# 3 Unit 6 PP

```Electronic Structure
and the Periodic Table
Unit 6 Honors Chemistry
Electromagnetic Waves:
Electromagnetic waves:
progressive, repeating disturbances that
come from the movement of electric
charges
Electromagnetic Waves &amp; Light
Wavelength and Frequency
 Wavelength (, lambda):
distance between any
two points in a wave
 measured in any
distance unit
(mainly nm or m:
1 nm = 1x10-9 m)
Wavelength Can be Measured in One
of Two Ways…
Wavelength and Frequency
Frequency (; pronounced nu):
the number of cycles
of the wave that pass through a
point in a unit of time
Measured in sec-1 (/sec)
1
sec-1 = 1 Hertz (Hz)
Illustration of Frequency
Wavelength is indirectly
proportional to frequency
As Wavelength increases, frequency
_________________.
As Wavelength decreases, frequency
_________________.
Amplitude
Note: height of wave is
amplitude (intensity or
brightness of wave)
Amplitude is INDEPENDENT of
frequency or wavelength!
Speed
Speed (c): The speed of light!
c = 3.00 x 108 m/s
(rounded to 3 sig figs)
Equation
 One equation relates speed,
frequency and wavelength:
c=
Example
 The wavelength of the radiation which
produced the yellow color of sodium vapor
light is 589.0 nm. What is the frequency of
The electromagnetic spectrum
 complete range of wavelengths and frequencies
 mostly invisible
What is color?
TED Talk: What is color?
The visible/continuous spectrum
 continuous spectrum: components of
white light split into its colors, ROY G
BIV
 from 390 nm (violet) to 760 nm (red)
Line Spectra
 Pattern of lines produced by light
emitted by excited atoms of an element
 unique for every element
 used to identify unknown elements
How do we see color?
TED Talk: How we see color
Max Planck
 Light is generated as a stream of
particles called PHOTONS
 Equation:
E (Energy of a photon)= h
(h =Plank’s constant=
6.626x10-34Js)
Relationships in Planck’s Eqn.
E = h•
High frequency, low λ, high E.
Low frequency, high λ, low E.
Photoelectric effect – Nobel Prize in
Physics 1921 to Einstein
Occurs when light strikes the surface of
a metal and electrons are ejected.
Practical uses:
Automatic
door openers
Photoelectric Effect: Conclusion
Light not only has wave
properties but also has particle
properties. These massless
particles, called photons, are
packets of energy.
Example 6.2
Using the frequency calculated in the previous
example, calculate the energy, in joules, of a
photon emitted by an excited sodium atom.
Calculate the energy, in kilojoules, of a mole of
excited sodium atoms.
Bohr’s Hydrogen Atom:
A Planetary Model
Niels Bohr: Proposed
planetary model.
Electrons “orbit” the nucleus like planets
around the sun.
NOT current model of atom but used to
explain some features of atom.
Ground State vs. Excited State
 ground state: all electrons in lowest
possible energy levels
 excited state: an electron that has
absorbed energy and moved to a
higher energy level
 This is a temporary state!!
Explanation of Line Spectra &amp;
Equation
Niels Bohr
 Energy of an
electron is
quantized: can
only have specific
values.
 Energy
proportional to
energy level.
Explanation of Line Spectra
Electron will
drop from
excited state to
ground state and
will emit energy
as a photon.
Explanation of Line Spectra
 Type of photon emitted by electron depends on
energy difference of energy levels
Elevel = -RH
1
– 1
(nhi)2
(nlow)2
AND Elevel = h = hc/
(h: Planck’s constant, 6.626 x 10-34 J sec/photon)
Flaw in Bohr’s Model
Only works well for 1 electron
species (H atom).
Does not explain fine structure of
line spectra.
