Modern Atomic Theory Notes
Electromagnetic radiation – energy that
travels through space as waves.
Waves have three primary
characteristics:
•
Wavelength (- lambda) –
distance between two
consecutive peaks or troughs
in a wave. Unit = meter
•
Frequency (  = nu) –
indicates how many waves
pass a given point per
second. Unit = Hertz (Hz)
•
Speed – velocity (c = speed
of light = 3 x 108 m/sec) indicates how fast a given
peak moves in a unit of time
c = 
Electromagnetic radiation (light) is divided
into various classes according to
wavelength.
Wave- Particle Theory
– Light as waves –
Light as photons
(de Broglie)
Photon/quantum
– packet of
energy – a
“particle” of
electromagnetic
radiation
Energy - (E – change in
energy) – Unit Joules (J)
Planck’s Constant –
(h = 6.626 x 10-34 J * s)
Ephoton = h
Change in Energy of a photon =
(Planck’s Constant) x (frequency)
c =  + Ephoton = h
= Ephoton = hc

Ex: What is the wavelength of light with a frequency of 6.5 x 1014 Hz? What
is the change in Energy of the photon?
E = hc ΔE = 4.3 x 10-19 J
= c = 3 x 108 m/sec
Given


= (6.626 x 10-34 J x s)(3 x 108 m/s)
6.5 x 1014 Hz
 = 6.5 x 1014 Hz
-7 m
4.6
x
10
-7
λ = 4.6 x 10 m
 = ? ΔE = ?
Wrap – up
So with light waves, you can convert between wavelength,
frequency, and energy with two equations:
= c
E=h
And two constants:
c = 3 * 108 m/s
h = 6.626 * 10-34 J s
In the visible part of the spectrum, different colors
correspond to different frequencies, wavelengths and
energies. Blue light has a short wavelength, high
frequency and high energy. Red light has a long
wavelength, low frequency, and low energy.
Learning Check
The wavelength of red light was measured
as 695 nm.
a. Calculate the frequency.
b. Calculate the change in energy.
Excited State – atom with
excess energy
Ground State – lowest
possible energy state
Wavelengths of light carry
different amounts of
energy per photon
Only certain types of
photons are produced
(see only certain colors)
Quantized – only certain
energy levels (and
therefore colors) are
allowed
Emission and Absorption Spectra
Intensity
Color
Emission Spectrum – bright lines on a dark background. Produced as
excited electrons return to a ground state – as in flame tests.
Absorption Spectrum – dark lines in a continuous spectrum. Produced
as electrons absorb energy to move into an excited state, only
certain allowable transitions can be made. Energy absorbed
corresponds to the increase in potential energy needed to move the
electron into allowed higher energy levels. The frequencies
absorbed by each substance are unique.
An Element’s Fingerprint
• When excited by heat or electricity, gases glow with
characteristic colors.
• A prism can be used to spread out the light from these
hot gases.
• This reveals a series of discrete lines, the element’s
fingerprint.
• Chemists use these fingerprints (called spectral lines) to
identify elements both in the lab and in space.
Here are some spectral lines
Learning Check
Now, try matching each of the spectra from column A with its corresponding line plot from column B.
A
B
Lab – How Do We Know
Bohr Model – suggested that electrons move around the
nucleus in circular orbits
Only Correct for Hydrogen
Wave Mechanical Model – Described by orbitals gives no
information about when the electron occupies a certain
point in space or how it moves *aka – Heisenberg's
Uncertainty Principle
Parts of the Wave Mechanical Model
1. Principle Energy Level (n) – energy level
designated by numbers 1-7.
-called principle quantum numbers
1
2
3
4
5
6
7
2. Sublevel – exist within each principle energy level
-the energy within an energy level is slightly different
-each electron in a given sublevel has the same
energy
-lowest sublevel = s, then p,
s
p
d
then d, then f
f
Parts of the Wave Mechanical Model
cont.
