Electron-Dot Structures

Chapter 7
Covalent Bonds and Molecular Structure
國防醫學院 生化學科
王明芳老師
2011-10-18; 2011-10-25
Chapter 7/1
Covalent Bonding in Molecules
Covalent Bond: A bond that results from the sharing of electrons
between atoms.
Covalent bonds are formed by sharing at least one pair of electrons.
A covalent H-H bond. The bond is the net result of
attractive and Repulsive electrostatic forces.
Chapter 7/2
Covalent Bonding in Molecules
Every covalent bond
has a characteristic
length that leads to
maximum stability.
This is the bond length.
A graph of potential energy versus internuclear
distance for the H2 molecule
Chapter 7/3
Strengths of Covalent Bonds
Energy required to break a covalent bond in an isolated
gaseous molecule is called the bond dissociation energy.
Chapter 7/4
Strengths of Covalent Bonds
Chapter 7/5
A Comparison of Ionic and
Covalent Bonds
Chapter 7/6
A Comparison of Ionic and
Covalent Bonds
Sodium chloride, an ionic
compound, is
a white, crystalline solid that melts
at 801℃.
Hydrogen chloride, a molecular
compound, is a gas at room
temperature.
Polar Covalent Bonds:
Electronegativity
Electronegativity: The ability of an atom in a molecule
to attract the shared electrons in a covalent bond.
The bonding continuum from nonpolar covalent to ionic.
[The symbol d means partial charge, either partial positive (d+)
or partial negative (d-)]
Chapter 7/8
Polar Covalent Bonds:
Electronegativity
Chapter 7/9
Polar Covalent Bonds:
Electronegativity
Chapter 7/10
Polar Covalent Bonds:
Electronegativity
Chapter 7/11
Polar Covalent Bonds:
Electronegativity
1. Bond polarity is due to electronegativity differences
between atoms.
2. Pauling Electronegativity: is expressed on a scale
where F = 4.0
3. Metals have low electronegativity; nonmetals higher
electronegativity.
4. % Ionic Character: As a general rule for two atoms in
a bond, we can calculate an electronegativity
difference (∆EN ).
∆EN = EN(Y) – EN(X) for X–Y bond.
5.
If ∆EN < 0.5 the bond is covalent.
If ∆EN > 2.0 the bond is ionic
If ∆EN < 2.0 the bond is polar covalent. Chapter 7/12
Polar Covalent Bonds:
Electronegativity
Electronegativity values and trends in the periodic table
unitless
Chapter 7/13
Electron-Dot Structures
Electron-Dot Structure (Lewis Structure): Represents an
atom’s valence electrons by dots and indicates by the
placement of the dots the way the valence electrons are
distributed in a molecule.
Chapter 7/14
Electron-Dot Structures
1. Using electron-dot (Lewis) structures, the valence
electrons in an element are represented by dots.
2. Valence electrons are those electrons with the
highest principal quantum number (n), F (7 valence
electrons), F2 (8 valence electrons).
3. The electron-dot structures provide a simple, but
useful, way of representing chemical reactions.
Chapter 7/15
Electron-Dot Structures
Single Bonds:
Ionic:
H
H
H
H
H
C
H
C
H
H
Double Bonds:
H
H
Covalent:
C
C
C
C
H
H
Triple Bonds:
C
C
H
C
C
Chapter 7/16
H
Electron-Dot Structures
Chapter 7/17
Electron-Dot Structures
Chapter 7/18
Electron-Dot Structures
Chapter 7/19
Electron-Dot Structures
Chapter 7/20
Electron-Dot Structures
Coordinate-covalent bond
1. This is an especial polar covalent bond, which is
represented by an arrow or double bond.
2. A Coordinate-covalent bond is one in which both
electrons come from the same atom; i.e., the bond can
be regarded as being formed by the overlap of an orbital
containing two electrons with an empty one.
coordinate-covalent bond
O
Me
S Cl
Me N
O polar covalent bond Me
O
Chapter 7/21
Example 7.1 Drawing an Electron-Dot Structure
Draw an electron-dot structure for phosphine, PH3.
Solution
1.
2.
The number of covalent bonds formed by a main-group element
depends on the element’s group number.
Phosphorus, a group 5A element, has five valence electrons and can
achieve a valence-shell octet by forming three bonds and leaving one
lone pair. Each hydrogen supplies one electron.
