covalent compounds. result from the sharing of electrons between two atoms ◦ a two-electron bond...
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Chapter 4Covalent Compounds
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result from the sharing of electrons between two atoms ◦ A two-electron bond in which the bonding atoms
share the electrons A molecule is a discrete group of atoms held
together by covalent bonds
Covalent Bonding
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Unshared electron pairs are called nonbonded electron pairs or lone pairs
Atoms share electrons to attain the electronic configuration of the noble gas closest to them in the periodic table◦ Main group elements share e- until they reach an
octet of e- in their outer shell◦ H shares 2 e-
Covalent Bonding
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Drawing Covalent Bonds
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Covalent bonds are formed when two nonmentals combine or when a metalloid bonds to a nonmetal
Atoms with one, two, or three valence e- form one, two, or three bonds respectively
Atoms with four or more valence e- form enough bonds to achieve an octet
Predicting the number of bonds
predicted number of bonds
= 8 – number of valence e−
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Predicting the number of lone pairs
Number of bonds Number of lone pairs+ = 4
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A molecular formula shows the number and identity of all of the atoms in a compound, but not which atoms are bonded to each other.
A Lewis structure shows the connectivity between atoms, as well as the location of all the bonding and nonbonding valence electrons ◦ General rules
Draw only valence electrons. Give every main group element (except H) an octet of e−
Give each hydrogen 2 e−
Lewis Dot Structures
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Lewis Dot StructuresStep [1]
Arrange the atoms next to each other that you think are bonded together.
Place H and halogens on the periphery, since they can only form one bond.
Step [2]
Count the valence electrons.
The sum gives the total number of e− that must be used in the Lewis structure.
Step [3] Arrange the electrons around the atoms.
Place one bond (two e−) between every two atoms.Use all remaining electrons to fill octets with lone pairs, beginning with atoms on the periphery.
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For CH3Cl◦ C brings 4 valence electrons = 4 e-
◦ Each H brings 1 valence electrons = 3 X 1 = 3 e-
◦ Cl brings 7 valence electrons = 7e-
Final diagram needs to have all 14 e- accounted for
H only forms one bond Cl (a halogen) only forms one bond Therefore start with C in the middle
Example
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Lewis Dot Structures
For CH3Cl:
C ClH
H
H
8 e−
on Cl2 e− oneach H
14 e−
4 bonds x 2e− = 8 e−
+ 3 lone pairs x 2e− = 6 e−
All valence e− have been used.
If all valence electrons are used and an atom still does not have an octet, proceed to Step [4].
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A double bond contains four electrons in two 2-e- bonds
A triple bond contains six electrons in three 2-e- bonds
Lewis Dot Structures
Step [4]
Use multiple bonds to fill octets when needed.
O O
N N
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[CO3]2-
◦ O brings 6 valence electrons = 3 X 6 = 18 e-
◦ Each C brings 4 valence electrons = 4 e-
◦ Overall negative charge adds 2 electrons = 2e-
Final diagram needs to have all 24 e- accounted for
Carbon can make 4 bonds so will start with C in the middle
Example
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I start by putting single bonds in place and filling out the rest of the electrons but C ends up without an octet around it even with the 24 e- all accounted for
Now will try making one of them a double bond
Carbonate example continued
C
O
OO
2-
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When I make one of the bonds a double bond I get an octet around C
I double check the oxygen and the total electron number and everything checks out
Therefore I am done and do not need to explore triple bonds
Carbonate example continued
C
O
OO
2-
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H is a notable exception, because it only needs 2 e- in bonding
Elements in group 3A do not have enough valence e- to form an octet in a neutral molecule
Exceptions
only 6 e− on B
B
F
FF
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Elements in the third row have empty d orbitals available to accept electrons
Thus, elements such as P and S may have more than 8 e- around them
Execptions
10 e− on P 12 e− on S
S
O
OHHO
O
P
O
OHHO
OH
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When drawing Lewis structures for polyatomic ions◦ Add one e- for each negative charge◦ Subtract one e- for each positive charge
Polyatomic ions
−
Each atomhas an octet.
Answer
C N
For CN– :
C N
1 C x 4 e− = 4 e−
1 N x 5 e− = 5 e−
–1 charge = 1 e−
10 e− total
All valence e−
are used, but C lacks an octet.
C N−
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Two Lewis structures having the same arrangement of atoms but a different arrangement of electrons
Two resonance structures of HCO3-
Neither Lewis structure is the true structure of HCO3
-
The true structure is a hybrid of the two resonance structures
Resonance Structures
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NamingHOW TO Name a Covalent
MoleculeExampl
eName each covalent molecule:
(a) NO2 (b) N2O4
Step [1]
Name the first nonmetal by its elementname and the second using the suffix“-ide.”
(a) NO2
nitrogen oxide
(b) N2O4
nitrogen oxide
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NamingStep [2]
Add prefixes to show the number of atoms of each element. Use a prefix from Table 4.1 for each element.
(a) NO2
nitrogen dioxide
(b) N2O4
dinitrogen tetroxide
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The prefix “mono-” is usually omitted.◦ Exception: CO is named
carbon monoxide If the combination
would place two vowels next to each other, omit the first vowel.◦ mono + oxide =
monoxide
Naming rules
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To determine the shape around a given atom, first determine how many groups surround the atom
A group is either an atom or a lone pair of electrons
Use the VSEPR theory to determine the shape◦ Valence shell electron pair repulsion
The most stable arrangement keeps the groups as far away from each other as possible
Molecular shape
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Any atom surrounded by only two groups is linear and has a bond angle of 1800
An example is CO2
Ignore multiple bonds in predicting geometry◦ Count only atoms and lone pairs
Molecular Shape
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Any atom surrounded by three groups is trigonal planar and has bond angles of 1200
An example is H2CO
Molecular Shape
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Any atom surrounded by four groups is tetrahedral and has bond angles of 109.50
An example is CH4
Molecular Shape
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If the four groups around the atom include one lone pair, the geometry is a trigonal pyramid with bond angles of 109.50
An example is NH3
Molecular Shape
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If the four groups around the atom include two lone pairs, the geometry is bent and the bond angle is 1050 (close to 109.50)
An example is H2O
Molecular Shape
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Molecular Shape
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Electronegativity is a measure of an atom’s attraction for e- in a bond
Polarity
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If the electronegativities of two bonded atoms are equal or similar, the bond is nonpolar
The electrons in the bond are being shared equally between two atoms
Polarity
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Bonding between atoms with different electronegativities yields a polar covalent bond or dipole
The electrons in the bond are unequally shared between the C (2.5) and the O (3.5)
e- are pulled toward O, the more electronegative element, this is indicated by the symbol δ−.
e- are pulled away from C, the less electronegative element, this is indicated by the symbol
Polarity
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Polarity
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Nonpolar molecules generally have◦ No polar bonds◦ Individual bond dipoles that cancel
Polar molecules generally have◦ Only one polar bond◦ Individual bond dipoles that do not cancel
Polar and Nonpolar
33Smith. General Organic & Biolocial Chemistry 2nd
Ed.