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8
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Maximum stability results
when an atom is isoelectric with a noble gas
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Lewis Dot Symbol: an elemental symbol surrounded by dots where each dot represents a
valence electron
*For the main group elements, the number of dots is the same as the group number (1A-8A)
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*The exact order in which the dots are placed around the element symbol is not important, but
the number of dots is
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*For main group metals, the number of dots in the Lewis dot symbol is the number of electrons
that are lost when the atom forms a cation that is isoelectric with the preceding noble gas
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*In addition to atoms, we can also represent atomic ions with Lewis dotskeletal structure
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Draw the symbols
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Count the valence
dot symbol of the atom and include the ion’s charge
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3
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2: Ionic Bonding
each bond
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Distribute the remaining
form anions
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Ionic Bonding: an electrostatic attraction that holds oppositely charged ions together in an ionic
5
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ions in the gas phase
6
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*The magnitude of lattice energy is a measure of an ionic compound’s stability
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Lattice Energies of Selected Ionic Compounds
CompoundLattice Energy (kJ/mol)Melting Point (°C)
LiF
1017
845
LiCl
860
610
LiBr
787
550

LiCl
LiBr
LiI
NaCl
NaBr
NaI
KCl
KBr
KI
MgCl2
Na2O
MgO

860
787
732
787
736
686
699
689
632
2527
2570
3890

610
550
450
801
750
662
772
735
680
714
Sub*
2800

*

Na2O sublimes at 1275°C
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• Smallest distance between ions = strongest attractive forces = largest lattice
energy
• Largest distance between ions = weakest attractive forces = smallest lattice
energy
*Lattice energy increases as the distance between ions decreases
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3: Covalent Bonding
Lewis Theory of Bonding: a chemical bond involves atoms sharing
electrons
Covalent Bonding: two atoms sharing a pair of electrons
Covalent Bond: a shared pair of electrons
Octet Rule: atoms will lose, gain, or share electrons to achieve a noble gas
electron configuration
Lone Pair: a pair of valence electrons that are not involved in covalent bond
formation
Lewis Structure: a representation of covalent bonding in which shared electron pairs are shown
either as dashes or as pairs of dots, and lone pairs are shown as pairs of dots on individual atoms
• Only valence electrons are shown in a Lewis structure
Single Bond: a pair of electrons shared by two atoms
Multiple Bond: a chemical bond in which two atoms share two or
more pairs of electrons

Multiple Bond: a chemical bond in which two atoms share two or
more pairs of electrons
Double Bond: a multiple bond in which the atoms share two pairs
of electrons
Triple Bond: a multiple bond in which the atoms share three pairs of electrons

*Multiple bonds are shorter than single bonds
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• Bond enthalpy can be used to quantify the intramolecular bonding force
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• Intermolecular forces are usually quite weak compared to the forces holding
atoms together within a molecule, so molecules of a covalent compound
are not held together tightly
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*The electrostatic forces holding ions together in an ionic compound are usually very strong, so
ionic compounds are solids at room temperature and have high melting points
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2
2
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5
Molar heat of vaporization* (kJ/mol)600
30
3
Density (g/cm )
2
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59
Solubility in water
High
Very low
Electrical conductivity
 Solid
Poor
Poor
 Liquid
Good
Poor
 Aqueous
Good
Poor
*

The molar heat of fusion and molar heat of vaporization are the amounts of heat needed to melt 1 mole of
the solid and to vaporize 1 mole of the liquid, respectively
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4: Electronegativity and Polarity
*Ionic bonds occur between a metal and a nonmetal
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Polar: having a nonuniform electron density
Polar Covalent Bond: bonds in which electrons are unequally shared

