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Molecular Geometry and Bonding Theories
AP Chemistry – Ch 9Mr. Christopherson
Bonding Theories & Geometry
• Molecular Geometry (shapes)• VSEPR Theory• Lewis Structures• Molecular Polarity (dipoles)• Covalent Bonds• Hybridization• Ionic Bonds
molecular formula
structural formula
molecular shape
ball-and-stick model
CH4 C
H
H
HH
H
H
H
H
109.5o
C
tetrahedrontetrahedralshape ofmethane
CH
H
H
H
109.5o109.5o
TetrahedronTetrahedron
CentralAtom
CentralAtom
CentralAtom
CentralAtom
SubstituentsSubstituents
SubstituentsSubstituents
Methane, CH4Methane, CH4
Tetrahedralgeometry
Tetrahedralgeometry
Methane, CH4Methane, CH4
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Methane & Carbon Tetrachloridemolecular formula
structural formula
molecular shape
ball-and-stick model
CH4 C
H
H
HH
H
H
H
H
109.5o
C
CCl4
space-filling model
C
Cl
Cl
ClCl
Molecular Geometry
H
H
H
H
109.5o
C
Linear Trigonal planar
Tetrahedral
Trigonal pyramidalBent
109.5o
107.3o104.5o
H2O CH4 AsCl3 AsF5 BeH2 BF3 CO2
180o
A Lone PairA Lone Pear
C109.5o
H
HHH
N107o HH
H
..
O104.5o H
H
..
..
CH4, methane NH3, ammonia H2O, water
..
O
O
O
lone pairelectrons
OOO
O3, ozone
Molecular ShapesThree atoms (AB2) Four atoms (AB3)
Five atoms (AB4)
Six atoms (AB5)
Seven atoms (AB6)
• Linear (180o)• Bent
• Trigonal planar (120o)• Trigonal pyramidal• T-shaped
• Tetrahedral (109.47o)• Square planar• Seesaw
• Trigonal bipyramidal (BeABe, 120o) & (BeABa, 90o)• Square pyramidal
• Octahedral
B BA
B
B
A
B
linear trigonal planarB
A
BB
B
tetrahedral
A Be
Be
Be
Ba
Ba
TrigonalbipyramidalB
B
B
B
B
BA
Bailar, Moeller, Kleinberg, Guss, Castellion, Metz, Chemistry, 1984, page 313.
Bonding and Shape of Molecules
Number of Bonds
Number of Unshared Pairs Shape Examples
2
3
4
3
2
0
0
0
1
2
Linear
Trigonal planar
Tetrahedral
Pyramidal
Bent
BeCl2
BF3
CH4, SiCl4
NH3, PCl3
H2O, H2S, SCl2
-Be-
B
C
N
:
O
:
:
CovalentStructure
Molecular ShapesAB2
Linear
AB3
Trigonal planar AB4
Tetrahedral
AB5
Trigonal bipyramidal
AB6
Octahedral
AB2EAngular or Bent AB3E
Trigonalpyramidal
AB2E2
Angular or Bent
AB4EIrregular tetrahedral(see saw)
AB3E2
T-shaped
AB2E3
Linear
AB6ESquare pyramidal
AB5E2
Square planar
Valence
Shell
Electron
Pair
Repulsion
Theory
Planar triangular
Tetrahedral
Trigonal bipyramidal
Octahedral
Valence
Shell
Electron
Pair
Repulsion
Theory
Planar triangular
Tetrahedral
Trigonal bipyramidal
Octahedral
......