Electron-Dot Structures of
Polyatomic Molecules
Compounds Containing Only Hydrogen and Second-Row Elements
Only one possible structure
Chapter 7/23
Electron-Dot Structures of
Polyatomic Molecules
Chapter 7/24
Example 7.2 Drawing an Electron-Dot Structure
Draw an electron-dot structure for hydrazine, N2H4.
Solution
1. Nitrogen, a group 5A element, has five valence electrons
and forms three bonds.
2. Join the two nitrogen atoms, and add two hydrogen atoms
to each.
Example 7.3 Drawing an Electron-Dot Structure
Draw an electron-dot structure for carbon dioxide, CO2.
Solution
The only possible structure contains two carbon–oxygen
double bonds.
Example 7.4 Drawing an Electron-Dot Structure
Draw an electron-dot structure for the deadly gas hydrogen cyanide, HCN.
Solution
The only way the carbon can form four bonds and the nitrogen
can form three bonds is if there is a carbon–nitrogen triple
bond.
Example 7.5 Identifying Multiple Bonds in Molecules
The following structure is a representation of histidine, an amino acid constituent of proteins. Only the connections
between atoms are shown; multiple bonds are not indicated. Give the chemical formula of histidine, and complete
the structure by showing where the multiple bonds and lone pairs are located (red=O, gray=C, blue=N, ivory=H).
1.
2.
Count the atoms of each element to
find the formula. Then look at each
atom in the structure to find what is
needed for completion.
Each carbon (gray) should have
four bonds, each oxygen (red)
should have two bonds and two
lone pairs, and each nitrogen (blue)
should have three bonds and one
lone pair.
Example 7.5 Identifying Multiple Bonds in Molecules
Continued
Solution
Histidine has the formula C6H9N3O2.
Electron-Dot Structures of
Polyatomic Molecules
Compounds Containing Elements below the Second Row
The octet rule sometimes fails
Chapter 7/30
Electron-Dot Structures of
Polyatomic Molecules
The octet rule sometimes fails
The octet rule occasionally fails for the main-group elements
shown in blue
Chapter 7/31
Electron-Dot Structures of
Polyatomic Molecules
Step 1: Valence Electrons
• Find the total number of valence electrons for
all atoms in the molecule.
• Add one additional electron for each negative
charge in an anion or subtract one for each
positive charge in a cation.
Chapter 7/32
Electron-Dot Structures of
Polyatomic Molecules
Step 2: Connect Atoms
• Draw lines to represent bonds between atoms.
• Hydrogen and halogens usually form only one
bond.
• Elements in the second row usually form a
certain number of bonds based upon the column
they occupy.
Chapter 7/33
Electron-Dot Structures of
Polyatomic Molecules
Step 2: Connect Atoms
• Draw lines to represent bonds between atoms.
• Hydrogen and halogens usually form only one
bond.
• Elements in the second row usually form the
number of bonds given in the next table.
• Elements in third row and lower are often a
central atom around which other atoms are
grouped and form more bonds than predicted by
the octet rule.
Chapter 7/34
Electron-Dot Structures of
Polyatomic Molecules
Chapter 7/35
Electron-Dot Structures of
Polyatomic Molecules
Step 3: Assign Electrons to the Terminal Atoms
• Subtract the number of electrons used for
bonding from the total number calculated in
step 1 to find the number that remain.
• Complete each terminal atom’s octet (except for
hydrogen).
Chapter 7/36
Electron-Dot Structures of
Polyatomic Molecules
Step 4: Assign Electrons to the Central Atom
• If unassigned electrons remain after step 3,
place them on the central atom.
Chapter 7/37
Electron-Dot Structures of
Polyatomic Molecules
Step 5: Multiple Bonds
• If no unassigned electrons remain after step 4 but
the central atom does not yet have an octet, use
one or more lone pairs of electrons from a
neighboring atom to form a multiple bond (either a
double or a triple).