3. Orbital – region within a sublevel or energy
level where electrons can be found
s sublevel – 1 orbital
p sublevel – 3 orbitals
d sublevel – 5 orbitals
f sublevel – 7 orbitals
- ** No more than two electrons can occupy an
orbital**
-an orbital can be empty, half-filled, filled
Electron Configuration – arrangement of the electrons
among the various orbitals of the atom
Ex: 1s22s22p6 = Neon
Sulfur = 1s2 2s2 2p63s2 3p4
Cd = 1s2 2s2 2p63s2 3p64s2 3d10 4p65s2 4d10
Na = 1s2 2s2 2p63s1
Ne
Na
Learning Check
Write electron configurations:
Oxygen
Chromium
Shapes of orbitals
All s orbitals are spherical
as the principle energy level
increases the diameter increases.
All p orbitals are dumbbell or figure-8 shaped – all have the
same size and shape within an energy level
4 of the d orbitals are 4-leaf clover shaped and the
last is a figure-8 with a donut – all have the same
size and shape within an energy level
f orbitals are complicated!!!!!
Electron Spin
Spin – motion that resembles earth rotating
on its axis– clockwise or counterclockwise
Pauli Exclusion Principle – two electrons in the same orbital must have
opposite spins
Hund’s Rule – All orbitals within a sublevel must contain at least one
electron before any orbital can have two
Orbital Diagram – describes the placement of electrons in orbitals
• use arrows to represent electrons with spin
• line represents orbital (s=1, p=3, d=5, f=7)
____ full
____ half-full
____ empty
Orbital Diagrams
Ex: Neon = 1s__2s__ 2p__ __ __
Carbon = 1s__2s__ 2p__ __ __
Zinc = 1s__2s__ 2p__ __ __3s__ 3p__ __ __
4s__ 3d__ __ __ __ __
Gallium =1s__2s__ 2p__ __ __3s__ 3p__ __ __
4s__ 3d__ __ __ __ __ 4p__ __ __
Learning Check
Draw orbital diagrams:
Oxygen
Chromium
• Noble Gas Configuration – Shorthand
configuration that substitutes a noble gas for
electrons
Na = 1s22s22p63s1 or
[Ne]3s1
Ex:
Sn = 1s22s22p63s23p64s23d104p65s24d105p2
or
[Kr]5s24d105p2
• Valence Electrons – Electrons in the outermost
(highest) principle energy level in an atom
Na = 1s22s22p63s1
1 Valence
Sn = 1s22s22p63s23p64s23d104p65s24d105p2
4 Valence
• Core Electrons – innermost electrons – not
involved in bonding
• Valence Configuration – shows just the valence
electrons
3rd Shell/1valence electron
Ex: Na = 3s1
Sn = 5s25p2
5th Shell/4 valence electrons
Learning Check
Write the noble gas configuration, valence
configuration, and number of valence electrons:
Oxygen
Chromium
Periodic Table
Dimitri Mendeleev-1869- developed the first
version of the periodic table.
He expressed the regularities as a periodic
function of the atomic mass.
Henry Moseley- revised Mendeleev periodic
table by describing regularities in physical
and chemical properties as periodic
functions of the atomic number
Groups (family) – vertical
column
Elements with similar
valence electrons
configurations
Group 1 – alkali metals – reactive
Group 2 - alkaline earth metals –
reactive
Group 3-12 – transition metals
Group 15 – nitrogen family
Group 16 – oxygen family –
reactive
Group 17 – halogens – very
reactive
Group 18 – noble gases
Periods – horizontal rows
•Period number
corresponds to the
principal quantum
number of valence
electrons
Periodic Trends
1. Atomic Radius/Size – size of an atom
Increases – down a group
Decreases – across a period
Size of ions
Cation
Anion
Ca+2/Ca
S-2/S
Ca larger because Ca+2 lost 2 electrons
S-2 larger because S-2 gained 2 electrons
2. Ionization Energy – energy required to
remove an electron from an individual
atom in a gas phase M(g)  M+(g) + e(energy to make a positive ion)
• Metals lose electrons to non-metals so
relatively low energy is needed
• High ionization energy means an electron
is hard to remove
Decreases – down a group
Increases - across a period
3. Electron Affinity – Electron affinity is the energy
involved when an electron is added to a
gaseous atom.