The VSEPR Model
O OC
Linear
The Shapes of Some Simple ABn Molecules
O OS
BentO O
S
O
Trigonalplanar
FF
F
N
Trigonalpyramidal
T-shaped Squareplanar
F FCl
F
F F
Xe
F FF
F
FP
F
F
Trigonalbipyramidal
Octahedral
FF
F
S
F
F
F
SF6
SO2
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 305
Molecular ShapesAB2
Linear
AB3
Trigonal planarAB4
Tetrahedral
AB5
Trigonal bipyramidal
AB6
Octahedral
AB2EAngular or Bent
AB3ETrigonal
pyramidal
AB2E2
Angular or Bent
AB4EIrregular tetrahedral(see saw)
AB3E2
T-shaped
AB2E3
Linear
AB5ESquare pyramidal
AB4E2
Square planar
Geometry of Covalent Molecules ABn, and ABnEm
AB2
AB2EAB2E2
AB2E3
AB3
AB3E
AB3E2
AB4
AB4E
AB4E2
AB5
AB5EAB6
222233
34
4
45
56
012301
20
1
20
10
LinearTrigonal planarTetrahedralTrigonal bipyramidalTrigonal planarTetrahedral
Triangular bipyramidalTetrahedral
Triangular bipyramidal
OctahedralTriangular bipyramidal
OctahedralOctahedral
LinearAngular, or bentAngular, or bentLinearTrigonal planarTriangular pyramidal
T-shapedTetrahedral
Irregular tetrahedral (or “see-saw”)Square planarTriangular bipyramidal
Square pyramidalOctahedral
CdBr2
SnCl2, PbI2
OH2, OF2, SCl2, TeI2
XeF2
BCl3, BF3, GaI3
NH3, NF3, PCl3, AsBr3
ClF3, BrF3
CH4, SiCl4, SnBr4, ZrI4
SF4, SeCl4, TeBr4
XeF4
PF5, PCl5(g), SbF5
ClF3, BrF3, IF5
SF6, SeF6, Te(OH)6, MoF6
TypeFormula
Shared Electron
Pairs
Unshared Electron
Pairs
IdealGeometry
ObservedMolecular Shape Examples
Bailar, Moeller, Kleinberg, Guss, Castellion, Metz, Chemistry, 1984, page 317.
Predicting the Geometry of Molecules
• Lewis electron-pair approach predicts number and types of bonds between the atoms in a substance and indicates which atoms have lone pairs of electrons but gives no information about the actual arrangement of atoms in space
• Valence-shell electron-pair repulsion (VSEPR) model predicts the shapes of many molecules and polyatomic ions but provides no information about bond lengths or the presence of multiple bonds
Introduction to Lewis Structures
Lewis dot symbols
1. Used for predicting the number of bonds formed
by most elements in their compounds
2. Consists of the chemical symbol for an element surrounded by dots that represent its valence
electrons
3. A single electron is represented as a single dot
Lewis Structures
1) Count up total number of valence electrons
2) Connect all atoms with single bonds
- “multiple” atoms usually on outside
- “single” atoms usually in center;
C always in center,
H always on outside.
3) Complete octets on exterior atoms (not H, though)
4) Check
- valence electrons math with Step 1
- all atoms (except H) have an octet; if not, try multiple bonds
- any extra electrons? Put on central atom
Molecules with Expanded Valence ShellsAtoms that have expanded octets have AB5 (trigonal bipyramidal)
or AB6 (octahedral) electron domain geometries.
Trigonal bipyramidal structures have a plane containing three electron pairs.
• The fourth and fifth electron pairs are located above and below this plane.• In this structure two trigonal pyramids share a base.
For octahedral structures, there is a plane containing four electron pairs.
• Similarly, the fifth and sixth electron pairs are located above and below this plane.• Two square pyramids share a base.
F
F
FP
F
F
FF
F
S
F
F
F
Trigonal Bipyramid F
F
FP
F
F• The three electron pairs in the plane are called equatorial.
• The two electron pairs above and below this plane are called axial.
• The axial electron pairs are 180o apart and 90o from to the equatorial electrons.
• The equatorial electron pairs are 120o apart.
• To minimize electron-electron repulsions, nonbonding pairs are always placed in equatorial positions, and bonding pairs in either axial or equatorial positions.
Octahedron• The four electron pairs in the plane are 90o to each other.
• The remaining two electron pairs are 180o apart and 90o from the electrons in the plane.
• Because of the symmetry of the system, each position is equivalent.
• The equatorial electron pairs are 120o apart.
• If we have five bonding pairs and one nonbonding pair, it doesn’t matter where the nonbonding pair is placed.
The molecular geometry is square pyramidal.
• If two nonbonding pairs are present, the repulsions are minimized by pointing them toward opposite sides of the octahedron.
The molecular geometry is square planar.