Wave-Particle Duality
Light has properties of both
WAVES and PARTICLES.
 most matter has undetectable wavelengths
(1000 kg car at 100 km/hr has  = 2.39 x 10-38 m)
 This work led to the development of the electron
microscope
Quantum Mechanics
 Quantum mechanics:
atomic structure based on wavelike properties of the electron
 Schr&ouml;dinger: wave equation that
describes hydrogen atom
Heisenberg Uncertainty
Principle
 The exact location of an electron cannot be
determined (if we try to observe it, we
interfere with the particle)
 You can know either the location or the
velocity but not both
 Electrons exist in electron clouds
and not on specific rings or orbits
Quantum Numbers
 Four quantum numbers are a mathematical
way to represent the most probable location
of an electron in an atom
 analogy...
state = energy level, n
city = sublevel, l
house number = spin, ms
Principal Quantum Number: n
 Always a positive integer (1,2, 3…7)
 Indicates size of orbital, or how far electron is
from nucleus
 Similar to Bohr’s energy levels or shells
 Larger n value = larger orbital or distance
from nucleus
The Periodic Table and
Shells
n = row number on periodic table for a given element
n=1
n=2
n=3
n=4
n=5
n=6
n=7
Angular Momentum Quantum Number: l
 positive integer from zero to n-1
 Sublevel within an energy level; indicates
shape of orbital
0 = s
1 = p
2 = d
3 = f
Types of Sublevels
s
p
d
Magnetic Quantum Numbers: ml
 integer from -l to +l
 Indicates orientation of orbital in space
 Orbital
= electron containing area
Spin Quantum Number: ms
 Two values only: + &frac12; or -&frac12;
2
electrons max. allowed in each orbital
 (Pauli Exclusion Principle)
 Indicates spin of electron; spins of each
electron must be opposite
REVIEW:
QUANTUM NUMBERS
Every Electron has four!
n ---&gt; level
1, 2, 3, 4, ...
l ---&gt; sublevel
0, 1, 2, ... n - 1
ml ---&gt; orbital
-l ... 0 ... +l
ms ---&gt; electron spin +&frac12; and -&frac12;
Orbitals
 No more than 2 e- assigned to an
orbital
 Orbitals grouped in s, p, d (and f)
subshells
s orbitals
d orbitals
p orbitals
Capacities of levels, sublevels, and
orbitals—see packet
Example
Example 6.6 Give the n and l values for the
following orbitals:
a. 3p
b. 4s
Example
Example 6.8 What are the possible ml values
for the following orbitals:
a. 3p
b. 4f
Shapes of
Atomic Orbitals
Shapes of Atomic Orbitals
s = spherical
p = peanut
d = dumbbell (clover)
f = flower
Multielectron Atoms
In the hydrogen
atom the subshells
(sublevels) of a
principal energy
level or shell are at
the same energy
level.
Previous Equation:
En = –RH /n2
Multielectron Atoms
In a
multielectron
atom, only the
orbitals are at the
same energy
level: the
sublevels are at
different energy
levels!
The increasing energy order of
sublevels is generally:
s&lt;p&lt;d&lt;f
Overlapping subshells
At higher
energy
levels,
sublevels
overlap.
Note:
4s vs. 3d!
Introduction to Electron Configuration
Definition: describes the distribution of
electrons among the various orbitals in
the atom
Represents the most
probable location of
the electron!
EOS
Electron Configurations
 The system of numbers and letters that
designates the location of the electrons
 3 major methods:



Full electron configurations
Abbreviated/Noble Gas configurations
Orbital diagram configurations
Full or Complete Electron Configuration
(uses spdf)
Uses numbers to
designate a principal
energy level and the
letters to identify a
sublevel; a superscript
number indicates the
number of electrons in a
designated sublevel.
EOS
Rules for Electron Configurations
The Aufbau principle:
Electrons fill from the lowest
energy level to the highest
(they don’t skip around)
1s22s22p63s23p64s23d10e
tc.
Pauli Exclusion Principle
No two electrons in the same
atom can have the same set of
4 quantum numbers.
That is, each electron has a unique
In other words, the maximum # of electrons an
orbital can hold is 2 e- (one with ms = +1/2 and one
with ms = -1/2)
HUND’S RULE
Orbitals of equal energy in a sublevel
must all have 1 electron before the
electrons start pairing up
a.k.a “creepy person on the bus rule”
*** also electrons in half-filled
orbitals have same spin
Why are these incorrect?
Why are these incorrect?
Why are these incorrect?
Full Electron Configuration
Example Notation:
 1s2 2s1 (Pronounced “one-s-two, two-s-one”)
A. What does the coefficient mean?
Principle energy level
B. What does the letter mean?
Type of orbital (sublevel)
C. What does the exponent mean?
# of electrons in that orbital
Steps to Writing Full Electron
Configurations
1. Determine the total number of electrons the atom
has (for neutral atoms it is equal to the atomic
number for the element).