Chapter 7/38
Electron-Dot Structures of
Polyatomic Molecules
Draw an electron-dot structure for H2O
Step 1: 2(1) + 6 = 8 valence electrons
H
H
Step 2:
O
H
Step 4:
O
H
bonding pair of electrons
lone pair of electrons
Chapter 7/39
Electron-Dot Structures of
Polyatomic Molecules
Draw an electron-dot structure for CCl4
Step 1: 4 + 4(7) = 32 valence electrons
Cl
Step 2: Cl
C
Cl
Cl
Cl
Step 3: Cl
C
Cl
Cl
Chapter 7/40
Electron-Dot Structures of
Polyatomic Molecules
Draw an electron-dot structure for H3O1+ (hydronium ion)
Step 1: 3(1) + 6 – 1 = 8 valence electrons
H
Step 2: H
O
+
H
H
Step 4:
H
O
H
Chapter 7/41
Electron-Dot Structures of
Polyatomic Molecules
Draw an electron-dot structure for CH2O
Step 1: 4 + 2(1) + 6 = 12 valence electrons
O
Step 2: H
C
O
H
Step 5:
H
C
H
O
O
Step 3: H
C
H
H
C
H
formaldehyde
Chapter 7/42
Electron-Dot Structures of
Polyatomic Molecules
Draw an electron-dot structure for SF6
Step 1: 6 + 6(7) = 48 valence electrons
F
F
F
F
S
Step 2:
F
F
S
Step 3:
F
F
F
F
F
F
Chapter 7/43
Electron-Dot Structures of
Polyatomic Molecules
Draw an electron-dot structure for ICl3
Step 1: 7 + 3(7) = 28 valence electrons
Cl
Cl
I
Step 2:
Cl
Cl
Cl
Step 4:
Cl
I
Cl
Cl
Step 3:
Cl
I
Chapter 7/44
Example 7.6 Drawing an Electron-Dot Structure
. Draw an electron-dot structure for phosphorus pentachloride, PCl5
Solution
Example 7.7 Drawing an Electron-Dot Structure
Draw an electron-dot structure for formaldehyde, CH2O, a compound used in
manufacturing the adhesives for making plywood and particle board.
Solution
Example 7.8 Drawing an Electron-Dot Structure
Draw an electron-dot structure for XeF5+, one of the very few noble-gas ions.
Solution
Electron-Dot Structures and
Resonance
Resonance Structures
• When multiple structures can be drawn, the actual
structure is an average of all possibilities.
• The average is called a resonance hybrid.
A straight double-headed arrow indicates
resonance.
Chapter 7/48
Electron-Dot Structures and
Resonance
Draw an electron-dot structure for O3 (ozone)
Step 1: 3(6) = 18 valence electrons
Step 2: O
O
O
Step 4:
O
O
O
Step 3: O
O
O
Step 5:
O
O
O
Chapter 7/49
Electron-Dot Structures and
Resonance
Move a lone pair from this oxygen?
Step 4: O
O
O
Or, move a lone pair from this oxygen?
O
O
O
O
O
O
Resonance
Chapter 7/50
Example 7.9 Drawing Resonance Structures
The nitrate ion, NO3–, has three equivalent oxygen atoms, and its electronic
structure is a resonance hybrid of three electron-dot structures. Draw them.
Strategy
Formal Charges
1. Determines the best resonance structure.
2. We determine formal charge and estimate the more
accurate representation.
Chapter 7/52
Formal Charges
# of
# of
# of
1
Formal
–
bonding – nonbonding
= valence e –
Charge
2
in free atom
e–
e–
Calculate the formal charge on each atom in O3.
O
1
6 – (4) – 4 = 0
2
O
O
1
6 – (6) – 2 = +1
2
1
6 – (2) – 6 = -1
2
Chapter 7/53
Example 7.10 Calculating Formal Charges
Calculate the formal charge on each atom in the following electron-dot structure for SO2:
Solution
Molecular Shapes: The VSEPR
Model
The approximate shape of molecules is given by
Valence-Shell Electron-Pair Repulsion (VSEPR)
mode.
Electrons in bonds and in lone pairs can be thought of
as “charge clouds” that repel one another and stay as
far apart as possible, thus causing molecules to
assume specific shapes.
Chapter 7/55
Molecular Shapes: The VSEPR
Model
Step 1
• Write an electron-dot structure for the molecule
and count the number of electron charge clouds
surrounding the atom of interest.
Step 2
• Predict the geometric arrangement of charge
clouds by assuming that the charge clouds are
oriented in space as far away from one another
as possible.
Chapter 7/56
Molecular Shapes: The VSEPR
Model
Two Charge Clouds:
Charge clouds point in opposite directions.
Chapter 7/57
Molecular Shapes: The VSEPR
Model
Three Charge Clouds:
Charge clouds lie in the same plane and point to the corners of an equilateral triangle.
Chapter 7/58
Molecular Shapes: The VSEPR
Model
Four Charge Clouds
Charge clouds point to the corners of a regular tetrahedron.
