Ch. 01 and 02: Structure and Bonding of Organic Molecules
I. Types of Chemical Bonds
A. Why do atoms forms bonds?
Atoms want to have the same number of electrons as the nearest noble gas atom (noble gas configuration). This requires having a completely full or completely empty valence shell of electrons.
Most main groups atoms will try to have eight valence electrons to completely fill their valence shell; this is the octet rule. There are some exceptions: hydrogen only needs two electrons (duet rule), and boron and aluminum are stable with only six valence electrons (incomplete octet).
Atoms form bonds in order to gain or lose electrons and satisfy the octet rule.
B. Two Types of Bonds
Atoms have two major strategies for gaining (or losing) electrons to fulfill the octet rule. These two different strategies lead to the two main types of chemical bonds.
1. Ionic Bonds
Na
One strategy is to completely transfer valence electrons between atoms:
Cl
Cl
Na
Oppositely charged ions attract (ionic bond)
This forms oppositely charged ions, which are then attracted and form ionic bonds.
2. Covalent Bonds
Cl
The other strategy is to share valence electrons between atoms:
Cl
Cl Cl
Atoms must stay together to share valence electons (covalent bond)
This strategy forces the atoms to remain in close proximity to continue sharing electrons.
Ch. 01 and 02: Structure and Bonding of Organic Molecules
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C. The type of bond formed depends on the electronegativity difference of the atoms
Electronegativity: the strength with which an atom attracts electrons. Electronegativity is a periodic property which increases from left to right and from bottom to top of the periodic table: electronegativity increases
H
2.1
Li
1.0
Be
1.6
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
Na
0.9
Mg
1.2
Al
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
K
0.8
Br
2.8
I
2.4
Atoms with an electronegativity difference larger than about 2.0 form ionic bonds.
This usually happens between atoms from opposite sides of the periodic table (between metals and non-metals).
II. Lewis Structures
Lewis structures show the connections between atoms that form covalent bonds. They are created by connecting atoms so that they share electrons to have octets.
In explicit Lewis structures, all atoms are shown using their atomic symbols, all covalent bonds are shown as dashes, and all unshared electrons are shown as dots.
To be correct, a Lewis structure must have the correct number of valence electrons, and in most cases should not violate the octet rule.
Ch. 01 and 02: Structure and Bonding of Organic Molecules
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III. Formal Charges
Formal charges must be calculated for each atom of a given Lewis structure. There are patterns for the formal charges of the important atoms in organic chemistry, and eventually we will begin to predict them without calculating.
To calculate the formal charge on an atom in a Lewis structure:
1. Draw a complete Lewis structure of the compound of interest. Include all lone pairs and unpaired electrons.
2. For each atom, determine the number of valence electron that the atom “wants” to own. This will be equal to the group number of that atom in the periodic table, or the number of valence electrons on the neutral atom.
3. Determine the number of valence electrons that the atom actually “owns” in the
Lewis structure:
An atom owns two electron for each of its lone pairs.
An atom owns one electron for each bond connected to it (a double bond counts as two bonds; a triple bond counts as three bonds).