Example: F
atomic # =
# of p+ = # of e- =
2. Fill orbitals in order of increasing energy (see
Aufbau Chart).
3. Make sure the total number of electrons in the
electron configuration equals the atomic number.
Aufbau Chart (Order of Energy Levels)
When writing electron
configurations:

d sublevels are n – 1
from the row they
appear in

f sublevels are n – 2
from the row they
appear in
Writing Electron
Configurations
Nitrogen:
Helium:
Phosphorous:
Rhodium:
Bromine:
Cerium:
Abbreviated/Noble Gas Configuration
i. Where are the noble gases on the periodic
table?
ii. Why are the noble gases special?
iii. How can we use noble gases to shorten
regular electron configurations?
Abbreviated/Noble Gas Configuration
Example: Barium
1. Look at the periodic table and find the noble
gas in the row above where the element is.
2. Start the configuration with the symbol for that
noble gas in brackets, followed by the rest of
the electron configuration.
Abbreviated/Noble Gas Configuration
Practice! Write Noble Gas Configurations for
the following elements:
Rubidium:
Bismuth:
Arsenic:
Zirconium:
Writing Electron
Configurations
Another way of
writing
configurations is
called an orbital
diagram.
(also called orbital
notation or orbital
diagrams)
ORBITAL BOX NOTATION
for He, atomic number = 2
2
1s
1s
Arrows
depict
electron
spin
One electron has n = 1, l = 0, ml = 0, ms = + &frac12;
Other electron has n = 1, l = 0, ml = 0, ms = - &frac12;
Orbital Diagrams
Orbital diagrams use boxes (sometimes circles)
to represent energy levels and orbitals. Arrows
are used to represent the electrons.
= orbital
sublevels
Orbital Diagrams
Don’t forget - orbitals have a capacity of two electrons!!
Two electrons in the same orbital must have opposite spin
so draw the arrows pointing in opposite directions.
Increasing Energy 
Example: oxygen
2p
2s
1s
1s22s22p4
Drawing Orbital Diagrams
1. First, determine the electron configuration for the element.
2. Next draw boxes for each of the orbitals present in the electron
configuration.
 Boxes should be drawn in order of increasing energy (see
the Aufbau chart).
3. Arrows are drawn in the boxes starting from the lowest energy
sublevel and working up. This is known as the Aufbau principle.
 Add electrons one at a time to each orbital in a sublevel
before pairing them up (Hund’s rule)
 The first arrow in an orbital should point up; the second
arrow should point down (Pauli exclusion principle)
4. Double check your work to make sure the number of arrows in
your diagram is equal to the total number of electrons in the atom.
 # of electrons = atomic number for an atom
Electron Configurations for
Nitrogen
Electron Configurations for
Nickel
Lithium
3p
3s
2p
2s
1s
Group 1A
Atomic number = 3
1s22s1 ---&gt; 3 total
electrons
Beryllium
3p
3s
2p
2s
1s
Group 2A
Atomic number = 4
1s22s2 ---&gt; 4 total
electrons
Boron
3p
3s
2p
2s
1s
Group 3A
Atomic number = 5
1s2 2s2 2p1 ---&gt;
5 total electrons
Carbon
3p
Group 4A
Atomic number = 6
1s2 2s2 2p2 ---&gt;
6 total electrons
3s
2p
2s
1s
Here we see for the first time
HUND’S RULE.
Nitrogen
3p
3s
2p
2s
1s
Group 5A
Atomic number = 7
1s2 2s2 2p3 ---&gt;
7 total electrons
Oxygen
3p
3s
2p
2s
1s
Group 6A
Atomic number = 8
1s2 2s2 2p4 ---&gt;
8 total electrons
Fluorine
3p
3s
2p
2s
1s
Group 7A
Atomic number = 9
1s2 2s2 2p5 ---&gt;
9 total
electrons
Neon
3p
3s
2p
2s
1s
Group 8A
Atomic number = 10
1s2 2s2 2p6 ---&gt;
10 total electrons
Note that we have
reached the end of
the 2nd period, and
the 2nd shell is full!
Exceptions to the Filling Order Rule
(Cr, Cu)—these will not be on test!