The tetrahedral geometry of an atom with four charge clouds
Chapter 7/59
Molecular Shapes: The VSEPR
Model
Four Charge Clouds
Chapter 7/60
Molecular Shapes: The VSEPR
Model
Five Charge Clouds
Charge clouds point to the corners of a trigonal bipyramid.
Chapter 7/61
Molecular Shapes: The VSEPR
Model
Five Charge Clouds
Chapter 7/62
Molecular Shapes: The VSEPR
Model
Five Charge Clouds
Chapter 7/63
Molecular Shapes: The VSEPR
Model
Six Charge Clouds
Charge clouds point to the corners of a regular octahedron.
Chapter 7/64
Molecular Shapes: The VSEPR
Model
Six Charge Clouds
Chapter 7/65
Molecular Shapes: The VSEPR
Model
Six Charge Clouds
Chapter 7/66
Chapter 7/67
Chapter 7/68
Chapter 7/69
Example 7.11 Using the VSEPR Model to Predict a Shape
Predict the shape of BrF5.
Solution
Six charge clouds imply an octahedral arrangement. The five attached
atoms and one lone pair give BrF5 a square pyramidal shape:
Valence Bond Theory
Valence Bond Theory: A quantum mechanical model
which shows how electron pairs are shared in a
covalent bond.
Chapter 7/71
Valence Bond Theory
Valence Bond Theory: A quantum mechanical model
which shows how electron pairs are shared in a
covalent bond.
Chapter 7/72
Valence Bond Theory
•
Covalent bonds are formed by overlap of atomic
orbitals, each of which contains one electron of
opposite spin.
•
Each of the bonded atoms maintains its own
atomic orbitals, but the electron pair in the
overlapping orbitals is shared by both atoms.
•
The greater the amount of overlap, the stronger
the bond.
Chapter 7/73
Hybridization and sp3 Hybrid
Orbitals
How can the bonding in CH4 be explained?
4 valence electrons
2 unpaired electrons
Chapter 7/74
Hybridization and sp3 Hybrid
Orbitals
How can the bonding in CH4 be explained?
4 valence electrons
4 unpaired electrons
Chapter 7/75
Hybridization and sp3 Hybrid
Orbitals
How can the bonding in CH4 be explained?
4 nonequivalent orbitals
Linus Pauling: Wave functions
from s orbitals & p orbitals could be
combined to form hybrid atomic orbitals.
Chapter 7/76
Hybridization and sp3 Hybrid
Orbitals
How can the bonding in CH4 be explained?
4 equivalent orbitals
Chapter 7/77
Hybridization and sp3 Hybrid
Orbitals
The formation of four sp3 hybrid orbitals by combination
of an atomic s orbital with three atomic p orbitals
Chapter 7/78
Hybridization and sp3 Hybrid
Orbitals
The bonding in methane (CH4)
Chapter 7/79
Hybridization and sp3 Hybrid
Orbitals
Tetrahedron
109.5o
trigonal pyramidal
107.3o
Bent
104.5o
Chapter 7/80
Other Kinds of Hybrid Orbitals
The formation of sp2 hybrid orbitals by combination of one s orbital and two p orbitals
Ethylene, H2C=CH2
Chapter 7/81
Other Kinds of Hybrid Orbitals
The structure of a carbon-carbon double bond
5 s bond + 1 p bond
Chapter 7/82
Other Kinds of Hybrid Orbitals
sp hybridization
Acetylene
H-C≡C-H
Chapter 7/83
Other Kinds of Hybrid Orbitals
Formation of a triple bond by two sp-hybridized atoms
3 s bond + 2 p bond
Chapter 7/84
Other Kinds of Hybrid Orbitals
Chapter 7/85
Example 7.12 Predicting The Hybridization of an Atom
Describe the hybridization of the carbon atoms in allene,
make a rough sketch of the molecule showing its hybrid orbitals.
Draw an electron-dot structure to find the number of charge clouds on each atom.
Then predict the geometry around each atom using the VSEPR model (Table 7.5).
and
Example 7.12 Predicting The Hybridization of an Atom
Continued
Solution
Because the central carbon atom in allene has two charge
clouds (two double bonds), it has a linear geometry and is sphybridized. Because the two terminal carbon atoms have
three charge clouds each (one double bond and two C — H
bonds), they have trigonal planar geometry and are sp2hybridized.
