Unit 3 – Bonding Review
Octet Rule!
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8 is the magic number! Most atoms want a complete valence shell with 8
electrons
Exceptions: For elements like hydrogen, helium, lithium etc, they can be
happy with 2 electrons which completely fill the first principle energy level
Atoms can gain or lose electrons OR share electrons in order to get this
complete Octet
Electronegativity is a measure of the tendency of an atom to attract electrons
towards itself. The higher the electronegativity, the more that atom attracts
electrons!
Electronegativity increases from left to right on the Periodic Table
Electronegativity decreases from top to bottom on the Periodic Table
**Fluorine is the most electronegative element**
Why is this important to know?!
1. Electrons with higher electronegativity tend to gain electrons to obtain
stable octet/complete valence shell
2. Electrons with lower electronegativity tend to lose electrons to obtain
stable octet/complete valence shell
3. Non-metals have higher electronegativities than metals; explains ionic
bonding
Ionization energy is the energy to remove one valence electron!
 Metals have low ionization energies
 Non-metal have high ionization energies
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Ion Formation and Properties
Metals
Non-Metals
Cation
Anion
Low electronegativity
 High electronegativity
Lose electrons
 Gain electrons
Form positive (+) ions
 Form negative (-) ions
Ion size is smaller than atom size
 Ion size is larger than atom size
Ion Formation:
1. Group 1 – Always forms +1 ions
2. Group 2 – Always forms +2 ions
3. Group 17 – Always forms -1 ions
4. Oxygen/Sulfur – Forms -2 ion; Nitrogen forms -3 ion
5. Non-Metals – Use the top charge/oxidation state on your periodic table
6. Transition Metals – Can form ions with multiple charges. Can only determine
charge in a compound
**When an atom gains or loses electrons, the resulting electron configuration
is typically the same configuration as a neutral noble gas, the stable Group 18
elements**
Dot Structures – Atoms and Ions
1. Draw the symbol for the atom
2. Count how many valence electrons the atom has
3. Begin adding electrons (represented by dots) to the top, right, bottom and
left side of the symbol
4. After you have added four dots or electrons, begin to pair them
Examples:
Energetics of Bonding!
1. Breaking a bond always takes energy – endothermic
2. Making a bond always released energy – exothermic
3. Systems tends to move towards more stable, lower energy states
4. More on this when we get to kinetics and thermodynamics!
Ionic Bonding
 Ionic bonding involves a transfer of electrons
 Electronegativity difference > 1.7
 Properties:
o High melting point/boiling point
o Hard
o Brittle
o Electrolytes (ions dissociate and can conduct electricity when
dissolved in water or another polar solvent)
Writing Ionic Formulae
Ionic compounds are held together by electrostatic forces, the force of attraction
between the + and – charged ions (opposite charges attract!). The overall resulting
compound is neutral. In order to write the formula for a compound, you must
balance the charges!
**Use the Kriss Kross method!!**
Remember: Ionic formulas are ratios of the anions and cations in the salt crystal. If
you can reduce, you must!
Polyatomic Ions
Polyatomic ions are compounds where the individual atoms are covalently bonding
but with an overall charge. They can therefore enter into ionic compounds with ions
of the opposite charge. Polyatomics ions are found in your reference table!
Dot Structures – Ionic Bonding
Covalent Bonding
A covalent bond is formed when two non-metals share electrons in order to form a
stable octet
Lewis Dot Structures
 Used to represent bonding among covalent compounds
 Show which atoms are bound together and how many bonds are present
Steps To Draw:
1. Determine # and types of atoms in compound
2. Count total number of valence electrons. This is the number you must
show in your picture!
3. Arrange skeleton and use single bonds (a line) to connect all atoms
a. Carbon, if present, is usually in the center
b. Hydrogen and the halides (Group 17 can never be in the center)
c. Otherwise, least electronegative atom in center
4. Start to add lone pair electrons.
5. If you cannot make each atom happy, i.e. give it a stable octet, you need to
start adding double and triple bonds
6. Count you electrons! They should add to the original # of electrons you
started with!!
Polar vs. Non-Polar Bonds
If the electronegativity difference between two atoms bound together is >0.4 and
<1.7, the electrons will be shared unevenly between them and the bond will be
polar.
 The more electronegative atom gets a partially negative sign as the electrons
spend more time around its nucleus
 The less electronegative atom gets a partially positive sign as the electrons
spend less time around its nucleus
Polar vs. Non-Polar Compounds
The polarity of a compound is determined by its structure. If the compound is
symmetrical - i.e. it has either a linear or tetrahedral shape AND the same atoms all
bound to the central atom - the overall compound EVEN though it might have polar
bonds can be non-polar. This is because the forces that are pulling the electrons
cancel out. An example of this would be CO2 or CCl4. In the later, the C-Cl BOND is
polar but the overall compound CCl4 is non-polar because it is symmetrical and the
forces pulling on the electrons cancel out.
Molecular Geometry
VSEPR theory – electrons clouds want to be as far apart from each other as possible.
Effects the overall shape of molecules.
# Atoms Bonded
To Central Atom
2
2
2
3
3
4
# Of Lone Pairs
Bonded To
Central Atom
0
1
2
0
1
0
Shape
Bond Angles
Linear
Bent
Bent
Trigonal Planar
Trigonal Pyramidal
Tetrahedral
180
120
109.5
120
109.5
109.5
Nomenclature – Ionic and Covalent Compounds
Ionic
1. First, state the name of the cation
(including polyatomic cations
such as ammonium)
2. ONLY IF the cation is a transition
metal, use a Roman Numeral to
write the positive charge on the
anion (e.g. FeO = Iron (II) Oxide)
3. Finally, state the name of the
anion with its ending changed to
“-ide” if it is a non-metal element.
