Warm-up Questions...
With Your Partners!
Chapter 8:
Covalent Bonding
Please discuss and answer the following
questions:
Ms. Nguyen
1) What type of bond occurs between
Sodium (Na) and Chlorine (Cl)? Why do
you think this?
2) Is Si and O an Ionic Bond? Why or
why not?
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8.1 Molecular Compounds
In the previous chapter you learned how metals
and nonmetals lose or gain electrons in an ionic
bond.
This chapter will introduce another type of
bonding that occurs between only nonmetals.
» Covalent Bonding: Atoms held together by
sharing electrons between nonmetals.
Monatomic vs. Diatomic Molecules
Molecules
Molecules can be monatomic or diatomic
In the previous chapter you learned that a metal
cation and a nonmetal anion are joined together by
an ionic bond (called a salt).
Monatomic Molecule: A molecule that consists of one
atom.
A neutral group of atoms joined together by a
covalent bond is called a molecule.
Example 1: Noble Gases (He, Ne, Ar, Kr, Xe, Rn)
Examples: Water (H2O), Carbon Dioxide (CO2),
Under the right conditions (STP), they can exist in the form
of a monatomic molecule because they are already stable.
Ammonia (NH3)
Example 2: Na+ (Sodium ion)
This is a monatomic ion because there’s a positive charge
after the elemental symbol, and there’s only one atom.
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Monatomic vs. Diatomic Molecules
Diatomic Molecule: A molecule that consists of two
atoms.
There are seven diatomic molecules: H2, N2, O2, F2, Cl2,
Br2, I2 .
They exist as diatomic molecules because they cannot
exist alone, to stabilize themselves they combine with
another atom of the same element.
When looking at the periodic table to find the 7 diatomic
molecules, remember that “H” is one of them and the
other six form a number seven and start at (N) Nitrogen.
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Molecular Compounds
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Properties of Molecular Compounds
Molecular compounds tend to have relatively lower
melting and boiling points than ionic compounds.
Atoms of different elements can combine
chemically to form compounds.
Many molecular compounds are gases or liquids at
room temperature.
A compound composed of molecules is called a
molecular compound.
They involve nonmetals.
The bonds don’t involve ions, so they are
nonconductors of electricity.
The molecules of a given molecular compound are
all the same.
Given what you know about ionic bonds, discuss with
your partners two different properties between ionic
and covalent bonds?
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Molecular Formula Practice
Molecular Formulas
How many atoms total and of each
element do the following molecular
compounds contain?
Molecular formulas: The chemical formula
of a molecular compound. It shows how many
atoms of each element a molecule contains.
The subscripts written after the symbol
indicate the number of atoms of each element
in the molecule.
Examples:
H 2O
3 atoms (2 Hydrogens; 1 Oxygen)
C2H6
8 atoms (2 Carbons; 6 Hydrogens)
• 2BF3
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8 atoms [(1 Boron; 3 Fluorines) x 2]
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1. H2
2 atoms (2 Hydrogen atoms)
2. 2CO
4 atoms [(1 Carbon atom; 1 Oxygen atom) x 2]
3. CO2
3 atoms (1 Carbon atom; 2 Oxygen atoms)
4. 3NH3
12 atoms [(1 Nitrogen atom; 3 Hydrogen atoms) x 3]
5. C2H6O
9 atoms (2 Carbon atoms; 6 Hydrogen atoms;
1 Oxygen atom)
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Warm-Up Questions
With Your Partners!
Please discuss and answer the following questions:
Section 8.2:
1. How many atoms total and of each element
does PCl5 (Phosphorus Pentachloride) contain?
The Nature of Covalent Bonding
2. How many atoms total and of each element
does 2BF3 contain?
Ms. Nguyen
3. What are the seven diatomic molecules and
how can you identify them on the periodic
table?
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8.2 The Nature of Covalent Bonding
Remember that ionic compounds either
gain or lose electrons in order to attain a
noble gas electron configuration.
Ionic
Bonding
-->
Covalent compounds form by sharing
electrons to attain a noble gas electron
configuration.
(Nonpolar) Covalent
Bonding
<--
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Review of Octet Rule and Valence Electrons
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Single Bonds
Two atoms held together by sharing a pair of electrons are
joined by a single bond.
Octet Rule: Atoms react by gaining or losing
electrons so as to acquire the stable electron
structure of a noble gas, usually eight electrons.
