# Mill Creek High School

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```					   Electron Configurations

How does the Bohr Model of
an atom compare to the
Quantum Mechanical Model
of an atom?
What is the wave nature of light?
• Visible light is part of
form of wave energy that
travels through empty space
in the form of alternating
electric and magnetic fields.

• All waves consist of crest
and troughs traveling away
from a source at a velocity
determined by the nature of
the wave & the material
through which the wave
passes.
How are frequency & wavelength
related?
• Frequency – rate of wave vibration & is the # of
waves that pass a given point per second.
Expressed as Hz – 1 Hz = 1 wave/sec

• Wavelength – determined by frequency &
velocity, it is the distance between points on a
continuous wave.
If frequency ↑ wavelength ↓
If frequency ↓ wavelength ↑
Wave A                 Wave B

1.   What wave has the shortest wavelength?
2.   Which wave has the longest wavelength?
3.   Which wave has the lowest frequency?
4.   Which wave has the highest frequency?
5.   Which wave do you think has the most
energy?
Wave A                 Wave B

1.   What wave has the shortest wavelength? A
2.   Which wave has the longest wavelength? B
3.   Which wave has the lowest frequency? B
4.   Which wave has the highest frequency? A
5.   Which wave do you think has the most
energy? A
6. Which electromagnetic radiation would cause the most damage, gamma-rays
The more light bends the shorter
the wavelength.
R
O
Y
G
B
I
V
•    Which color of light bends the most?
•    Describe its wavelength and frequency.

•    Which color of light has the lowest frequency?
•    Which color has the longest wavelength?
The more light bends the shorter
the wavelength.
R
O
Y
G
B
I
V

•    Which color of light bends the most? Purple
•    Describe its wavelength and frequency. Short
Wavelength High Frequency

•    Which color of light has the lowest frequency? Red
•    Which color has the longest wavelength? Red
Niels Bohr
(1913)
Bohr
• placed p+ & no in the center or nucleus of
an atom
• Placed e- in rings around the nucleus
• worked in Rutherford‟s Lab                Electron

Electron

e- gets excited       Proton

Ground state
e- relax
Bohr
• A Quantum model of H based on its
spectrum.
• Predicted correctly frequencies of the lines
H atomic emission spectrum.
• Proposed – H atom has certain allowable
energy states
Lowest level – ground state
Gain energy – excited states
• H has many different excited states
Emission Spectrum of H(sun), H-2
He, Hg, U
Visible Spectrum
ROY G BIV

Hydrogen has 4 colors
Red
Aqua
Blue
Violet
(colored lines are
called “spectral lines”
Proposed
•   Ground state is the 1st energy level & atom doesn‟t radiate energy

•   Add energy from an outside source e- moves to higher energy level (orbit) to
excited states

•   An atom emits a photon (color of light) corresponding to the difference
between energy levels of two orbits.

Light (has certain
wavelength &
frequency = color)

electron that absorbed                      electron in ground state
energy and is “excited”                     that emitted light energy

ground state

nucleus                             nucleus
Draw the Bohr model of the atom
• The number of protons and neutrons are listed
in the center of the atom
• The number of electrons must be determined
• The electrons are placed in concentric rings
around the nucleus
–   1 The first ring can only hold 2 electrons
–   2 The next ring can only hold 8 electrons
–   3 The next ring can only hold 8 electrons
–   4 The next ring can only hold 18 electrons
–   5 The next ring can only hold 18 electrons
–   6 The next ring can only hold 32 electrons
–   7 The next ring can only hold 32 electrons
Bohr Model WS
• Fill in the information and place the correct
amount of e-‟s in each ring

• Fill in the orbitals from the inside (ground
State) out (excited states)
Problem
• Failed to explain spectrum of other
elements.
• Did not fully account for atoms behavior
• Research shows e- do not move in circular
orbits

•   Did lay groundwork for other models
What is the quantum mechanical
model for atoms?