• Negative values of energy mean that energy
was released during the process. Atoms with
negative values of electron affinity have a very
strong attraction for electrons.
• Positive values of electron affinity have very little
attraction for electrons.
(energy involved in negative ions)
Decreases – down a group
Increases - across a period
4. Electronegativity is the tendency of an atom to
draw electrons to itself when in a covalent bond.
Consequently, the trends are the same as for
electron affinity.
The atoms with the highest electronegativity are
fluorine, then oxygen, then nitrogen. It is also
important to know that the electronegativity of
hydrogen is slightly less than that of carbon.
Decreases – down a group
Increases - across a period
5. Metallic Character
Increases – down a group
Decreases – across a period
Electronegativity
Electronegativity
Learning Check
Put the following elements in order of increasing
atomic radius:
a. Ge, Se, Fe, Ca
b. C, Pb, Sn, Si, Ge
Put the following elements in order of increasing
electronegativity:
a. Ge, Se, Fe, Ca
b. C, Pb, Sn, Si, Ge
Notes- Chemical Bonding
Bond- force that holds groups of two or more atoms
together and makes them function as a unit
bond energy- energy required to break the bond (tells the
bond strength)
Ionic bonding- between ionic compounds which contain a
metal and a nonmetal
• Atoms that lose electrons relatively easily react with an
atom that has a high affinity for electrons
• Transfer of electrons
Covalent bonding- between two nonmetals
• Electrons are shared by nuclei
Polar Covalent bonding- unequal sharing of electrons
• positive end attracted to the negative end
• (delta) indicates partial charge
• electronegativity-(p. 362) relative ability of an atom in a molecule to
attract shared electrons to itself
• The higher the atom’s electronegativity value, the closer the shared
electrons tend to be to that atom when it forms a bonds
Increases – across a period
Decreases- down a group
Electronegativity
difference
Bond type
Covalent
character
Ionic character
Zero (0-.4)
Covalent
Decreases
Increases
Intermediate
(.4 – 1.4)
Polar covalent
Decreases
Increases
Large (<1.4)
Ionic
Decreases
Increases
H-H = 2.1 - 2.1 = 0
Ex. List the following in order of increasing polarity. O-H = 3.5 - 2.1 = 1.4
Cl-H = 3.0 - 2.1 = .9
H-H, O-H, Cl-H, S-H, F-H
H-H, S-H, Cl-H, O-H, F-H
S-H = 2.5 - 2.1 = .4
F-H = 4.0 - 2.1 = 1.9
• Dipole moment- has a
center of positive
charge and a center
of negative charge
• Represented by an
arrow
• Arrow points toward
the negative charge
Chemical Formula – type of notation made with
numbers and chemical symbols
– indicates the composition of a compound
– indicates the number of atoms in one molecule
Molecule - Bonded collection of two or more atoms
of the same element or different elements
- monatomic molecule – one atom molecules
- diatomic molecule – two atom molecules
(seven) MEMORIZE
Br, I, N, Cl, H, O, F
Nonmetals
METALS
Metals
Semi-metals
Location: Left side of Periodic Table
Properties:
Ductile – drawn into wires
Malleable – hammered into sheets
Metallic Luster – shine
Good Conductors of Heat and Electricity
Nonmetals
Location: Right side of Periodic Table
Properties:
Brittle
Lack Luster – not shiny
Poor Conductors of Heat and Electricity
Semi-metals
Location: Along Stair-step
Properties:
Have properties of metals and nonmetals
also called METALLOIDS Si, Ge, As, Sb, Te, Po, At
Molecular Nomenclature
Molecular Compounds (molecules) – compounds made from two nonmetals
- electrons are shared by two atoms
Naming Molecular
Prefixes: (MEMORIZE)
Mono-1
tetra-4
hepta-7
deca-10
di-2
penta-5
octa-8
tri-3
hexa-6
non-9
prefixes are used with both the first named and second named element. Exception:
mono- is not used on the first word
second word ends in –ide
If a two syllable prefix ends in a vowel, the vowel is dropped before the prefix is attached
to a word beginning with a vowel
monooxide
Writing molecular formulas
Translate prefixes
Examples:
N2O = Dinitrogen monoxide
dihydrogen monoxide = H2O
Si8O5 = Octasilicon pentoxide
tetrasulfur hexachloride = S4Cl6
NH3 = Nitrogen trihydride
carbon monoxide = CO
P3I10 = Triphosphorus deciodide
carbon dioxide = CO2
Learning Check
Write the name:
a. C2O4
b. P2O5
Write the formula:
a. Dihydrogen monoxide
b. Phosphorus trihydride
1
Valence electrons are used in
2
bonding. 3 4 5 6 7
8
2
• Stable elements want to achieve 8 electrons similar to the noble gases
• If it’s a metal it wants to achieve the configuration for the noble gas
before.
