COVALENT COMPOUNDS
Chapter 6
Unit Essential Question:
1) HOW DO COVALENT
COMPOUNDS DIFFER FROM
IONIC COMPOUNDS?
Lesson Essential Questions:
1) WHY DO ATOMS SHARE
ELECTRONS?
2) HOW DO BONDS AND
PROPERTIES RELATE?
Section 1: Covalent Bonds

Covalent- electrons are shared.

Simplest covalent bonds are diatomic
molecules.
Always 2 atoms; atoms can be different or
the same (diatomic elements).
 Ways to remember diatomic elements:
BrINClHOF, HON & the halogens, #7


Ex: H2
 Each H atom shares its valence
electron so each can have 2.
 Called a shared pair of electrons..
Molecular Orbitals

Where shared electrons are located.


Region of high probability of finding the
shared electrons- recall quantum model.
Form when covalent bonds occur and
atomic orbitals overlap.
Energy and Stability

Non-bonded atoms have high PE and low stability.
 Become more stable when bonded.
 E is released when bonds form, PE decreases.
 This determines the bond length.

Ideal length is when the two bonded atoms are at
their lowest PE.
Energy and Stability Cont.

Covalent bonds are flexible.



Bond energy – energy required to
break a bond.


Bond length is an average distance.
Bond length oscillates like a spring
between the two nuclei.
Different from lattice energy!
More E is needed to break a shorter
bond.

Shorter bond length = stronger bond.
Re-visiting Electronegativity

How do we know if atoms will transfer
or share electrons in a bond?


Check electronegativity values!
If the difference is greater than 2.1 it is
considered ionic.
 Ex: KCl
Electronegativities: K= 0.8 and Cl= 3.2
Subtract: 3.2-0.8 = 2.4, so the bond is
ionic.
Re-visiting Electronegativity



If the difference is less than 2.1 it is
considered covalent.
Ex: CO2
Electronegativities: C=2.6 and O=3.4
Subtract: 3.4 – 2.6 = 0.8, so the bond is
covalent.
You do not need to memorize
electronegativity values! They will be given
to you on tests and quizzes.

pg.194 Figure 6 shows electronegativity
values
Re-visiting Electronegativity




It is important to note that this cut-off
value of 2.1 is arbitrary! Therefore, it is
not 100% accurate.
Properties of the compound must be
investigated for better classification.
Ex: magnesium chloride
Electronegativities: Mg=1.3 and Cl=3.2
Subtract: 3.2 – 1.3 = 1.9
It looks like it should be a covalent bond,
but properties indicate it is actually ionic.
Polar vs. Nonpolar Covalent


If the electronegativity difference is less
than 0.5 we will consider the bond to be
nonpolar covalent.
 Means the atoms are essentially sharing
the electrons equally.
If the difference is between 0.5 and 2.1,
we will consider it to be polar covalent.
 Means one atom attracts the electrons
more towards itself; there is unequal
sharing of electrons.
Determining Bond Type
Practice Classifying Bonds
A) Classify the following bonds as ionic, polar
covalent, or nonpolar covalent:
1)
2)
3)
4)
5)
Li – Cl 2.2 ionic
B – C 0.6 polar covalent
N – O 0.4 nonpolar covalent
Mg – Br 1.7 ionic
C – F 1.4 polar covalent
B) Since F has the highest
electronegativity value, can F
ever form a nonpolar covalent
bond? Why or why not?
Simple Polar Molecules


Ions have full charges (electrons are
transferred).
Polar molecules have partial charges
(unequally shared electrons).
 Use Greek symbol delta δ (and +/-) to
indicate partial charges.
 Dipole – molecule that contains both
positive and negative partial charges.

Ex: HCl
Simple Polar Molecules Continued

Use electronegativity values to determine
partial charges.
Atom with larger e-neg value = δ Atom with lower e-neg value = δ+
 Ex: HCl
Elecronegativities: H=2.2 and Cl=3.2
So: H is δ+ and Cl is δ In other words, the electrons are more likely
to be found at the Cl atom than the H
because it has a larger electronegativity;
this makes Cl more negative.

