energy level number, the
number that comes before the sublevel
letter designation, is the same as the
period number for the s and p sublevels.
• For the d sublevels, the principal energy
level number is one less than the periodic
number.
The Principal Quantum Number and
Sublevel on the Periodic Table
Figure 7.21
66
An Example of Electron Configuration
P: 1s22s22p63s23p3
Mn: 1s22s22p63s23p64s23d5
Abbreviated Electron Configuration
Notice that phosphorus’ electron configuration starts
out with neon’s electron configuration
Phosphorus 1s22s22p63s23p3
Neon
1s22s22p6
We use the symbol for Neon to represent neon’s
electron configuration in the configuration for
phosphorus.
The abbreviated electron configuration for P would be:
[Ne] 3s23p3
• The outermost electrons are often shown
in this way:
• [Ar] 4s23d10 4p4
• The most important electrons are the
ones with the highest number, or energy
level
Valence Electrons
• The electrons on the outside edge of the atom
• This is where the action is- where bonding takes place
• Atoms have no more than 8 valence electrons
Neon
Argon
Radon
Energy
Level
(Shell)
Maximum
Number of
Electrons
1
2
3
4
5
6
7
2
8
18
32
50
72
98
Max
number of
Valence
Electrons
2
8
8
8
8
8
8
The Octet Rule:
• Atoms will combine to form compounds in
order to reach eight electrons in their
outer energy level.
• Atoms with less than 4 electrons tend to
lose electrons.
• Atoms with more than 4 electrons tend to
gain electrons.
Electron-dot diagrams can be used to give
the number of valence electrons
• The number of valence electrons is equal to
the element’ group number
•
1A 2A
3A 4A 5A 6A 7A 8A
Write the electron-dot symbols for the
following elements:
iodine
phosphorus
gallium
argon
Valence Electrons, review
• The highest principal energy level is called
the valence level, or valence shell
• The valence level contains electrons that
are furthest from the nucleus (on the
outside edge) and highest in energy
• Inner-level electrons are called core
electrons
• Valence electrons are important in chemical
reactions