*The δ symbol is used to denote partial charges on the atoms
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*In general, electronegativity increases from left to right across a period in the periodic table, as
the metallic character of the elements decreases
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*Transition metals do not follow the above trends
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*Atoms of elements with widely
different electronegativities tend
to form ionic compounds with
each other because the atom of the
less electronegative element gives
up its electrons to the atom of the more
electronegative element
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*Only atoms of the same element
which have the same electronegativity,
can be joined by a pure covalent bond
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5 is generally considered purely covalent or nonpolar
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5 to 2
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*A bond between atoms whose electronegativities differ by 2
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Dipole Moment (μ): a quantitative measurement of the polarity of a bond
• μ is always positive
• Q = charge
• r = distance between the charges/bond length (in meters)
*Bond lengths are usually expressed in angstroms (Å) or picometers (pm) à so it’s generally
necessary to convert to meters
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*Dipole moments are usually expressed in debye units (D)
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0622 10-19C

Bond Lengths and Dipole Moments of the Hydrogen Halides
MoleculeBond Length (Å)Dipole Moment (D)
HF
0
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82
HCl
1
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08
HBr
1
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82
HI
1
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44
Percent Ionic Character: the ration of experimentally-measured dipole moment to calculated
dipole moment – multiplied by 100%

8
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From the molecular formula, draw the skeletal structure of the compound, using
chemical symbols and placing bonded atoms next to one another
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In general, the least electronegative atom will be the central atom
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2
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a
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a
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3
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4
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If there is
more than one type of terminal atom, complete the octets of the most
electronegative atoms first
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If any electrons remain after step 4, place them in pairs on the central atom
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If the central atom has fewer than eight electrons after completing steps 1-5, move
one or more pairs from the terminal atoms to form multiple bonds between the
central atom and the terminal atoms
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) Like Lewis dot symbols for atomic anions, Lewis structures for
polyatomic anions are enclosed by square brackets
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6: Lewis Structure and Formal Charge
Formal Charge: method of electron “bookkeeping” in which shared electrons are divided
equally between atoms that share them
• Can be used to determine the most plausible Lewis structure when more
than one possibility exists for a compound
*Formal charge is determined by comparing the number of electrons associated with an atom in a
Lewis structure with the number of electrons that would be associated with the isolated atom
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*To determine the number of electrons associated with an atom in a Lewis structure…
1
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2
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*The sum of formal charges must equal the
overall charge on the species
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*To determine the best skeletal arrangement use formal charges and…
1
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2
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The best skeletal arrangement of atoms will give rise to Lewis structures in
which the formal charges are consistent with electronegativities
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For example, the more electronegative atoms should have the more
negative formal charges
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7: Resonance
Resonance Structures: two or more equally valid Lewis structures for a single
molecule that differ only in the positions of electrons (not in the positions of their
atoms)
*The double-headed arrow indicates a resonance structure
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8: Exceptions to the Octet Rule
*The octet rule almost always holds for second-period elements
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1
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a
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Properties do not differ from those of a normal covalent bond
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The central atom has fewer than eight electrons due to an odd number of electrons
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Free Radicals (Radicals): a molecule with an odd number of electrons
3
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8
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4
H─N 393
H─O 460
H─S 368
H─P 326
H─F
568
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9
H─Br
366
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3
C─H 414
C─C 347
C═C 620
C≡C 812
C─N 276
C═N 615

Bond Bond Enthalpy
(kJ/mol)
C≡O
1070
C─P 263
C─S 255
C═S 477
C─F 453
C─Cl 339
C─Br 276
C─I 216
N─N 193
N═N 418
N≡N
941
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7
O─P 502
O═S 469
O─F 190
O─Cl 203
O─Br 234
O─I 234
P─P 197
P═P 489
S─S 268
S═S 352
F─F
156
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7
Cl─F
193

C─N
C═N
C≡N
C─O
C═O†

276
615
891
351
745

N─F 272
N─Cl 200
N─Br243
N─I 159

Cl─Cl
Cl─F
Br─Br
I─I

242
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5
151
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The C═O bond enthalpy in CO2 is 799 kJ/mol
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