Different Properties for Different Bonds

Besides electronegativity, you can
predict bond type by the type of
elements involved in the bond.




Metallic = metal atoms (K, Cu, etc.)
Ionic = 1 metal (typically) + 1 nonmetal
Covalent = 2 or more nonmetals
Properties of these compounds are
determined by bond type.

Recall ionic compound properties and
how they stem from the crystal lattice.
Covalent Properties

Some are soluble in water and some are
not.



Depends if the bonds are polar or not.
Poor conductivity.
Covalent compounds can exist as a solid,
liquid, or gas.

Depends upon the polarity of bonds.
Tend to have low melting and boiling
points.
*Remember: covalent compounds are made
of molecules, not ions!

Bond type
Metallic
Ionic
Covalent
Potassium
Potassium
chloride
Chlorine
Melting point
63 oC
770 oC
-101 oC
Boiling point
760 oC
1500 oC
(sublimes)
-34.6 oC
Model
Example
Properties
•Soft, grey
lustrous
•Conductor as
solid
•Crystalline
white solid
•Conductor
when molten or
in solution
•Green-yellow
gas
•Insulator
Lesson Essential Question:
HOW ARE COVALENT
COMPOUNDS REPRESENTED?
Section 2: Drawing and Naming Molecules

Lewis structures (electron dot diagrams)
use dots and lines to show valence
electrons.



Atoms only use dots:
 Each dot represents a valence e-.
Molecules use both dots and dashes:
 Each line represents a bond.
Help to determine and show bonds that will
form in covalent compounds.
Lewis Structures- Atoms

The 4 sides of the element symbol
represent the s and three p orbitals.




s = 2 e- and p = 6 e- so total = 8
No more than 8 dots per symbol
Put one dot on each side of the
atom before pairing:
Draw the Lewis structure for the
following: Al, Br, N, Ne
Lewis Structures- Molecules

When drawing Lewis structures for
molecules, give each atom an octet.


Needs 4 e-,
so it will form
4 bonds.

Except for H!
Looking at the Lewis structures for
each atom can help to determine how
many bonds it will form.
Ex: CH4
Each H
needs 1 e-,
so they will
form 1 bond.
Electron Pairs
•
Shared pairs: e- are shared between
two atoms, forming a bond.



•
1 shared pair = single bond
2 shared pairs = double bond
3 shared pairs = triple bond
or
Unshared/lone pairs: e- are not shared
between two atoms, and are not involved
in a bond.
Steps for Drawing Lewis Structures
1) Find the total # of valence electrons for all
atoms in the compound.
Ex: NH3
N: 5e-
H: 1e-
total: 5 + 1(3) = 8e-
2) Arrange atoms/determine ‘backbone’ of the
compound.
H
and halogens tend to be on the end.

C
They don’t like to share more than 1 pair of e-.
is almost always in the center, if present.

Likes to form 4 bonds.
Except
for C, the atom with the lowest
electronegativity will be the center.
H will NEVER be in the center!
 Ex: NH3
N will be in the center:
Steps for Drawing Lewis Structures
3) Keep track of how many e- are used and how
many remain.
Each bond used to build the backbone uses
2 electrons. NH3 has 3 bonds, so 6 e- used.
 8 e- - 6 e- = 2 e- left.

4) Distribute remaining e- by giving each atom
an octet.
Give the most electronegative atoms e- first.
 Ex: 2 e- left; N is the only one that needs an
octet, so it gets the remaining 2 e-.
5) Verify the structure!
 Make sure all atoms have an octet and check that
all valence electrons have been used.

Polyatomic Ions
1) The bonds that hold together polyatomic ions are
covalent bonds.
2) When drawing the Lewis structure for a
polyatomic ion, you must take into account its
charge when determining the total number of
valence electrons you’re working with:
• negative ion: add electrons to the total
• positive ion: subtract electrons from the total
3) Finally, put brackets around the ion and include
its charge.
Ex: Draw the Lewis structure for
the sulfate ion.
Practice Problem #2: Ions
Draw the Lewis structure for the
ammonium ion.
Practice Problem #4: Multiple Bonds
1) It is possible to run out of remaining
electrons to give to atoms as lone pairs.