• If it’s a nonmetal it wants to achieve the configuration for the noble gas
after.
Lewis Structure- representation of a molecule
Shows how the valence electrons are arranged
among the atoms in the molecule.
For an element:
s
pz X px
py
Oxygen
1s22s22p4
•• •
•O• •
For a compound:
Li
+ Cl
For a molecule:
•• ••
•• F F ••
•• ••
 [Li]+1 + [ Cl ]-1
Duet rule- only two electrons in the full shell
H & He
Octet rule- surrounded by eight electrons
Happy Eight!!!!!
Bonding pair- electrons shared with other
atom
Line (-) = 2 electrons
Lone pair or unshared pair- not involved in
bonding dots (••) = 2 electrons/each dot is
one electron
5 Steps for Covalently Bonded Lewis Structures
1. Find the total number of valence electrons.
2. Calculate the number of “needed” electrons to give each atom 8
electrons, except for H which wants 2.
3. Subtract valence electrons from the “needed” electrons. This is the
number of bonding electrons.
4. Divide the bonding electrons by 2, to find the number of bonds.
5. Subtract the bonding electrons from the valence electrons to find
the non-bonding electrons or lone pairs.
6. Choose a central atom and assemble the pieces to make all atoms
involved stable.
Ex. GeBr4
Valence = 1(4) + 4(7) = 32
Needed = 1(8) + 4(8) = 40
Bonding = 40 – 32 = 8
Bonds = 8/2 = 4 lines
Lone e- = 32 – 8 = 24 dots
Central atom = Ge
••
•• Br
••
• •Br• •
• •
Ge
•• Br ••
••
•• •
Br •
••
• Single bond- involves two atoms sharing one pair
• Double bond- involves two atoms sharing two pairs
• Triple bond- involves two atoms sharing three pairs
Ex. CH4
C2H4
C2H2
1. 1(4) + 4(1) = 8
2. 1(8) + 4(2) = 16
3. 16 - 8 = 8
4. 8/2 = 4 lines
5. 8 – 8 = 0 dots
Central atom = C
H
H
C
H
1. 2(4) + 4(1) = 12
2. 2(8) + 4(2) = 24
3. 24 - 12 = 12
4. 12/2 = 6 lines
5. 12 – 12 = 0 dots
Central atom = C
H
H
C C
H
H
1. 2(4) + 2(1) = 10
2. 2(8) + 2(2) = 20
3. 20 - 10 = 10
4. 10/2 = 5 lines
5. 10 – 10 = 0 dots
Central atom = C
H C C H
H
Resonance- more than one Lewis structure
can be drawn for the molecule
Ex. CO2
•
•
1. 1(4) + 2(6) = 16
2. 1(8) + 2(8) = 24
3. 24 - 16 = 8
4. 8/2 = 4 lines
5. 16 – 8 = 8 dots
Central atom = C
•O
••
••
••O
••
••O
C
O••
•
C
O••
C
•••
O•
••
Exceptions to the Octet Rule
1. boron and beryllium- tend to be electron
deficient
– boron can hold 6 total electrons
– beryllium can hold 4 total electrons
•
•
ex. BF3
BeH2
• •
1. 1(3) + 3(7) = 24
2. 1(6) + 3(8) = 30
3. 30 - 24 = 6
4. 6/2 = 3 lines
5. 24 – 6 = 18 dots
Central atom = B
• F•
B
•
•
• F • • F•
•• •
• ••
1. 1(2) + 2(1) = 4
2. 1(4) + 2(2) = 8
3. 8 - 4 = 4
4. 4/2 = 2 lines
5. 4 – 4 = 0 dots
