Covalent bonds

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Covalent Bonds
The joy of sharing!
Covalent Bonds


Covalent bonds: occur between
two or more nonmetals; electrons
are shared not transferred (as in
ionic bonds)
The result of sharing electrons is
that atoms attain a more stable
electron configuration.
Covalent Bonds

Most covalent bonds involve:



2 electrons (single covalent bond),
4 electrons (double covalent bond, or
6 electrons (triple covalent bond).

Lewis structures (electron dot structures)
show the structure of molecules. (Bonds
can be shown with dots for electrons, or
with dashes: 1 dash = 2 electrons)

H2
HBr

CCl4
O2

N2
CO

Lewis structures (electron dot structures)
show the structure of molecules. (Bonds
can be shown with dots for electrons, or
with dashes: 1 dash = 2 electrons)

H2
HBr

CCl4
O2

N2
CO

Lewis structures (electron dot structures)
show the structure of molecules. (Bonds
can be shown with dots for electrons, or
with dashes: 1 dash = 2 electrons)

H2
HBr

CCl4
O2

N2
CO

Lewis structures (electron dot structures)
show the structure of molecules. (Bonds
can be shown with dots for electrons, or
with dashes: 1 dash = 2 electrons)

H2
HBr

CCl4
O2

N2
CO

Lewis structures (electron dot structures)
show the structure of molecules. (Bonds
can be shown with dots for electrons, or
with dashes: 1 dash = 2 electrons)

H2
HBr

CCl4
O2

N2
CO

Lewis structures (electron dot structures)
show the structure of molecules. (Bonds
can be shown with dots for electrons, or
with dashes: 1 dash = 2 electrons)

H2
HBr

CCl4
O2

N2
CO
Octet Rule

Octet Rule: The representative
elements achieve noble gas
configurations (8 electrons) by
sharing electrons.
Writing Lewis Structures
1)
2)
3)
4)
5)
6)
Select a skeleton for the molecule (the least
electronegative element is usually the
central element).
Calculate N (the # of valence e- need by all
atoms in the molecule of polyatomic ion.
Calculate A (the # of electrons available).
Calculate S (the # of electrons shared in the
molecule) S = N – A
Place the S electrons as shared pairs in the
skeleton.
Place the additional electrons as unshared
pairs to fill the octet of every representative
elements (except hydrogen!).
Lewis Structure Examples:

CO2-

CBr3-

OH-

N22-

NO3-

CO32-

SO42-

NH4+
Lewis Structure Examples:

CO2-

CBr3-

OH-

N22-

NO3-

CO32-

SO42-

NH4+
Lewis Structure Examples:

CO2-

CBr3-

OH-

N22-

NO3-

CO32-

SO42-

NH4+
Lewis Structure Examples:

CO2-

CBr3-

OH-

N22-

NO3-

CO32-

SO42-

NH4+
Lewis Structure Examples:

CO2-

CBr3-

OH-

N22-

NO3-

CO32-

SO42-

NH4+
Lewis Structure Examples:

CO2-

CBr3-

OH-

N22-

NO3-

CO32-

SO42-

NH4+
Lewis Structure Examples:

CO2-

CBr3-

OH-

N22-

NO3-

CO32-

SO42-

NH4+
Lewis Structure Examples:

CO2-

CBr3-

OH-

N22-

NO3-

CO32-

SO42-

NH4+
Electronegativity

We’ve learned how valence
electrons are shared to form
covalent bonds between elements.
So far, we have considered the
electrons to be shared equally.
However, in most cases, electrons
are not shared equally because of a
property called electronegativity.
Electronegativity


The ELECTRONEGATIVITY of an
element is: the tendency for an
atom to attract electrons to itself
when it is chemically combined with
another element.
The result: a “tug-of-war”
between the nuclei of the
atoms.
Electronegativity


Electronegativities are given numerical
values (the most electronegative element
has the highest value; the least
electronegative element has the lowest
value)
**See Figure 6-18 p. 169 (Honors)


Most electronegative element:
Fluorine (3.98)
Least electronegative elements:
Fr (0.70), Cs (0.79)
Electronegativity

Notice the periodic trend:



As we move from left to right across a
row, electronegativity increases (metals
have low values nonmetals have high
values – excluding noble gases)
As we move down a column,
electronegativity decreases.
The higher the electronegativity
value, the greater the ability to
attract electrons to itself.
Nonpolar Bonds


When the atoms in a molecule are
the same, the bonding electrons are
shared equally.
Result: a nonpolar covalent bond

Examples: O2, F2, H2, N2, Cl2
Polar Bonds



When 2 different atoms are joined by a
covalent bond, and the bonding electrons
are shared unequally, the bond is a polar
covalent bond, or POLAR BOND.
The atom with the stronger electron
attraction (the more electronegative
element) acquires a slightly negative
charge.
The less electronegative atom
acquires a slightly positive charge.
Polar Bonds

Example: HCl

Electronegativities:


H = 2.20
Cl = 3.16
+ H
Cl
-
Polar Bonds

Example: H2O

Electronegativities:


H = 2.20
O = 3.44
Polar Bonds

Example: H2O

Electronegativities:


H = 2.20
O = 3.44
Polar Bonds

Example: H2O

Electronegativities:


H = 2.20
O = 3.44
Predicting Bond Types

Electronegativities help us
predict the type of bond:
Electronegativity
Difference
0.00 – 0.40
0.41 – 1.00
1.01 – 2.00
2.01 or higher
Type of Bond
covalent
(nonpolar)
covalent
(slightly polar)
covalent
(very polar)
ionic
Example
H-H
H-Cl
H-F
Na+Cl-
Polar Molecules


A polar bond in a molecule can
make the entire molecule polar
A molecule that has 2 poles
(charged regions), like H-Cl, is
called a dipolar molecule, or dipole.
Polar Molecules


The effect of polar bonds on the polarity of
a molecule depends on the shape of the
molecule.
Example:
O=C=O
CO2
shape: linear
*The bond polarities cancel because they
are in opposite directions; CO2 is a
nonpolar molecule.
Polar Molecules


The effect of polar bonds on the
polarity of a molecule depends on the
shape of the molecule.
Water, H2O, also has 2 polar bonds:


But, the molecule is bent, so the bonds
do not cancel.
H2O is a polar molecule.

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