Bonds Part I
Molecular Compound
9.1 Key points
• Describe how a covalent bond forms, including the
energy change involved in the process.
• Use the octet rule to draw Lewis electron dot structures
for simple molecules. Know how and when to
incorporate double and triple bonds into the structures.
• Understand how a coordinate covalent bond differs from
other covalent bonds.
• Be able to draw Lewis structures for polyatomic ions.
• Understand the concept of resonance.
• Know some common exceptions to the octet rule.
• Relate bond energy to the stability and reactivity of
molecules.
LET’S
FIRST
REVIEW
IONIC
BONDING
In an IONIC bond,
electrons are lost or gained,
resulting in the formation of IONS
in ionic compounds.
K
F
The compound potassium fluoride
consists of potassium (K+) ions
and fluoride (F ) ions
K
+
_
F
The ionic bond is the attraction
between the positive K+ ion
and the negative F- ion
Compounds and Molecules
• Compound: a substance that is made from the
atoms of two or more elements that are
chemically bonded.
• Notice: The type of bond is not important, can
be ionic, covalent or metallic
Examples:
H2O, CO2, NaCl, C6H12O6
Non-examples:
I2, O2, Na, Si
Compounds and Molecules
• Molecule: a neutral group of a least two
atoms held together by covalent bonds
– Now the type of bond is important:
• Only covalent bonds
**Notice it only has to be two atoms**
• It can have two or more atoms of the same element
or two or atoms of different elements
Examples:
• H2O, CO2, F2, H2, C6H12O6
Non-Examples:
• NaCl, MgO, Al2O3,
3 Types of Chemical Bonds
1. Ionic Bonds – a metal cation transfers
valence electrons to a nonmetal anion
2. Metallic Bonds – postive cations in a sea of
mobile valence electrons
3. Covalent Bonds – the bonds we will study in
this chapter
All three types of chemical bonds are intramolecular
forces : the forces between atoms within a compound
Covalent Bonds
• Covalent Bonds – “Co-Workers”
Nonmetal + Nonmetal
• two atoms share valence electrons to form a
stable octet
• Examples: H2O, CO2, NO2, SF6
– Covalently bonded compounds are called
molecules
Covalent Bonds
• Molecular Formula: shows how many atoms
of each element a molecule contains.
– Examples:
• Diatomic Elements - O2, H2, Cl2
• Molecules - CH4, NH3, H2O
Oxygen
molecule
O2
Benzene
C6H6
Molecular Formulas
• The formula for water is written as H2O
What do the subscripts tell us?
• Molecular formulas do not tell any information
about the…..
structure!
(the arrangement of the various atoms).
Covalent bonds
• Why do nonmetals share electrons?
– Remember Nonmetals
• Hold on to their valence electrons
• Cannot give away electrons to bond.
• Still want to form a stable octet.
• By sharing valence electrons both nonmetal
atoms get to count the electrons toward a
stable octet.
So
what
are
covalent
bonds?
In covalent bonding,
atoms still want to
achieve
a noble gas
configuration
(the octet rule).
In covalent bonding,
atoms still want to achieve
a noble gas configuration
(the octet rule).
But rather than losing or gaining
electrons,
atoms now share an electron pair.
Showing Covalent bonding
• Show the bonding of Cl2
Chlorine
forms
a
covalent
bond
with
itself
Cl2
Cl
Cl
How
will
two
chlorine
atoms
react?
Cl Cl
The octet is achieved by
each atom sharing the
electron pair in the middle
O2
Oxygen is also one of the diatomic molecules
O
O
How will two oxygen atoms bond?
O O O= O
For convenience, the double bond
can be shown as two dashes.
Important Covalent Compounds
• 7 Diatomic Elements *Memorize*
O2
N2
F2
Cl2
Br2
I2
H2
These elements are NEVER found as
individual atoms.
Ex: The oxygen gas we breathe is O2
Types of Chemical Bonds
• Polar Vs Nonpolar
– Nonmetals do not always equally share their
electrons
– Some nonmetals can have a stronger pull on the
shared pair of electrons—like tug of war of e-
– These 2 types of covalent bonds are called
polar and nonpolar.
• Covalent Bonds: Polar and Nonpolar
– Polar: a covalent bond in which the bonded
atoms have an unequal attraction for the shared
pair of electrons
– Nonpolar: a covalent bond in which the two
bonding electrons are shared equally by the
bonded atoms.
• Electronegativity: How bad an element
wants an electron
– Using electronegativity differences to predict
polarity and the bond type
• Electronegativity Difference: (in Packet p. 13)
– 0.0 - 0.4 = Nonpolar Covalent
– 0.4 – 1.7 = Polar Covalent
– > 1.7 = Ionic
Electronegativity
*0.0 - 0.4 = Nonpolar * 0.4 – 1.7 = Polar * > 1.7 = Ionic
• Partial negative: element is partially neg.
• Partial positive: element is partially pos.
