Periodic Table and Bonding Notes
Electromagnetic radiation – energy that travels through space as waves.
Waves have three primary characteristics:
1. _______________ (- lambda) – distance between two consecutive peaks
or troughs in a wave. Unit = meter
2. _______________ (  = nu) – indicates how many waves pass a given
point per second. Unit = Hertz (Hz)
3. ________________ – 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.
_______________________ – Light as waves – Light as photons (de Broglie)
Photon/quantum – packet of energy – a “___________” of electromagnetic radiation
Energy - (E – change in energy) – Unit Joules (J)
Planck’s Constant – (h = 6.626 x 10-34 J * s)
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?
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 ____________ wavelength,
__________ frequency and _____________ energy. Red light has a ___________
wavelength, __________ frequency, and ___________ energy.
__________________ – atom with excess energy
__________________ – lowest possible energy state
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Wavelengths of light carry different amounts of energy per photon
Only certain types of photons are produced (see only certain colors)
_______________ – only certain energy levels (and therefore colors) are
allowed
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.
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
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, then d, then f
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
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Electron Configuration – arrangement of the electrons among the various orbitals of t
y
the atom
Ex:
Shapes of orbitals
- All s orbitals are spherical as the principle energy level increases the
diameter increases.
-
All p orbitals are dumbbell shaped – all have the same size and shape
within an energy level
-
All d orbitals are flower (clover) shaped and a donut – all have the same
size and shape within an energy level
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
Ex:
Noble Gas Configuration – Shorthand configuration that substitutes a noble gas for
electrons
Ex:
Valence Electrons – Electrons in the outermost (highest) principle energy level in an
atom
Core Electrons – innermost electrons – not involved in bonding
Valence Configuration – shows just the valence electrons
Ex
Periodic Table
Dimitri Mendeleev-1869- developed the first version of the periodic table.
He expressed the regularities as a periodic function of the _______________.
Henry Moseley- revised Mendeleev periodic table by describing regularities in
physical and chemical properties as periodic functions of the ____________.
Groups (family) – vertical column
Elements with similar ________________________ 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 ____________________________ of
valence electrons
Periodic Trends
1. Atomic Size (radius)
Increases – down a group
Decreases – across a period
Size of ions
Cation
Ca+2/Ca
Ca larger because Ca+2 lost 2 electrons
Anion
S-2/S
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• 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
Summary of Trends
Chemical Bonding Notes
Bond- force that holds groups of two or more atoms together and makes them
function as a unit
bond energy- energy required to __________ the bond (tells the bond strength)
Ionic bonding- between ionic compounds which contain a ____________ and a
____________________
 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 ______________ by nuclei
Polar Covalent bonding- ________________ sharing of electrons
 positive end attracted to the negative end
 (delta) indicates partial charge
_______________________ - 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
Zero
Intermediate
Large
Covalent character
Ionic character
Decreases
Decreases
Decreases
Increases
Increases
increases
Bond type
Covalent
Polar covalent
Ionic
Ex. List the following in order of increasing polarity.
H-H, O-H, Cl-H, S-H, F-H
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 ________________ and
chemical symbols
- indicates the composition of a compound
- indicates the number of atoms in one molecule
Molecule – covalently 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
Metals
Location: ____________ side of Periodic Table
Properties: Ductile – drawn into wires
Malleable – hammered into sheets
Metallic Luster – ________________
Good Conductors of Heat and Electricity
Nonmetals
Location: ___________ 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
1. Prefixes: (MEMORIZE)
Mono-1
tetra-4
hepta-7
di-2
penta-5
octa-8
tri-3
hexa-6
non-9
deca-10
2. prefixes are used with both the first named and second named
element. Exception: mono- is not used on the first word
3. second word ends in –ide
4. 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
Writing molecular formulas
Translate prefixes
Examples:
N2 O
Si8O5
NH3
P3I10
dihydrogen monoxide
tetrasulfur hexachloride
carbon monoxide
carbon dioxide
Valence electrons are used in bonding.
 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:
For a compound:
For a molecule:
Duet rule- only two electrons in the full shell
Octet rule- surrounded by eight electrons
Bonding pair- electrons shared with other atom
Lone pair or unshared pair- not involved in bonding
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 nonbonding electrons or lone pairs.
6. Choose a central atom and assemble the pieces to make all atoms involved
stable.
Ex. GeBr4
Single bond- involves two atoms sharing one pair
Double bond- involves tow atoms sharing two pairs
Triple bond- involves two atoms sharing three pairs
Ex. CH4
C2H4
C2H2
Resonance- more than one Lewis structure can be drawn for the molecule
Ex. CO2
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
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
PH3
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
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
Structure
Molecular (geometric) structure- three-dimensional 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.
1.
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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
2.



Bent
Four total pairs
Two shared pairs and two unshared pairs
Bond angle = 104.5
Ex. H2O
3. 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
4.
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Tetrahedral
Four total pairs
Four shared pairs no unshared pairs
Bond angle = 109.5
Ex. CCl4
5. Trigonal pyramid
 Four total pairs
 Three shared pairs and one unshared pair
 Bond angle = 107
Ex. NH3