• (1924 – Louis de Broglie 1892-1987)
• Proposed a model to account for fixed
energy levels of Bohr‟s Model
• Thought that light waves have both wave
& particle like behavior.
• If waves have particle like behavior could
particles have wavelike behavior?
Erwin Shrödinger
• Austrian physicist; using complex
equation, developed quantum mechanical
model (electron cloud) of the atom

– Electrons treated as waves
– Electrons not in circular orbits
– Mathematically predicts probable location of
an electron
Quantum Mechanical Model
• e- are treated as
waves
• limits e- to certain
values
• doesn‟t describe e-
path
• 3-dimensional
design
• atomic orbital –
describes e-
probable location
e- wiggle as the move around nucleus
Werner Heisenberg (1901- 1976)
Heisenberg Uncertainty Principle
• Impossible to make any measurement on
an object without disturbing the object,
even a little.
• States it is fundamentally impossible to
know precisely both the velocity & position
of a particle @ the same time.

Where is the e-?
Assign Principle Quantum #‟s (n)
(n) increases as orbitals become larger
• e- away from nucleus have increased
energy
• n = principle energy level ground state
(PEL)
•     Ex: n = 1 ground state
H has a n = 1-6
PEL contains
• energy sublevels s,p,d,f according to the
shapes of the atom‟s orbitals.

s - spherical

p - dumbbell shaped

d - f have different shapes
Energy Levels
• While Bohr‟s model of an atom is an easy
representation it is not correct. Electrons exist in
energy levels and must be assigned correctly to
determine the valence electrons

• There are 4 suborbitals each holding a different
number of electrons
s – holds 2 electrons
p – holds 6 electrons
d – holds 10 electrons
f – holds 14 electrons
To determine the electron configuration
for any atom simply look at the periodic
table

•   Columns = Groups = Family
•   Groups 1 - 2 are s orbitals (He)
•   Groups 13 – 18 are p orbitals
•   Groups 3 – 12 are d orbitals
•   Lanthanide and Actinide are f orbitals
Periodic Table
• Color your periodic table indicating the s,
p, d, f elements
• Pp. 161 in book
s = yellow
p = purple
d = blue
f = green
Fill in WS
• Energy levels are numbered from 1 to 7 and are
where electrons are located. Energy levels can
only hold a certain # of electrons. Electrons fill
inner energy levels before they fill outer energy
levels. Electrons in the outer most energy level
determine the behavior of an element. Another
name for these electrons would be valence
electrons. Atoms with full outer shells are stable
or inert. Generally, that means 8 electrons and is
called the octet rule. The exception to this rule is
He which has only 2 valence electrons.
Sublevels:

Sublevels are areas within an energy level.
They are designated by the letters s, p, d, f.

s holds 2 electrons             __
p holds 6 electrons          __ __ __
d holds 10 electrons      __ __ __ __ __
f holds 14 electrons __ __ __ __ __ __ __

Each ___ is a ring in the sublevel & holds 2 e-
Sublevels per Energy Level
•   n = 1 = 1 sublevel    1s
•   n = 2 = 2 sublevels   2s   2p
•   n = 3 = 3 sublevels   3s   3p   3d
•   n = 4 = 4 sublevels   4s   4p   4d     4f

• n = 5-7 are the same as the 4th energy
level
Electron Configuration

• A method of keeping track of how many
electrons are located at each energy level and
which sublevel they are in. The number
represents the energy level. The letter
represents the sublevel. The superscript
represents how many electrons are located at
a certain place.
• The period (row number on PT) is important in
determining the energy level

• To write an electron configuration you must write
all the electrons in the atom

Energy level→1s2   superscript (number of electrons)

Sublevel
Aufbau Diagram
•   1s
•   2s   2p
•   3s   3p   3d
•   4s   4p   4d   4f   Energy pathway for electrons

•   5s   5p   5d   5f
•   6s   6p   6d
•   7s   7p
Aufbau Diagram
•   1s                  Notice the pathway goes from
4s to 3d to 4p. This is b/c the
•   2s   2p             4s electrons have more energy
than the 3d, but less energy
•   3s   3p   3d        than the 4p electrons.

•   4s   4p   4d   4f   d suborbitals are in 1 energy
level less than the period
•   5s   5p   5d   5f   (row) the element is in.

