Download 7.1 The Covalent Bond 7.2 Strengths of Covalent Bonds

Dr. Bottaro Chem 1010 Fall 2005
Chapter 7: Covalent Bonds and
Molecular Structure (7.1-7.7, 7.9, 7.11, 7.12)
Chapter Goals: Be Able to:
Predict which compounds are ionic and which are molecular.
Use the periodic table to predict which of two elements is more
electronegative, thus if a given bond is ionic, polar covalent, or
nonpolar covalent.
Write Lewis structures (electron-dot structures) for atoms and
molecules
For each atom in an electron-dot structure give the number of bonded
electron pairs and the number of nonbonded electron pairs.
For a given Electron-dot structure, give the number of single bonds,
double bonds, and triple bonds (bond order).
Draw Electron-dot resonance structures as needed.
Use the VSEPR model to predict geometries of molecules and
polyatomic ions, including those with more than one central atom.
Sketch and identify the orbitals used by each atom to form bonds in
molecules and polyatomic ions.
7.1 The Covalent Bond
covalent bond - formation of covalent, sharedelectron bond between atoms
Bond length - the
optimum distance
between nuclei - net
attractive forces are
maximized (minimum
energy) and the H-H
molecule is most stable
7.2 Strengths of Covalent Bonds
Bond dissociation energy
- the amount of energy that must be supplied to break
a chemical bond in an isolated molecule in the
gaseous state
- are always +ve (energy must be supplied to break a
bond)
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Dr. Bottaro Chem 1010 Fall 2005
7.3 Comparison of Ionic and
Covalent Compounds
ionic compounds (e.g. NaCl)
high-melting point solids
3-dimensional networks of ions
requires large amount of energy to
overcome attractions
covalent compounds (e.g. HCl)
low-melting point solids, liquids or gases
discrete molecules
only small amounts of energy required to
overcome attractive forces between
molecules
bonds between atoms are often very strong
(delta)
a) + (partial positive) - atom that has a smaller share of
bonding electrons
b) - (partial negative) - atom that has a larger share of the
bonding electrons (i.e. the electrons
are on average closer to this
nucleus.)
7.4 Polar Covalent Bonds:
Electronegativity
ionic bonds - electrons completely transferred
covalent bonds - electrons shared equally
polar covalent bonds
- bonding electrons are shared unequally between
two atoms but not completely transferred
Cl : Cl
Na+ Cla nonpolar
an ionic bond
covalent
bond
+H-Cl [H :Cl]
a polar covalent bond
(bonding e are attracted
more strongly by Cl than H)
Electronegativity (EN)
the ability of an atom in a molecule to
attract the shared e- in a covalent bond
causes bond polarity
Generally:
metallic elements have a low electronegativity,
alkali metals are the least electronegative,
reactive nonmetals have a high electronegativity,
fluorine, oxygen, nitrogen and chlorine are most
electronegative.
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Dr. Bottaro Chem 1010 Fall 2005
Predicting Bond Polarity
nonpolar covalent bonds
- formed between atoms with the same or similar
electronegativities
ionic bonds
- formed between atoms whose electronegativities
differ by more than about 2 units
polar covalent bonds
- formed between atoms whose electronegativities
differ by less than 2 units
Practice Problems:
Predict which bond is more polar and assign partial
charges where appropriate.
a) C H or C Br
c) N Cl or N Mg
b) Si Li or Si Cl
Representations in Lewis Structures:
1. element s symbol represents atomic nucleus with its
core electrons
2. valence electrons (represented by dots) are placed
around symbol one at a time until they are used up or all
four sides have at least one electron
3. the remaining valence electrons are paired with ones
already there.
e.g.