Molecular Orbital Theory: The
Hydrogen Molecule
Atomic Orbital: A wave function whose square gives
the probability of finding an electron within a given
region of space in an atom.
Molecular Orbital: A wave function whose square
gives the probability of finding an electron within a
given region of space in a molecule.
The molecular orbital (MO) model provides a
better explanation of chemical and physical
properties than the valence bond (VB) model.
Chapter 7/88
Molecular Orbital Theory: The
Hydrogen Molecule
s bonding orbital (additive combination, lower energy)
s* antibonding orbital (subtractive combination, higher
energy, sigma star)
Formation of molecular orbitals in the H2 molecule
Chapter 7/89
Molecular Orbital Theory: The
Hydrogen Molecule
1. Molecular orbital theory says that there are two
ways for the orbital interaction to occur –- an
additive way and a subtractive way.
2. Additive combination of orbitals (s) is lower in
energy than two isolated 1s orbitals and is called a
bonding molecular orbital.
3. Subtractive combination of orbitals (s*) is higher in
energy than two isolated 1s orbitals and is called an
antibonding molecular orbital.
Chapter 7/90
Molecular Orbital Theory: The
Hydrogen Molecule
s bonding orbital
s* antibonding orbital
Bond Order =
(# bonding e– – # antibonding e–)
2
Chapter 7/91
Molecular Orbital Theory: The
Hydrogen Molecule
2–0
Bond Order =
=1
2
Energy levels of molecular orbitals for the H2 molecule.
Chapter 7/92
Molecular Orbital Theory: The
Hydrogen Molecule
Bond Order:
2–1
1
=
2
2
2–2
=0
2
Energy levels of molecular orbitals for (a) the stable H2- ion
and (b) the unstable He2 molecule
Chapter 7/93
Key Ideas of Molecular Orbital
Theory
1. Molecular orbitals are to molecules what atomic orbitals are to atoms.
2. Molecular orbitals are formed by combining atomic orbitals on different
atoms.
3. Molecular orbitals that are lower in energy than the starting atomic
orbitals are bonding , and Molecular orbitals that are higher in energy
than the starting atomic orbitals are antibonding.
4. Electrons occupy molecular orbitals beginning with the Molecular
orbital of lowest energy. A maximum of two electrons can occupy each
orbital, and their spins are paired.
5. Bond order can be calculated by subtracting the number of electrons
in antibonding MOs from the number in bonding MOs and dividing the
difference by 2.
Chapter 7/94
Molecular Orbital Theory:
Other Diatomic Molecules
O2
O
O
Diamagnetic: All electrons are spin-paired. It is
weakly repelled by magnetic fields.
Paramagnetic: There is at least one unpaired
electron. It is weakly attracted by magnetic fields.
Oxygen, O2, is predicted to be diamagnetic by
electron-dot structures and valence bond theory.
Chapter 7/95
Molecular Orbital Theory:
Other Diatomic Molecules
O2
O
O
However, it is known to be paramagnetic.
Why does liquid O2 stick to the poles of a magnet?
Chapter 7/96
Molecular Orbital Theory:
Other Diatomic Molecules
Energy levels of molecular orbitals for (a) N2 and (b) O2 and F2
Chapter 7/97
Molecular Orbital Theory:
Other Diatomic Molecules
Chapter 7/98
Molecular Orbital Theory:
Other Diatomic Molecules
Energy levels of molecular orbitals for the second-row diatomic molecules
(a) N2 (b) O2, and (c) F2
Chapter 7/99
Bond Order
1. Bond Order is the number of electron pairs shared
between atoms.
2. Bond Order is obtained by subtracting the number
of antibonding electrons from the number of
bonding electrons and dividing by 2.
3. Bond Order
N2
O2
F2
(8-2)/2 = 3
(8-4)/2 = 2
(8-6)/2 = 1
Chapter 7/100
Combining Valence Orbital Theory
and Molecular Orbital Theory
1. Valence bond theory is better because of its simplicity
and ease of visualization, but MO theory is better
because of its accuracy.
2. Ozone is a resonance hybrid of two equivalent structure,
both of which have two O-O s bonds and one O=O p
bond.
3. The actual structure of O3 is an average of the two
resonance.
4. The s bonds are best described in valence bond
terminology as being localized between pairs of atoms,
and the p electrons are best described by MO theory as
being delocalized over the entire molecule.
Chapter 7/101
Combining Valence Orbital Theory
and Molecular Orbital Theory
The structure of ozone
Chapter 7/102