• Examples: Determine the electronegativity
difference, the bond type and indicate partial
positive and partial negative charges.
a.) H and I H= ___ I=___ , Δ = ____
Bond type=_______________
H-I
b.) K and Br K=____ Br=_____, Δ = _____
Bond type=_______________
K - Br
Ex: Draw the electron dot diagram for the
covalent bonds
**Remember Hydrogen needs only 2
electrons to fill the outer shell.
F2
CH4
Bonds
• 2 valence electrons = 1 bond
• Hydrogen can only form one single bond
WHY??
Single Bond
• Single bond: when atoms share 1 pair of
electrons (2 electrons total)
Draw lewis dot for H2O, then show bonds
~Tips for writing lewis dot structures for
molecules with more than 2 atoms:
• Central atom: is the 1st element in the
compound or molecule (except H)
1. **The central atom ALWAYS goes in the
middle!!! ***
2. Rearrange dots so that every element has 8
valence electrons (H and He only need 2 val)
Structural Formulas
• structural formula: Showing bonds.
HOH
Double Bond
**Two atoms can share more than one pair of
valence electrons.
• Double bond: when atoms share 2 pairs of
electrons (4 electrons total)
Ex 1: Draw the lewis dot for CO2, then show
structural formula
Double bonD cont…
Ex 2: Draw the lewis dot for H2CO, then show
structural formula.
Triple Bond
~ Triple bond: when atoms share 3 pairs of
electrons (6 electrons total)
Draw the lewis dot for HCN and show structural
formula.
How to find the # of bonds in a
lewis structure
1. Find the total # of valence electrons.
2. Use the formula to find the number of bonds.
# of val e- needed (all have 8 or 2 e-)
- # of val e- available
= ____/2 to find the # of bonds
1. Find the total # of valence electrons.
2. Use the formula to find the number of bonds.
# of val e- needed (all have 8 or 2 e-)
- # of val e- available
= ____/2 to find the # of bonds
Ex: Find the number of bonds for each molecule or
compound and write the lewis dot and structural formula:
a.) CO
b.) C2F4
c.) C2H6
Exceptions to Octet rule
• For some molecules, it is impossible to
satisfy the octet rule
• Yet the stable molecules do exist
• Two types of exceptions:
– Atoms that cannot hold 8 valence electrons
• Hydrogen, helium, beryllium, boron, aluminum
– Atoms that can hold more than 8 valence
electrons
• Phosphorus, sulfur, iodine, xenon, krypton
Exceptions to the Octet Rule
1. Most covalent compounds of Beryllium: the number
of valence electrons needed for Be is 4.
•BeF2
2. Most covalent compounds of Group 13: Primarily
Boron & Aluminum - the number of valence electrons
needed is 6
•AlF3
•BF3
Exceptions to Octet rule
3. Sometime when Phosphorus, Sulfur, Iodine,
Xenon & Krypton are the central element they
can hold more than 8 electrons:
•PCl5
•I3
•SF6
I–I–I
Review on charges on bonding:
• Ionic Bonds:
– Have a full positive or full negative charge.
– Ionic bonds do NOT have partial charges.
Why?
• Polar Covalent Bonds:
– Have partial positive or partial negative charges.
Why?
• Nonpolar Covalent Bonds:
– Have NO partial positive or partial neg. charge.
Why?
Intermolecular Forces (IMF)
• Attractive forces between molecules.
• Much weaker than chemical bonds.
Intramolecular forces
are within a molecule. (bonds)
Types of IMF
• London Dispersion Forces:
– Occurs between nonpolar molecules (diatomics)
– Caused by motion of electrons ( “e- sloshing” ),
they create a temporary dipole (slight charge)
– Weakest of all forces.
View animation online.
Types of IMF
• Dipole-Dipole Forces:
– Occurs between polar molecules
– Where one side is partial positive and one is
partial negative.
– Stronger than London Dispersion forces.
-
+
View animation online.
Types of IMF
• Hydrogen Bonding:
– When Hydrogen bonds to Nitrogen, Oxygen or
Fluorine (NOF)
– Strongest of all intermolecular forces!
Types of intermolecular forces:
Examples of intermolecular forces:
Classify as London, Dipole or Hbonding.
• NCl (nonpolar)
• CO (polar)
• HF (polar)
Properties Molecular Compounds
• Low melting points and boiling points.
– The IMF between molecular compounds are
weaker than ionic or metallic compounds
– This means that only a small amount of energy is
required break the bonds
Strongest Bonds  Weakest Bonds
Heat and electrical conductors
• Covalent bonds: poor electrical and
thermal conductivity.
– No mobile electrons to conduct current
Review of bonds:
Covalent:
Ionic:
Metallic:
Draw Lewis dot diagrams for
polyatomic ions: p.6 in packet
1. SO42-
2. PO43-
Molecular Geometry
Lewis structures fail to indicate threedimensional shapes of molecules.
The shape of a molecule controls some of its
chemical and physical properties.
VSEPR
Valence Shell Electron Pair Repulsion Theory - predicts
the shapes of a number of molecules and polyatomic
ions.
•Electron pairs move to create the most stable
arrangement.
-The repulsions between electron pairs causes
molecular shapes to adjust so that the
electron pairs stay as FAR APART as
possible.