H
C
N
O
H
H C H
H
H C H
H
H
Cl
7.6 Electron-Dot Structures
(Lewis Structures)
Electron-dot structures (Lewis Structures)
a way of keeping track of valence electrons for maingroup elements
Noble gases
have completely filled outer-most shell very stable
configuration (relatively unreactive)
octet rule - have eight valence electrons (except He)
valence electrons
those electrons that participate in bonding to other atoms
core electrons
electrons from inner-shells which do not participate in
bonding
include the inner-shell e- which have a noble-gas
configuration
Valence electrons in molecules and ions ordinarily
occur in pairs:
Two kinds of electron pairs:
1. bonding electrons
- pair of e- shared btw two atoms (shown as straight line
between bonded atoms)
- form a covalent bond
2. lone pair (or nonbonding pair) electrons
- unshared pair of electrons (shown as pair of dots on the
atom)
- belongs entirely to one atom
Single and multiple covalent bonds:
1. single bond - bonded atoms sharing one e- pair
2. double bond - bonded atoms sharing two e- pairs
3. triple bond - bonded atoms sharing three e- pairs
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Dr. Bottaro Chem 1010 Fall 2005
Steps for writing Lewis Structures:
1. Consider how the atoms are attached together:
central atom
- often the one with the lowest electronegativity (except H)
- usually the one written first in the formula
terminal atoms
- often are H, O or a halogen (Cl, Br, or I)
2. Determine the total number of valence electrons.
3. Form single bonds between each atom pair.
4. Distribute remaining electron pairs as lone pairs so that they
obey the octet rule.
5. Share lone pairs between atoms to make multiple bonds
where necessary to obey the octet rule.
Drawing Lewis Structures (Electron Dot Structures)
NH2F Amino Fluoride: In this molecule, nitrogen is the
central atom.
Rule 1: Number of electrons = 5 + (2 x 1) + 7 = 14 = 7 pairs
H N H
H N H
H N H
F
F
F
Rule 2
Rule 3
Compounds that Do Not Obey the
Octet Rule
electron-deficient molecules
- molecules in which some elements may have less than an
octet
- most commonly Al, B (Group 3A) and Be
expanded valence shell
- elements of 3rd or higher period (row) may be surrounded
by more than 4 valence pairs in certain compounds
- use empty outer d-orbitals and available lone pairs to
allow for an additional bonding.
e.g. BrF3
Bond order
- number of electron pairs being shared between any two
bonded atoms
single bond - bond order of 1 (1 shared pair of electrons)
double bond - bond order of 2 (2 shared pairs)
triple bond - bond order of 3 (3 shared pairs)
Rule 4
Practice Problems.
Draw the electron-dot structures for the following molecules
or ions: CCl4, CS2, NBr3, BF4-, O22-, SF4, SF6 , XeOF4,
XeF5+, XeF4
What are the bond orders for the bonds in these compounds?
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Dr. Bottaro Chem 1010 Fall 2005
Types of covalent bonds:
covalent bond
- formed when two atoms each contribute one electron
The electron-dot structures provide a simple way of
representing chemical reactions.
Ionic:
Covalent:
coordinate covalent bond
- formed when one atom donates both electrons (a lone pair)
to another atom that has a vacant valence orbital
7.6 Electron Dot Structures of
Polyatomic Ions
Most biological molecules contain only H, and a
combination of C, N and O.
These molecules can be complex, containing many atoms.
However, they are easy to draw because they almost
always follow the octet rule.
7.7 Electron-Dot Structures and
Resonance:
resonance structures
- sometimes it is possible to write more than one Lewis
Structure to represent a molecules
- some molecules with multiple bonds adjacent to single
bonds
e.g. O3
Practice Problems:
Draw electron-dot structures for:
C3H8
H2O2
CO2
N2H4
CH5N
C2H4
C2H2
Cl2CO
resonance hybrid
- actual electronic structure which is an average of
the different possible resonance structures
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Dr. Bottaro Chem 1010 Fall 2005
Notes:
1. resonance forms do not imply different kinds of molecules
(structure is intermediate btw those of the resonance
forms written.
