Covalent Bonding
What is a covalent bond?
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A chemical bond that results from the ________________ of electrons, to form a stable ______
or __________ (Hydrogen only needs 2 to be stable)
Molecule = ______ or more atoms that are held together by _______________ bonds
Majority of covalent bonds form between _______________ (CLOSE together on periodic table)
Diatomic molecule
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molecule containing the ______ _____ atoms
Some elements _________ exist this way because they are _____ stable than the individual atoms
The Seven Diatomic Elements :___________________________________
Bonds in _____ the ________________ ions and diatomics are all covalent bonds
Single Covalent Bonds
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Two atoms held together by a sharing of _____ pair of
electrons are joined together by a single covalent bond.
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An electron dot structure represents the __________
pair of electrons of the covalent bond by ______ dots.
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A ______________ formula represents the covalent bonds by _________ and shows the arrangement
of covalently bonded atoms.
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A pair of valence electrons that is ____ ___________ between atoms is called an unshared pair, also
known as a ______ pair of a _________________ pair.
Double & Triple Covalent Bonds
o
Atoms form _______ or ____________ covalent bonds if they can attain a noble gas structure by
sharing two or three pairs of electrons.
o
A double bond involves _________ ________ pairs of electrons.
o
A __________ bond involves sharing ___________ pairs of electrons
Bond Length
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From a study of various Nitrogen containing compounds ______ _____________ as a function of bond
type can be summarized as follows:
N-N 147pm
N=N 124pm
N≡N 110pm
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As a general rule, the _________ e that are shared:
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the __________ the covalent bond (N≡N > O=O > F–F)
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the _____________ the covalent bond (N≡N < O=O < F–F)
Single
Double
Triple
Covalent Nomenclature
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1.
Prefix System (binary molecules)
Add prefixes to indicate # of atoms. Omit mono- prefix on first element.
1. Change the ending of the
second element to -ide.
2. Second element ALWAYS gets a prefix.
Prefixes:
1- _______________________
2- _______________________
3- _______________________
4- _______________________
5- _______________________
6- _______________________
7- _______________________
8- _______________________
9- _______________________
10__________________
Examples:
•
CCl4 _______________________
•
arsenic trichloride _______________
•
N2O _______________________
•
dinitrogen pentoxide _____________
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SF6 _______________________
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tetraphosphorus decoxide __________
Lewis Diagrams
1. Arrange atoms
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___________ atom is usually in the center (often Carbon)
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If not Carbon, _______________ electronegative atom is in center
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Hydrogen is always _____________ (on the side, not a central atom)
2. Find ___________ # of e- available to bond (_____________ e-! only )
3. Place a __________ of electrons between _____________ atom and each terminal atom
4. Place remaining electrons in pairs around _____________ atoms (except H) to satisfy octet rule
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Any ______________ pairs are assigned to central atom
5. Determine whether or not central atom satisfies _____________
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If not, __________ one or more lone pairs from terminal atoms to double or triple bonds
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Certain atoms can be ___________ to octet rule – H, Be, B, S, P, Xe
Ex 1 : CF4
Ex 2: CO2
Polyatomic Ions:

To find ___________ # of valence e-:
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_________ 1e- for each negative charge
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____________ 1e- for each positive charge
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Place ______________ around the ion and label the charge
Ex 3: ClO4-1
Octet Rule Exceptions:
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______________  2 valence e-
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Boron & Beryllium get ____ & ____ valence e- respectively
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_____________ octet  more than 8 valence e- (e.g. S, P, Xe)
Ex 4: BeCl2
Ex 5: SF6
Resonance Structures
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Molecules that ____ be correctly represented by a _________ Lewis diagram
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Actual structure is an _____________ of all the possibilities
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Show all possible structures separated by double-headed ___________
Ex: SO3
Bond Polarity
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Most bonds are a blend of ________ & __________ characteristics.
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____________ in electronegativity determines bond type.
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Electronegativity
o
______________ an atom has for a shared pair of electrons.
o
___________ e-neg atom  -
o
___________ e-neg atom +
Electronegativity Difference
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If the difference in electronegativities is between:
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1.7 to 4.0: ___________
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Greater than 0.3 & less
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than 1.7: _____ Covalent
0 to 0.3: __________
________________
The type of bond can usually be calculated by finding the difference in electronegativity of the two atoms
that are bonded.
o
Compound:
F2 or F-F
CF4
LiF or Li - F
o
Difference:
_______
_______
_______
o
Type of Bond:
_______
_______
_______
Nonpolar Covalent Bond
o
e- are shared _____________
o
________________ e- ______________
o
usually _____________ atoms
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Ex: H2 or Cl2
Polar Covalent Bond
o
e- are shared ________________
o
_____________ e- density
o
results in ________ ___________ (dipole)
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Ex: H2O
VSPER Theory
o
________________________________________________________
o
Electron pairs ____________ themselves in order to _______________ repulsive forces
Types of e- Pairs
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__________________ – form bonds
o
Lone pairs – _______________ e-
o
Total e- pairs– _________ + _______ pairs
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Lone pairs ________ more ____________ than bonding pairs!!!
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Lone pairs __________ the bond angle between atoms
Determining Molecular Shape
o
Draw the ____________ Diagram
o
Tally up e- pairs on central atom (bonds + lone pairs)
o
double/triple bonds = ONE pair
o
Shape is determined by the _________ of bonding pairs and lone pairs
Common Molecular Shapes:
Ex: BeH2
Ex: BF3
Molecular Polarity:
 ___________ molecule = one end slightly __________ and one end slightly __________
 Molecule with 2 poles = ___________ molecule or ___________
 _____________, symmetry and bond polarity determines molecular polarity
 H – O bond is __________ and water is ______________, so H2O is polar
 C – Cl bond is polar, but CCl4 is _________________, so molecule is _____________
Identify each molecule as polar or nonpolar
o
O2 : Nonpolar Bonds __________
o
CS2 : Linear  ______________
o
CF4 : Tetrahedral  __________
o
H2O : Bent  _______________