Download Covalent Bonding

COVALENT BONDING
8.1 Molecules & Molecular Compounds
• Molecule: a neutral group of atoms joined by covalent
bonds
• Diatomic Molecule: two atoms joined by a covalent bond
• Examples: H2, Cl2, O2, NO, CO
• Diatomic elements: Dr. Brinclhof
• Molecular Compounds: Compounds composed of
molecules (covalent bonds)
Comparison of Molecular & Ionic
Compounds
Molecular
Ionic
Bonding
Covalent
Ionic
Melting point
Lower
Higher
Electrolyte
Usually weak or
non
Strong
Physical state
@ room temp
(s), (l), (g)
(s)
Molecular Formulas
• Show number & type of atoms in a molecule
• CH4
• H2S
• HNO3
• C6H6
• C3H7OH
• (NH4)3PO4
Structural Formulas
• Show the arrangement of atoms in a molecule
8.2 Nature of Covalent Bonding
• Introduction to Lewis Theory
• Lewis theory
• Octet rule is a guide
• Some exceptions will occur
• Boron accepts less than an octet
• Phosphorus & Sulfur can accept more than an octet
• “expanded octet”
• Electron pairs are shared to form a covalent bond
• In most cases, octets are completed by sharing pairs of electrons
Formation of a Single Covalent Bond
• Formed when two atoms share one pair of
electrons
Why do some elements form diatomic
molecules?
Single Covalent Bonds
The hydrogen and oxygen atoms attain noble-gas configurations by
sharing electrons.
Ammonia, NH3
Drawing Electron Dot (Lewis) Structures
Lewis structure is a type of structural formula that
depicts all the valence electrons in the
molecule or ion
See Tutorial
1. Determine the total # ve
2. Connect atoms in such a way that all have a
noble gas configuration (octet rule)
3. Carbon is often a central atom
4. Check
Draw Lewis Structures for these
Molecular Compounds
• HCl hydrogen chloride
• Cl2
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•
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chlorine
I2
iodine
H2O2 hydrogen peroxide
PCl3 phosphorous trichloride
CH4 methane
Single, Double and Triple
Covalent Bonds
• Sometimes atoms share more than one pair of ve’s
• A bond that involves one shared pair of e-s is a single
covalent bond
• Two shared pairs of electrons is a double covalent bond.
• Three shared pairs of electrons is a triple covalent bond.
Acetylene
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•
•
A gas used in cutting steel
Molecular formula is C2H2
Draw the Lewis structure for acetylene
1.
2.
3.
4.
Connect the atoms
Calculate ve’s
Form single covalent bonds between atoms
Complete octets until remainder of ve’s are
used
Form double or triple bonds if needed to
complete octets.
5.
Polyatomic Ions
• Same process except…
• Add or subtract e-s to account for the charge of the ion, for
•
•
•
•
example
[NH4]+
[SO4]2[ClO][ClO2]-
ammonium ion
sulfate ion
hypochlorite ion
chlorite ion
Coordinate Covalent Bonds
• Bonds in which one of the shared pair comes
completely from one of the bonding atoms
• Carbon Monoxide
Bond Energies
• Energy required to break a chemical bond
• Energy released when a bond is formed
• Is a measure of the strength of the bond
• Large bond energies = strong bonds
Type of bond
Bond Energy
(kJ/mol)
C─C
347
C=C
657
C≡C
908
Resonance
• Resonance occurs when two or more valid Lewis
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•
•
•
structures are possible for a compound or ion
Often occurs with placement of a double bond about a
central atom
Resonance structures are all the valid structures
The actual structure is a hybrid of all the possible
resonance structures
i.e. the bonding present in the particle is a hybrid of those
shown in the resonance structures
Ozone
• Is an allotropic form of oxygen
• Molecular formula is O3
• Is a pollutant (smog)
• Protects earth by absorbing UV radiation
• Draw the resonant Lewis structures for ozone
Nitrogen Dioxide
• Formed by lightning strikes
• Molecular formula NO2
• Also a pollutant in automobile exhaust
• Draw the Lewis structures for NO2
• Why is this an exception to the octet rule?
Exceptions to Octet Rule
• When there is an odd number of ve,
NO2
• Less than an octet:
• Boron
BF3
• More than an octet:
• Phosphorous
PCl5
• Sulfur
SF6
• Unfilled d-shells accept additional electrons, creating an
“expanded” octet
8.3 Bonding Theories
• Molecular orbitals
• When covalent bonds form, atomic orbitals merge to form
molecular orbitals
Sigma and Pi Bonds
• Sigma form when atomic orbitals merge along the axis
between nuclei (internuclear axis)
• Pi bonds result when atomic orbitals merge to surround
the internuclear axis
Sigma Bonds
σ bonds are present in single covalent bonds.
Sigma bond: p-orbital overlap
Pi Bonds
π bonds are present in double and triple
covalent bonds
Sigma and Pi Bonds
C2H2
VSEPR Theory
• Valence Shell Electron Pair Repulsion Theory
• The big idea:
• Because covalent bonds and non-bonding pairs of
electrons are areas of negative charge, they repel one
another
• Covalent bonds and non-bonding electrons are called
“electron domains”
VSEPR Predicts the shape of small
molecules
According to VSEPR theory, the repulsion between electron pairs
causes molecular shapes to adjust so that the valence-electron
pairs stay as far apart as possible.
How to predict the shape of the following molecules:
1.
Draw the Lewis structure
2.
Count the electron domains around the central atom
3.
Determine the domain geometry
4.