2. resonance can be anticipated when it is possible to write
two or more Lewis structures that are about equally
plausible.
3. resonance forms differ only in distribution of e-, not in
arrangement of atoms.
Practice problem:
Draw as many electron-dot resonance structures as possible
for: SO2, CO32 , SO42 , PO43 .
Effect of Unshared Pairs of Molecular
Geometry
electron-pair geometry
- geometry describing the arrangement of all electron pairs
(bonding and nonbonding) around the central atom
molecular geometry
- geometry describing the arrangement of bonded atoms
around the central atom
- the positions of unshared pairs of electrons are not
included
7.9 Molecular Shapes:
The VSEPR Model
VSEPR - Valence-Shell Electron-Pair Repulsion Theory
- all molecules have a 3-dimensional shape
- VSEPR used to predict approximate shape of a molecule
- shape plays a crucial role in determining the molecule s
chemistry
**According to VESPR theory, valence e- pairs (either
bonding or nonbonding) surrounding an atom repel one
another.
Orbitals containing valence electrons orient themselves
around the central atom so that they are as far away from
one another as possible
Non-equivalence of electron pairs:
1. bonding pairs (bp)
- localized to a region around the bond axis due to their
attraction to both atoms
2. lone pairs (lp)
- more diffuse and in a sense occupy more space
Note: for a central atom with a mixture of lp and bp
interactions (repulsions) are not all equal and the
geometry of the molecule will be affected
i.e. Groups do not compete equally for space:
Lone Pair > Triple Bond > Double Bond > Single Bond
pair requiring
pair requiring
most space
least space
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Dr. Bottaro Chem 1010 Fall 2005
Determining Geometries by VSEPR
1. Count the total electron groups from the electron dot
structure.
2. Arrange electron groups to maximize separation.
Molecular geometry of diatomic molecules :
since there are only two points to a straight line, the
molecule must be linear, bond angles: 180o
e.g. Cl2 and HCl
Cl Cl
H Cl
Geometries for Molecules with Two
Charge Clouds
- Central atom has two electron clouds
e.g. BeF2
Central atom Be Two electron clouds
electron-pair geometry: linear
F-Be-F
molecular geometry: linear
bond angles: 180o
- also include compounds such as CO2 and HCN
Ideal Geometries with Two to Six Charge Clouds on the
Central Atom (3 or more atoms in a molecule)
- geometry is not obvious from molecular (or ionic) formula
- must consider bonding and non-bonding electron pairs
- must also consider bond angles (angles between bonds)
Geometries for Molecules with Three
Charge Clouds
Three electron groups lie in the same plane and point to
the corners of an equilateral triangle.
e.g. BF3 and H2CO and SO2
BF3 : electron-pair geometry: trigonal planar
molecular geometry: trigonal planar
bond angles: 120o
H2CO :
electron-pair geometry: trigonal planar
molecular geometry: trigonal planar
bond angles: H-C-O > 120o
H-C-H < 120o
SO2
electron-pair geometry: trigonal planar
molecular geometry: bent
bond angles: O-S-O < 180o
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Dr. Bottaro Chem 1010 Fall 2005
Geometries for Molecules with
Four Charge Clouds
- includes species that follow the octet rule
- charge clouds are directed towards the four corners
of a tetrahedron
e.g. CH4 electron-pair geometry: tetrahedral
molecular geometry: tetrahedral
bond angles: 109.5o
NH3
electron-pair geometry: tetrahedral
molecular geometry: trigonal pyramidal
bond angles: < 109.5o
H2O
electron-pair geometry: tetrahedral
molecular geometry: bent
bond angles: < 109.5o
Geometries for Molecules with Five
Charge Clouds