Central atom = Be
H Be H
2. Electrons are small spinning electric
charges that create magnetic fields
– Diamagnetic- substances which have paired
electrons that cancel out the magnetic field
– Paramagnetic- substances the have one or
more unpaired electrons that show great
attraction to the magnetic field
Ex. O2
••O O••
••
••
PH3
••
H P H
H
3. Odd number of electrons
–
You cannot write electron dot structures that fulfill the octet
rule, when the total number of valence electrons is odd
Ex. NO
1. 1(5) + 1(6) = 11
No Drawing
4. Expanded Octet- expand the valence shell to include
more than 8 electrons
–
–
Phosphorus and sulfur can expand to include 10 or 12
electrons
You will know you have an expanded octet when you don’t
have enough bonds for the atoms present
Ex. SF6
••
F• ••
•
••
•• F • • F ••
S ••
••
••
••••
F •F• F •
•• •••• •••
Lab – Lewis Structures
Structure (shape)
Molecular (geometric) structure- threedimensional arrangement of the atoms in a
molecule
VSEPR model- valence shell electron pair
repulsion
• Lone pairs of electrons like to be as far
away from each other as possible
• Double and triple bonds “act” like a single
shared pair for shape.
Linear
Linear- two pairs of
electrons are present
around an atom
– One total pair – one shared
pair
– Two total pairs – two
shared pairs
– Bond angle = 180
Ex. BeCl2
1. 1(2) + 2(7) = 16
2. 1(4) + 2(8) = 20
3. 20 - 16 = 4
4. 4/2 = 2 lines
5. 16 – 4 = 12 dots
Central atom = Be
•
•
•
•
• Cl Be Cl •
•••
•
••
Bent
Bent
– Four total pairs
– Two shared pairs and
two unshared pairs
– Bond angle = 104.5
Ex. H2O
1. 2(1) + 1(6) = 8
2. 2(2) + 1(8) = 12
3. 12 - 8 = 4
4. 4/2 = 2 lines
5. 8 – 4 = 4 dots
Central atom = O
•
•
•• O H
H
Trigonal planar
Trigonal planar- whenever
three pairs of electrons are
present they should be placed
at the corners of a triangle
– Three total pairs
– Three shared pairs
– Bond angle = 120
Ex. BCl3
1. 1(3) + 3(7) = 24
2. 1(6) + 3(8) = 30
3. 30 - 24 = 6
4. 6/2 = 3 lines
5. 24 – 6 = 18 dots
Central atom = B
•
•
••Cl••
B
•
•
•Cl• •Cl•
•• •
• ••
Tetrahedral
Tetrahedral
– Four total pairs
– Four shared pairs no
unshared pairs
– Bond angle = 109.5
Ex. CCl4
Valence = 1(4) + 4(7) = 32
Needed = 1(8) + 4(8) = 40
Bonding = 40 – 32 = 8
Bonds = 8/2 = 4 lines
Lone e- = 32 – 8 = 24 dots
Central atom = C
••
•• Cl
••
• •Cl• •
• •
C
•• Cl ••
••
•• •
Cl •
••
Trigonal pyramid
Trigonal pyramid
– Four total pairs
– Three shared pairs
and one unshared pair
– Bond angle = 107
Ex. NH3
1. 1(5) + 3(1) = 8
2. 1(8) + 3(2) = 14
3. 14 - 8 = 6
4. 6/2 = 3 lines
5. 8 – 6 = 2 dots
Central atom = N
••
H N H
H