If it is a polyatomic ion, keep its
full name and original ending (ide, -ate or –ite)
Covalent
1. Write the prefix and name of the
least electronegative non-metal.
Exception: if the least
electronegative metal has a
subscript 1, do not use the
prefix mono- for the first
element named.
2. Write the prefix (including monoif the subscript is only 1) of the
second, more electronegative
non-metal and change the ending
to –ide
H20 = dihydrogen monoxide
Metallic Bonds
 Between one or more metals
 Metals have few valence electrons and low electronegativities and ionization
energies
 Metal atom is considered to have a central portion on “kernel” made of
nucleus and on valence electrons
 Positively (+) charged kernals are held together in a “sea of electrons” –
valence electrons freely float among these kernels in the metallic crystal
 Properties:
o Ductile
o Malleable
o Hard
o High melting points
o High boiling points
o Conduct heat and electricity well (even in solid state)
Coordinate Covalent Compounds
A coordinate covalent compound forms when one atom contributes both electrons
in the bond to the other atom.
Example: NH3 + H+ ------------ NH4+1
Network Solids
Network solids are covalent compounds where a large number of atoms are all
bonding together. These extremely large covalent compounds are extremely hard.
The classic two examples are diamond (a large number of carbon atoms all bound
together) and –(SiO2)- (quartz).
Allotropes
When an element can exist in more than one form, the multiple forms are called
allotropes.
Example: Carbon can be amorphous carbon, graphite or diamond; oxygen
can be the kind we breathe (O2) or ozone (O3)
Properties of Ionic versus Covalent Bonds!
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Ionic solids
hard
brittle
high melting point
form electrolyte solutions in water (solutions that conduct electricity)
form crystals with regular geometric shapes that cleave along cleavage
planes - salt
Covalent substances
Solids
Network solids - a large macromolecule held together by covalent bonds
 Hard
 Brittle
 High melting point
 Don't conduct electricity solid or melted
 Form crystals with regular geometric shapes that cleave along cleavage
planes - diamonds (C), sand (SiO2)
Molecular solids - discrete molecules held together by intermolecular forces
 May be soft
 May have a low melting point
 Don't conduct electricity solid or melted
Liquids
 Polar compounds dissolve best in other polar compounds (acids and water)
 Nonpolar compounds are often immiscible in water and dissolve better in
nonpolar compounds (tar in gasoline)
Intermolecular Forces
Intermolecular forces are the forces BETWEEN molecules that hold them together.
These forces are much weaker than the strength of covalent and ionic bonds. They
are all due to the attraction of partial positive charges and partial negative charges
in a molecule because attracted to the opposite partial sign in another molecule.
Three types of intermolecular forces can operate between covalent molecules:
1. Dispersion Forces - also known as London Forces (named after Fritz
London who first described these forces theoretically 1930) or as Weak
Intermolecular Forces or as van der Waal's Forces (named after the person
who contributed to our understanding of non-ideal gas behavior).
2. Dipole-dipole interactions – between two polar molecules
3. Hydrogen bonds – compounds that have F, O or N bonded to H
A fourth intermolecular force you should know is ion-dipole (also known as
molecule-ion)! This is when an ionic compound is dissolved into a polar substance,
such as water, and the ionic crystal dissolved.
4. Ion-Dipolar (Molecule-Dipole) – when an ionic compound dissolves in a
polar covalent substance. + ions attracted to partially negative side of
molecule and – ions attracted to partially positive side of molecule. Example:
NaCl in water.
Relative strength of Intermolecular Forces:
dispersion forces < dipole-dipole interactions < hydrogen bonds < ion-dipole
Dispersion Forces (London Forces, Weak Intermolecular Forces, van der Waal's
Forces)
 VERY WEAK forces of attraction between molecules resulting from a
momentary dipoles occurring due to uneven electron distributions in
neighboring molecules as they approach one another (imagine the electron
cloud temporarily behind concentrated mostly on one side of the molecule
and therefore temporarily induced a partial negative and partial positive
charge)
Dispersion forces are the only type of intermolecular force operating between
non-polar molecules!!!!!!!!!
As the mass of the molecules increases, so does the strength of the dispersion force
acting between the molecules, so more energy is required to weaken the attraction
between the molecules resulting in higher boiling points
Dipole-dipole Interactions
 are stronger intermolecular forces than Dispersion forces
 occur between molecules that have permanent net dipoles (polar molecules),
for example, dipole-dipole interactions occur between SCl2 molecules, PCl3
molecules and CH3Cl molecules.
The partial positive charge on one molecule is electrostatically attracted to the
partial negative charge on a neighboring molecule.
Hydrogen bonds
 occur between molecules that have a permanent net dipole resulting from
hydrogen being covalently bonded to either fluorine, oxygen or nitrogen.
 Strongest type of intermolecular force between molecules (ion-dipole is
stronger but contains ions)
 Important in biology!!! Holds your DNA together! Allows you to
breathe!
 Important in biology!!! Explains high boiling point of water!
H-bonds are stronger intermolecular force than either Dispersion forces or dipoledipole interactions since the hydrogen nucleus is extremely small and positively
charged and fluorine, oxygen and nitrogen being very electronegative so that the
electron on the hydrogen atom is strongly attracted to the fluorine, oxygen or
nitrogen atom, leaving a highly localized positive charge on the hydrogen atom and
highly negative localized charge on the fluorine, oxygen or nitrogen atom. This
means the electrostatic attraction between these molecules will be greater than for
the polar molecules that do not have hydrogen covalently bonded to fluorine,
oxygen or nitrogen.
Effect of Intermolecular forces on melting and boiling points of molecular
covalent substances:
 Explains extremely high boiling and melting points of H2O, NH3, etc.