- includes species display an expanded valence
- charge clouds are directed towards the corners of a
geometric figure known as a trigonal pyramid
Examples:
PCl5 electron-pair geometry: trigonal bipyramidal
molecular geometry: trigonal bipyramidal
bond angles: equatorial 120o
axial 90o
SF4
electron-pair geometry: trigonal bipyramidal
molecular geometry: see-saw
bond angles: equatorial <120o
axial <90o
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Dr. Bottaro Chem 1010 Fall 2005
ClF3
electron-pair geometry: trigonal bipyramidal
molecular geometry: T-shaped
bond angles: <90o
Geometries for Molecules with Six
Charge Clouds
- includes species display an expanded valence
- charge clouds are directed towards the corners of a
geometric figure known as a octahedron
I3
electron-pair geometry: trigonal bipyramidal
molecular geometry: linear
bond angles: 180o
Examples:
SF6 electron-pair geometry: octahedral
molecular geometry: octahedral
bond angles: 90o
SbCl52electron-pair geometry: trigonal bipyramid
molecular geometry: square pyramidal
bond angles: equatorial 90o
axial <90o
XeF4 electron-pair geometry: octahedral
molecular geometry: square planar
bond angles: 90o
Note: Summary of the molecules discussed is
available in your text in Table 7.4 on pages
268-269
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Dr. Bottaro Chem 1010 Fall 2005
7.11 & 7.12 Hybridization of Orbitals:
sp, sp2, sp3, sp3d, sp3d2
Electron Groups
Geometry
Examples
2
Lone Pairs Bonds
0
2
Linear
BeCl2
3
0
3
Trigonal planar
BF3
3
1
2
Bent
SO2
4
0
4
Tetrahedral
CH4
4
1
3
Trigonal pyramidal
NH3
4
2
2
Bent
H2O
5
0
5
Trigonal bipyramidal
PCl5
5
1
4
See-saw
SF4
5
2
3
T-Shaped
ClF3
5
3
2
linear
I3-
6
0
6
Octahedral
SF6
6
1
5
Square pyramidal
SbCl52-
6
2
4
Square planar
XeF4
hybrid orbital - new kind of atomic orbital formed
Note: - when considering hybridization, the number of
hybrid orbitals formed is equal to the number of atomic
orbitals mixed
- the energies of the hybrid orbitals are in between those
of the atomic orbitals from which they are derived
e.g. BeF2
- as fluorine atoms approach, the atomic orbitals of Be
atom undergo significant change
- 2s orbital mixed (hybridized) with 2p orbital to form two
new sp hybrid orbitals
one s atomic orbital + one p atomic orbital = two sp hybrid
orbitals
Wave functions from s orbitals & p orbitals can be combined
to form hybrid atomic orbitals.
sp3 hybridization
e.g. CH4
Carbon undergoes
orbital hybridization
to allow the formation
of four bonds.
sp hybrid orbitals
sp and p-orbitals
Example: Acetylene, H-C C-H
bond - head-on overlap of sp hybridized orbitals
two bonds - side by side overlap of p orbitals
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Dr. Bottaro Chem 1010 Fall 2005
Types of Bonding
sigma bond ( ) - a covalent bond in which the shared
electrons are centered about the axis between the two
nuclei
sp2 hybrid orbitals (
bond)
e.g. BF3
one s atomic orbital + two p atomic orbital
= three sp2 hybrid orbitals
pi bond ( ) - a bond in which shared electrons occupy a
region above and below a line connecting the two nuclei
e.g. H2C=O
sp3 hybrid orbitals
e.g. Methane, CH4
e.g. CH4, NH3 and H2O
one s atomic orbital + three p atomic orbital
= four sp3 hybrid orbitals
Generally, unshared as well as shared e- pairs
can be located in hybrid orbitals
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Dr. Bottaro Chem 1010 Fall 2005
sp3d hybrid orbitals
sp3d2 hybrid orbitals
eg. PCl5
one s atomic orbital + three p atomic orbital
+ one d atomic orbital
= five sp3d hybrid orbitals
e.g. PCl5
one s atomic orbital + three p atomic orbital
+ two d atomic orbital
= six sp3d2 hybrid orbitals
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