Determine the molecular geometry (the way the atoms are
arranged
Methane, CH4
Tetrahedron, bond angles of 109.5°
Ammonia, NH3
Trigonal pyramid, 107°
Why is this not trigonal planar?
Why is the H-N-H bond angle not 109.5 °?
Water, H2O
• Draw the Lewis structure
• Determine the total domains
• Determine the bonding domains
• Determine the shape of the molecule
• Why is water a bent molecule and not a linear one?
Molecular Shape of Water
Water is a bent molecule
Triatomic bent
Hybrid Orbitals
• When covalent bonds form, atomic orbitals mix together
to form hybrid orbitals
• Atomic orbitals involved in bonding often contain a single
unpaired electron
• When the orbitals hybridize, a pair of electrons is shared
• These hybrid orbitals are equal in number to the atomic
orbitals which made them
Covalent Bond formation in CH4
In order for carbon’s 4 ve
to be used in bonding, one
2s2 electron is promoted to
2p.
This results in 4 unpaired
ve, which can then bond
with unpaired e’s of other
atoms.
In order to accomplish this,
the atomic orbitals of C
containing these ve
hybridize.
One s and three p orbitals
hybridize to form four
equivalent orbitals, called
sp3 orbitals
Covalent bonding in CH4
• The s (one) and p (three) orbitals in the valence shell of C
hybridize (merge) to form four equivalent sp3 orbitals.
• They are called sp3 orbitals because they are formed from
one s orbital and three p orbitals
Formation of Hybrid Orbitals
• http://www.mhhe.com/physsci/chemistry/essentialchemistr
y/flash/hybrv18.swf
Hybrid Orbitals
•Hybridization Involving Single Bonds
Hybrid Orbitals
•Hybridization Involving Double Bonds
Hybrid Orbitals
•Hybridization Involving Triple Bonds
How to Determine Hybridization about an
Atom
• The principle: the number of hybrid orbitals must
equal the number of atomic orbitals hybridized
• Count the number of covalent bonds about an
atom
• This must equal the number of hybridized orbitals
• Beginning with s, continue to add orbitals until the
total equals the number of covalent bonds about
the atom
Hybridization Chart
# bonds
Hybridization
2
sp
3
sp2
4
sp3
5
??
6
??
Predicting Hybridization
• What hybridzation would be found about carbon in the
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•
•
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following molecules?
HC≡CH
sp
H2C=CH2
sp2
H3C-CH3
sp3
8.4 Polar Bonds and Molecules
• Electrons in a covalent bond are attracted to
the nuclei of both atoms. Why?
Unequal Sharing of Bonding Electrons
• When covalently bonded to another atom, some atoms
attract electrons more strongly than others
• These atoms have greater “electronegativity”
• When bonded atoms differ in electronegativity, they do
not share the bonding electrons equally
Bonding Electrons in HCl
• Bonding e’s spend
more time near Cl than
H
• What does this imply
about Cl?
• What does this imply
about the distribution
of electrical charge in
HCl?
Polar Covalent Bonds
• When bonded atoms are sufficiently different in
electronegativity, the bond develops negative (-)
and positive (+) ends
• Why? Because the bonding e’s spend more time
around the more electronegative element
• i.e. the bonding e’s are not shared equally
• This unequal distribution of (-) charge is called a
dipole
• The bond is called a polar covalent bond
Polar Bonds and Molecules
• Bond Polarity
• Bond polarity has to do with unequal distribution of shared
electrons caused by differences in electronegativity
between bonded atoms
• This causes one end of the bond to have a “partial
positive” (δ+) charge and the other to have a “partial
negative” (δ-)charge
• These polar covalent bonds and possess a dipole
moment
• The dipole moment is symbolized as -|-------->
Bond Character
• Describes the type of charge distribution in a
chemical bond
• Based upon differences in electronegativity
Differences in Electronegativity and Bond
Character
Polar Molecules
• Molecules containing polar bonds may have a net dipole
• The molecule may have a (+) and (-) side
• Depends upon two factors
• Presence of polar bonds
• Geometry (shape) of molecule
Polarity of Molecules
• A molecule as a whole has a dipole depending upon
• The presence of polar bond(s)
• The geometry of a molecule
• Examples:
• CH4
• CO2
• H 2O
Polar Molecules
Intermolecular Forces
• Types of intermolecular forces account for differences
between ionic and molecular substances.
Intermolecular Forces of Attraction
• Not chemical bonds
• Much weaker than covalent or ionic bonds
• Van der Waals Forces
dipole-dipole interactions
London dispersion forces
• Hydrogen Bonds
very important
Hydrogen Bonds
• Hydrogen bonds
• Attraction between a hydrogen covalently bonded to a very
electronegative atom to an unshared electron pair of another
electronegative atom
• May involve different molecules or occur within very large
molecules like proteins or nucleic acids
http://www.chem.ucla.edu/harding/IGOC/H/hydrogen_bond_acceptor.html
Hydrogen Bonding
• Hydrogen bonding accounts for the unusual properties of
water.
Hydrogen Bonding in Water
http://www.mikeblaber.org/oldwine/BCH4053/Lecture03/Lecture03.htm
Network Solids
https://opentextbc.ca/chemistry/wp-content/uploads/sites/150/2016/05/CNX_Chem_10_05_NtwrkSolid.jpg
Network Solids are Molecular Solids with
High Melting Points
• Molecular solids have lower melting points because only
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weak intermolecular forces have to be broken in order for
them to melt
In network solids, all atoms are covalently bonded to one
another throughout the solid
In order to melt, much stronger covalent bonds must be
be broken (requires more heat)
This is why network solids have very high mp’s (>1000
°C)
In some cases, network solids will decompose rather than
melt