But they still need an octet!
2) If this happens, you take lone pairs from a
neighboring atom and form another bond to
the atom that needs an octet.


The neighboring atom donates both electrons to
be shared in this bond.
This is how double and triple bonds are formed.
Ex: Draw the Lewis structure for O2.
Practice Problems #5-6: Multiple Bonds

Draw the Lewis structure for CO2.

Draw the Lewis structure for N2.
Resonance Structures
•1) Some compounds can be represented with
more than one Lewis structure.
 All Lewis structures are drawn.
 A double-headed arrow is put between them
to indicate all possibilities.
•2) Compound is an average of the possible
Lewis structures (resonance hybrid).
Naming Binary Covalent Compounds
1) Name the first element listed in the
formula.
 The name does not change.
2) Second element always has ‘-ide’ ending.
3) Prefixes are used to indicate numbers of
atoms of both elements.
 Exception: do not use ‘mono-’ for the first
element listed if there is only one.
Naming Binary Covalent Compounds
Take off the ‘a’ at the end of the prefix if
the element begins with a vowel.
Ex: pentoxide, not pentaoxide.
Ex #1: CCl4
carbon tetrachloride
Ex #2: P2O5
diphosphorus pentoxide
Ex #3: H2O
dihydrogen monoxide
Ex #4: CO
carbon monoxide
Lesson Essential Question:
HOW CAN THE THREE
DIMENSIONAL SHAPE OF A
MOLECULE BE DETERMINED?
WHY ARE THE SHAPES OF
MOLECULES USEFUL?
Section 3: Molecule Shapes
1) VSEPR theory: Valence Shell Electron Pair
Repulsion
 Predicts shape based on electron repulsion.
 e- want to be as far apart as possible!
 Look at the central atom to determine the shape.
2) Unshared pairs repel more than shared
pairs.
 All electrons repel each other, whether
they are bonded or lone pairs.
3) In addition to bond polarity, shape also
plays a role in determining the properties
of a substance.
Predicting Shapes


Count the number of shared and unshared pairs
on the central atom in the Lewis structure.
Count double/triple bonds as 1 shared pair.
Total pairs
Shared pairs
Unshared
pairs
Shape name
4
4
0
Tetrahedral
4
3
1
Trigonal
pyramidal
4
2
2
Bent
3
2
1
Bent
2
2
0
Linear
3
3
0
Trigonal planar
Angles



Bond angles also show repulsion of electrons
in molecules.
Tetrahedral angles = 109.5o
Base all other angles off of this.
 From here, we can predict >, < , or
= 109.5o
 Depends upon shared vs. lone pairs.
(Additional factors that affect bond
angles: atom size, multiple bonds take
up more space. We won’t worry about
these.)
Angles Cont.




Trigonal Pyramidal: <109.5o
 Remember- tetrahedral was 109.5, now
we’ve replaced a shared pair with a lone pair
that repels more! So angles decrease!
Linear: > 109.5o (180o is much bigger!)
Trigonal Planar: >109.5o (120o is bigger!)
Bent – depends on # of lone pairs


Bent with one lone pair is like taking trigonal
planar and replacing a shared pair with a lone pair,
so angle is <120 but >109.5.
Bent with two lone pairs is like taking tetrahedral
and replacing two shared pairs with two lone pairs,
so angles are much <109.5. This is most common!
Polarity

Tetrahedral, trigonal planar, and
linear molecules can be nonpolar.



Polar bonds can ‘cancel’ out other polar
bonds because of the ‘symmetrical’
shape.
This happens if the bonds are all the
same.
 Ex: CH4 & CO2
They can also be polar if the bonds
are not all the same.

Bonds do not completely cancel out.
 Ex: CH3F
Polarity Cont.

Trigonal pyramidal and bent molecules
tend to be polar.




Polar bonds are not cancelled out because of
the ‘asymmetrical’ shape.
Example: H2O
Even if the bonds are not polar, the lone pair
makes it polar.
Let’s relate this back to covalent
compound properties from section 1.


Methanol is soluble in water.
Oil is not soluble in water.