Example: Hydrogen gas (H2) consists of diatomic molecules
whose atoms share only one pair of electrons, forming a single
covalent bond.
The octet rule still applies to covalent bonds.
Atoms usually acquire a total of eight electrons
by sharing electrons.
An electron dot structure such as H H for Hydrogen
gas represents the shared pair of electrons of the
covalent bond by two dots.
Valence Electrons: The electrons in the highest
occupied energy level of an element’s atoms.
Hydrogen gas can also be represented by a dash H-H in a
structural formula.
Example: Water (H2O)
The molecular formula for Hydrogen gas is H2 which
indicates the number of atoms in the molecule.
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Practice With Single Bonds
Let’s Practice Drawing!
Example 1: F2
Please draw the structures for each of the
covalently bonded atoms. Also list the amount
of lone pairs that may exist for each?
Solution: We know that the Fluorine atom has 7 valence
electrons and it only needs 1 more to fulfill the octet
rule (usually 8 electrons).
Notice that the two fluorine atoms share only one pair
of valence electrons.
1. Cl2
Important note: When drawing structural formulas for
molecules, we want to draw symmetrical structures.
2.H2
A lone pair is a pair of valence electrons that is not
shared between atoms.
3.I2
Example 2: F2
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Let’s Review!
Multiple Covalent Bonds
Atoms form double or triple covalent bonds if they
can attain a noble gas structure by sharing two
pairs or three pairs of electrons.
Please draw the electron dot structure for each
of the following atoms:
Double bond: A bond that involves two shared
pairs of electrons.
Example: CO2
Triple bond: A bond formed by sharing three pairs
of electrons.
Example: N2
B
N
Al
He
Mg
Be
Li
Ar
P
Question: Will Hydrogen ever form double or triple
bonds? Why or why not?
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Lewis (Dot) Structures- 5 Steps
Lewis (Dot) Structures- 5 Steps
Step 4: Fill the valence shells
Step 1: Determine the number of valence electrons
in the Lewis structures
Look at the chart from step 1 and see if the atoms
have the correct number of valence electrons.
Step 2: Determine the number of bonds in the
structure
Step 5: Check your structure
# of electrons needed = # of bonds in molecule
2 electrons per bond
A) Do we have the same number of total valence
electrons?
Step 3: Draw a simple and symmetrical Lewis
structure
B) Does each atom fulfill the octet rule?
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Lewis (Dot) Structures
Step 4: Fill the valence shells
Example: H2O
Step 1: Atoms
Valence e-
Look at the chart from step 1, do the atoms have the correct
number of valence electrons? If not, remember to include lone pairs
to correct the structure.
# e- needed
H
1
1
H
1
1
O
6
2
Total
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H—O—H
Step 5: Check your structure by answering the following questions:
a) Do we have the same number of total valence electrons as step 1?
2 bonds + 2 lone pairs = 8 electrons
Step 2: 4 e- needed/ 2 e- per bond = 2 bonds
b) Does each atom fulfill the octet rule?
Step 3: Draw the symmetrical structures.
H—H—O
H—O—H
Each Hydrogen has 2 valence electrons around it, and Oxygen has 8.
O—H—H
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Polyatomic Ions
Let’s Practice Drawing!
Example 1: Cl2
Polyatomic Ions: A tightly bound group of
atoms that has a positive or negative
charge and behaves as one unit.
Example 2: CO2
Example: NH4+ (Ammonium Ion)
The ammonium ion forms when a positively
charged hydrogen ion (H+) attaches to the
unshared electron pair of an ammonia
molecule (NH3).
Example 3: NF3
Example 4: CH4
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Bond Dissociation Energies
Bond Dissociation Energy: The energy
required to break the bond between two
covalently bonded atoms.
Section 8.4:
A strong bond dissociation energy
corresponds to a strong covalent bond.
Polar Bonds and Molecules
Ms. Nguyen
Example: Carbon- Carbon has a strong
bond dissociation so it’s not very reactive.
It’s unreactive because the dissociation
energy for each of these bonds is high.
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Warm-Up Questions
With Your Partners!
8.4 Polar Bonds and Molecules
There are two types of covalent bonds:
•
a)
Please discuss and answer the following
questions:
Nonpolar covalent bonds
b) Polar covalent bonds (or polar bonds).
The bonding pairs of electrons in (nonpolar) covalent
bonds are pulled equally.
•
How many valence electrons does
oxygen have?
•
How do you draw the electron dot
structure for methane (CH4)?