•   6s   6p   6d        f suborbitals are in 2 energy
levels less than the period
•   7s   7p             (row) the element is in.
Electronic Configurations
• The arrangement of electrons in an atom is
called the atom‟s configuration.

• Electrons will arrange themselves in a way that
gives the atom the lowest possible energy.

Energy level→1s2   superscript (number of electrons)

Sublevel
3 rules define how electrons are
arranged…
• The Aufbau principle states that electrons will
occupy the lowest energy level available

• The Pauli exclusion principle Two e- within same orbit
must have opposite spins

• Hund‟s rule single e- with the same spin must occupy each
equal energy orbit before additional e- with opposite spins can
occupy the same orbital
Practice Electron Configurations
• Fill in arrows to represent electrons

• Follow Aufbau pathway when filling orbitals

• Transfer arrows to other WS and write electron
configuration notation for each atom
Identify each element
Identify each element

Fluorine     Phosphorus
Electron Configuration
Examples
•   H     1s1
•   He    1s2
•   Li    1s2 2s1
•   Be    1s2 2s2
•   B     1s2 2s2 2p1
•   C     1s2 2s2 2p2
•   N     1s2 2s2 2p3
•   O     1s2 2s2 2p4
•   F     1s2 2s2 2p5
•   Ne    1s2 2s2 2p6
Electron Configuration
Examples
•   Na    1s2 2s2 2p63s1
•   Mg    1s2 2s2 2p63s2
•   Al    1s2 2s2 2p63s23p1
•   Si    1s2 2s2 2p63s23p2
•   P     1s2 2s2 2p63s23p3
•   S     1s2 2s2 2p63s23p4
•   Cl    1s2 2s2 2p63s23p5
•   Ar    1s2 2s2 2p63s23p6
•   K     1s2 2s2 2p63s23p64s1
•   Ca    1s2 2s2 2p63s23p64s2
Electron Configuration
Examples
•   Sc    1s2 2s2 2p63s23p64s23d1
•   Ti    1s2 2s2 2p63s23p64s23d2
•   V     1s2 2s2 2p63s23p64s23d3
•   Cr    1s2 2s2 2p63s23p64s23d4
•   Mn    1s2 2s2 2p63s23p64s23d5
•   Fe    1s2 2s2 2p63s23p64s23d6
•   Co    1s2 2s2 2p63s23p64s23d7
•   Ni    1s2 2s2 2p63s23p64s23d8
•   Cu    1s2 2s2 2p63s23p64s23d9
•   Zn    1s2 2s2 2p63s23p64s23d10
Electron Configuration
Examples
• Cs        1s2 2s2 2p63s23p64s23d104p6
5s24d105p66s1

• Hg        1s2 2s2 2p63s23p64s23d104p6
5s24d105p66s24f145d10
Noble Gas (Short Hand) Notation
of Electron Configurations
1. Find the noble gas in the row above the
element you need to write the notation
for and put in brackets [ __ ].

Ex: Looking for S (sulfur) the noble
gas would be Neon

S[Ne]
Noble Gas (Short Hand) Notation
of Electron Configurations
2. Write the notation starting in the period
(row) the element is in.

S[Ne]3s23p4

Write noble gas notation using these
elements:
Br, Ag, U, Xe
What are Valence Electrons?
• e-„s that determine chemical properties of
an element

• These e- are in the outermost orbitals
usually the highest principle energy level
How are valence e- determined
• Find which group
(column) element is
in and that is how    1   2   13 14 15 16 17 18
many valence
electrons the         1   2   3   4   5   6   7   8
element has.

• Works for groups
1-2 (He has only 2
ve-) 13-18
How are valence e- determined
from electron configurations

S[Ne] 3s2 3p4 (6 e- in outermost level)

Cs[Xe] 6s1 (1 e- in outermost level)
What are Electron Dot or Lewis
Structures?
(1875-1946)
• Li[He] 1s1       Li    1 ve-

• Ne[He] 2s2 2p6   Ne    8 ve-
What are Electron Dot or Lewis
Structures?
(1875-1946)
• Li[He] 1s1       Li    1 ve-

• Ne[He] 2s2 2p6   Ne    8 ve-

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