Valence electrons
Importance and definition:
Definition: Electrons in the outermost energy levels;
they determine the chemical properties of an
element.
Write the noble gas configuration...the valence
electrons are the ones beyond the core.
Example: Sulfur
Valence Electrons and Core
Configuration (Shorthand)
What is the shorthand notation for S?
Sulfur has six valence electrons
EOS
Configurations of Ions
Cations: Formed when metals lose e– in
highest principal energy level.
Example:
(Z = 11) Na
(Z = 11) Na+
EOS
Configurations of Ions
Anions: Formed when non-metals gain e–
to complete the p sublevel.
Example:
Z= 18
Cl
EOS
Transition Metals
Transition metals (and p block metals) lose e–
from the highest principal energy level (n)
FIRST, then lose their d electrons!
Zr: [Kr] 5s24d2
Zr+2 : [Kr] 4d2
EOS
Isoelectronic Species
Definition: Ions or atoms that have the same
number of electrons
Example: Neon, O2-, F-, Na+, Mg2+, Al3+
all have the same configuration (1s22s22p6) and
are isoelectronic
Electron Spin and
Magnetism
•Diamagnetic: NOT
attracted to a magnetic
field
•Paramagnetic:
substance is attracted to
a magnetic field.
•Substances with
unpaired electrons are
paramagnetic.
Examples
 Mg
 Cl
Write orbital notation: if it has an unpaired e- it is
paramagnetic.
Periodic Properties &amp; Trends
 Electronegativity

Ability of an atom to pull e- towards itself

Increases going up and to the right


Across a period  more protons in nucleus = more
positive charge to pull electrons closer
Down a group  more electrons to hold onto =
element can’t pull e- as closely
Periodic Properties &amp; Trends
Electronegativity



Ability of an atom to pull e- towards itself
Across a period  more protons in nucleus = more
positive charge to pull electrons closer
Down a group  more electrons to hold onto = protons
in nucleus can’t pull e- as closely
Definition:
&frac12; experimental distance between
centers of two bonded atoms
Trend in a family:
Size increases
down a group.
(More principal
energy levels)
Trend in a period:
Size decreases across a period, e- more
strongly attracted to nucleus.
Transition metals:
Size stays
relatively constant
across a period; eadded to inner
energy level.
Memory Device
LLLL: Lower Left, Larger Atoms
Sizes of Ions
+
Li,152 pm
3e and 3p
Li + , 78 pm
2e and 3 p
 CATIONS are SMALLER than the
atoms from which they are formed.
 Size decreases due to increasing he
electron/proton attraction.
Sizes of Ions
F, 71 pm
9e and 9p
F- , 133 pm
10 e and 9 p
 ANIONS are LARGER than the atoms
from which they are formed.
 Size increases due to more electrons
in shell.
Trends in Ion Sizes
Trends in ion sizes are the same
as atom sizes.
Active Figure 8.15
First Ionization Energy
Definition: energy required to remove an
electron from an atom in the gas phase.
Mg (g) + 738 kJ ---&gt; Mg+ (g) + e-
First Ionization Energies
Trend in a group:
Decreases going down a
group (e- further away;
easier to remove)
Trend in a period:
Increases going across a
period (e- held more
tightly).
EOS
Memory Device
LLLL: Lower Left,
Larger Atoms;
Looser electrons
Second Ionization Energy
Definition: energy required to remove 2nd
electron from an atom in the gas phase.
Takes
e- is
Mgmore
(g) + energy
738 kJ because
---&gt; Mg+ (g)
+ eremoved from increasingly positive ion.
Mg+ (g) + 1451 kJ ---&gt; Mg2+ (g) + e-
Electron Affinity
Some elements GAIN electrons to form
anions.
Electron affinity is the energy involved
when an atom gains an electron to form
an anion.
A(g) + e- ---&gt; A-(g) E.A. = ∆E
Trends in Electron Affinity
Trend in a group:
Affinity for edecreases going
down a group
Trend in a series
or period:
Affinity for eincreases going
across a period
Electron Affinity
Note that the
trend for E.A.
is the SAME as
for I.E.!
Trends in Metallic Properties
Most metallic means easiest loss of electrons!
Metals are on left, nonmetals on right of p.t.
A Summary of Periodic Trends
Remember LLLL!!
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