What are the ideal arrangements of electron
pairs to minimize repulsions?
1) We need to identify the number of regions of high electron
density, called the steric number, on the central atom.
2) Regions of high electron density include:
Single bonds
Double bonds
Triple bonds
Unshared (lone) pairs of electrons
**Double and triple**
bonds only count as
ONE region of high
electron density just
like a single bond or a
lone pair.
Examples: Draw the Lewis Dot
Structure and fill in the following:
1. CH4
– Steric # ____
– # of lone pairs _____
2. H2O
– Steric # ____
– # of lone pairs _____
3. CO2
– Steric # ____
– # of lone pairs _____
1. CH4
Examples: Use table to determine
molecular shape and bond angle.
– Steric # 4
Molecular Shape: __________
– # of lone pairs 0
Bond angle: _________
2. H2O
– Steric # 4
Molecular Shape:_____________
– # of lone pairs 2 Bond angle:________________
3. CO2
– Steric # 2
Molecular Shape:______________
– # of lone pairs 0 Bond angle: ______________
How does Molecular Geometry affect Polarity?
1. One polar bond on central atom
Molecule polar?
Molecule nonpolar?
2. More than one polar bond on the central atom will cancel
out polarities if they have equal electronegativities.
Molecule polar?
Molecule nonpolar?
How does Molecular Geometry affect Polarity cont..
3. One lone pair on the central atomPolar? Nonpolar?
4. Two or more lone pairs on the central atom
Polar? Nonpolar?
Water
(asymmetrical)
Xenon difluoride
(symmetrical)
Xenon tetrafluoride
(symmetrical)
Two regions of high electron density
•AX2 notation
•Steric # is 2
•No lone pairs
•Geometry is linear
•Bond Angle is 180
Look at the example of the BeF2(g) molecule.
The Lewis Structure is:
Example: BeH2
• Steric # _____
H : Be :
• # of lone pairs
• Bond angle _________
• Molecular Geometry __________
H
Example: CO2
• Steric # _____
• # of lone pairs ____
• Bond angle _________
• Molecular Geometry __________
• Is the molecule polar?
•Electronegativity Difference between
Carbon & Oxygen is .89
•So the bonds are polar
•But is the molecule?
Example: CO2
• Is the molecule polar? WHY?
•
Example: HCN
• Steric # _____
• # of lone pairs
• Bond angle _________
• Molecular Geometry __________
• Is the molecule polar? WHY?
Three regions of high electron density
•AX3 notation
•Steric # is 3
•No lone pairs
•Geometry is
trigonal planar
•Bond Angle is
120
Example of BF3 molecules.
The Lewis Structure is:
Example: BF3
• Steric # _____
• # of lone pairs ______
• Bond angle _________
• Molecular Geometry __________
• Is the molecule polar? _______
•AX2E noation
•Steric # is 3
•# of lone pairs is 1
•Geometry is bent
•Bond angle is 120
Example is GeF2
• Steric # _____
• # of lone pairs ______
• Bond angle _________
• Molecular Geometry __________
Is this molecule polar? ____
Four regions of high electron density
•AX4 notation
•Steric number is 4
•No lone pairs
•Geometry is tetrahedral
•Bond angle is 109.5
Look at the example of CH4
molecules.
The Lewis Structure is:
• Steric # _____
• # of lone pairs ______
• Bond angle _________
• Molecular Geometry __________
Is the molecule POLAR? _________
•AX3E notation
•Steric # is 4
•#of lone pairs is 1
•Geometry is
trigonal pyramidal
•Bond angle is 107
Example NH3
The Lewis structure is:
NH3
•Steric # _____
• # of lone pairs _____
• Bond angle _________
• Molecular Geometry __________
Is the molecule POLAR? _________
•AX2E2 notation
•Steric # is 4
•#of lone pairs is 2
•Geometry is bent
•Bond angle is 105
Example H2O.
The Lewis structure is:
H2O
•Steric # _____
• # of lone pairs _____
• Bond angle _________
• Molecular Geometry __________
Is the molecule POLAR? _________
FIVE regions of high electron density
•AX5 notation
•Steric Number 5
•No lone pairs
•Geometry is trigonal
bipyramidal
•Bond angle is 90/120
Example of PF5 molecules.
PF5
•Steric # _____
• # of lone pairs _____
• Bond angle _________
• Molecular Geometry __________
Is the molecule POLAR? _________
SIX regions of high electron density
•AX6 notation
•Steric # is 6
•No lone pairs
•Geometry is octahedral
•Bond angle is 90
Example SF6 molecules.
SF6
•Steric # _____
• # of lone pairs ______
• Bond angle _________
• Molecular Geometry __________
Is the molecule POLAR? _________
London Dispersion
Dipole Dipole
Hydrogen Bonding
Octet Rule
Electronegativity
Polar
Nonpolar
Sharing
Transfer
Gaining
Molecular Formula
Formula Unit
Lone Pair
Chemical Bonds
Double Bond
Molecule
Intramolecular forces
Between
Sea of electrons
Cation