Example: Cl2, N2, H2
Diatomic molecules are always nonpolar
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Polar Covalent Bonds (or polar bonds): A covalent bond
between atoms in which the electrons are shared
unequally.
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Electronegativity
Electronegativity: The tendency of an atom
in a molecule to attract electrons to itself.
The more electronegative atom attracts
electrons more strongly and gains a
slightly negative charge.
Electronegativity values generally increase across a period table
(from left to right), and also decreases going down a group.
The less electronegative atom has a
slightly positive charge.
Nonmetals have the highest electronegativity values.
Metals have the lowest electronegativity values.
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Electronegative Differences
Classification of Bonds
Example: HCl
You can determine the type of bond artificially by
calculating the difference in electronegativity between
elements
Step 1: Using your chart, find the electronegative
values for Hydrogen and Chlorine.
Important note: Consider what you know about ionic
and covalent bonds as well!
Type of Bond! !
Step 2: Subtract the larger electronegative value from
the smaller electronegative value.
Electronegativity Difference
Nonpolar Covalent!
0 ! 0.4
Polar Covalent!!
0.5 ! 1.9
Ionic! ! !
2.0 ! 4.0
Step 3: Use the “Electronegativity Difference Table” to
understand the value difference.
Cl - H = Electronegativity Difference
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! 3.0 - 2.1 = 0.9 indicating a polar covalent bond
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Let’s Practice!
Let’s Practice!
1.
N and H
3.0 - 2.1 = 0.9 (Polar Covalent Bond)
2.
H and H
2.1 - 2.1 = 0 (Nonpolar Covalent Bond)
3.
Ca and Cl
3.0 - 1.0 = 2.0 (Ionic Bond)
4.
Al and Cl
3.0 - 1.5 = 1.5 (Polar Covalent Bond)
5.
Mg and O
3.5 - 1.2 = 2.3 (Ionic Bond)
6.
H and F
4.0 - 2.1 = 1.9 (Ionic Bond)
Using electronegativity values, place the following
bonds in order of increasing polarity:
N—N
O—H
Cl—As
Bond
Electronegativity Diff.
N-N
O-H
Cl-As
O-K
3.0 - 3.0 = 0.0
3.5 - 2.1 = 1.4
3.0 - 2.0 = 1.0
3.5 - 0.8 = 2.7
O—K
Type of Bond
Nonpolar Covalent
Polar Covalent
Polar Covalent
Ionic
Least Polar --------------------------> Most Polar (ionic)
Which is more polar N—H or H—F?
N—N
Cl—As
O—H
O—K
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Let’s Practice Drawing Dipoles!
Dipoles
Please identify the polar covalent bonds by
assigning slightly positive and slightly negative
symbols to the atoms below
Dipole: A molecule that has two poles
In looking at the previous example (HCl), we learned that
there is a significant difference in electronegative values.
This means that the chlorine atom acquires a slightly
negative charge (minus sign). The hydrogen atom acquires a
slightly positive charge (plus sign).
Si (1.8) – Br (2.8)
When there is unequal sharing of electrons a dipole exists
Se (2.4) – F (4.0)
It can also be represented by an arrow pointing to the
more electronegative atom.
!+
H
!"
Cl
or
H
N (3.0) – H (2.1)
Cl
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Hydrogen Bonding
Attractions Between Molecules
Intermolecular attractions are weaker than either
ionic or covalent bonds.
Hydrogen Bonds - attractive forces in which
a hydrogen covalently bonded to a very
electronegative atom is also weakly bonded
to an unshared electron pair of another
electronegative atom
These attractions are responsible for determining
whether a molecular compound is a gas, liquid, or
solid at a given temperature.
Van der Waals forces – consists of the two
weakest attractions between molecules
dipole interactions – polar molecules attracted
to one another
dispersion forces – caused by motion of
electrons (weakest of all forces)
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Intermolecular Attractions
Hydrogen Bonding
This other atom may be in the same
molecule or in a nearby molecule, but
always has to include hydrogen
A few solids that consist of molecules do
not melt until the temperature reaches
1000ºC or higher called network solids
(Example: diamond, silicon carbide)
Hydrogen Bonds have about 5% of the
strength of an average covalent bond
Network Solid – solids in which all of the
atoms are covalently bonded to each other
Hydrogen Bond is the strongest of all
intermolecular forces
Melting a network solid would require
breaking covalent bonds throughout the
solid
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