FF
F
S
F
F
F
F F
Xe
F F
Electron-Domain GeometriesNumber of Electron Domains
Arrangement ofElectron Domains
Electron-DomainGeometry
Predicted Bond Angles
2
3
4
5
6
Linear
Trigonalplanar
Tetrahedral
Trigonal-bipyramidal
Octahedral
180o
120o
109.5o
120o
90o
90o
A Be
Be
Be
Ba
Ba
B
B
B
B
B
BA
B
B
A
B
B
A
BB
B
B BA
Number of electron domains
Electron-domain geometry
Predicted bond angles
TetrahedralTrigonalplanar Tetrahedral
109.5o 120o 109.5o
C C OH H
H
H O
4 3 4
Acetic Acid, CH3COOH
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 314
Hybridization of central atom sp3 sp2 none
Intermolecular Forces
Ion-ion (ionic bonds)
Ion-dipole
Dipole-dipole
Hydrogen bonding
London dispersion forces
+ −− +
+ − + −
HO
H
HO
H
H
O
H
−+
London Dispersion Forces
• London dispersion forces are created when on molecule with a temporarily dipole causes another to become temporarily polar.
+
−
+
− +
−
Molecular Polarity
Molecular Structure
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Electronegativity
+ – 0 0
H Cl H H
Na
Ionic vs. Covalent• Ionic compounds form repeating units.• Covalent compounds form distinct molecules.• Consider adding to NaCl(s) vs. H2O(s):
HO
H Cl
Cl
Cl
ClH
OH
H
O H
• NaCl: atoms of Cl and Na can add individually forming a compound with million of atoms.
• H2O: O and H cannot add individually, instead molecules of H2O form the basic unit.
Na Na
Na
Cl ClNa Na
Holding it togetherQ: Consider a glass of water.
Why do molecules of water stay together?A: There must be attractive forces.
Intramolecular forces occur between atoms
Intermolecular forces occur between molecules
• Intermolecular forces are not considered in ionic bonding because there are no molecules.
• The type of intramolecular bond determines the type of intermolecular force.
Intramolecular forces are much
stronger
I’m not stealing, I’m sharing unequally• We described ionic bonds as stealing electrons• In fact, all bonds share – equally or unequally.• Note how bonding electrons spend their time:
• Bonding electrons are shared in each compound, but are NOT always shared equally.
• The greek symbol indicates “partial charge”.
H2 HCl LiCl
+ –0 0 + –
covalent (non-polar) polar covalent ionic
H H H Cl [Li]+[ Cl ]–
Dipole Moment
• Direction of the polar bond in a molecule.• Arrow points toward the more
electronegative atom.
H Cl+ -
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Dipole-induced dipole attraction
Dipole-induced dipole attraction
The attraction between a dipole and an induced
dipole.
The attraction between a dipole and an induced
dipole.
Oxygen, O2Oxygen, O2
Oxygen, O2Oxygen, O2
NonpolarNonpolar
Water, H2OWater, H2O
Water, H2OWater, H2O
d-d-++
d-d-++
d-d-++
d-d-++
++--
d-d-++
++ --
++ --d-d-++
++ --d-d-++
++ --d-d-++
++ --
++
induced dipole
induced dipole
DipoleDipole
d-d-
++ --d-d-++
++ --d-d-++
++ --d-d-++
d-d-++
++ --
d-d-++
++--
d-d-++
++--
++d-d-
++--
++d-d-
induced dipole
induced dipoleDipoleDipole
PolarPolar NonpolarNonpolar
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Determining Molecular Polarity
• Depends on:– dipole moments– molecular shape
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
+–
+ –
+ –
+ –
H Cl
+ –
BF3
F
F F
B
Determining Molecular Polarity
• Nonpolar Molecules– Dipole moments are symmetrical and cancel
out.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
H2OH H
O
Determining Molecular Polarity
• Polar Molecules– Dipole moments are asymmetrical and don’t
cancel .
netdipolemoment
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
CHCl3
H
Cl ClCl
Determining Molecular Polarity
• Therefore, polar molecules have...– asymmetrical shape (lone pairs) or – asymmetrical atoms
netdipolemoment
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Dipole Moment
Nonpolar
Polar
....
H H
O
C OO
Bond dipoles
Overall dipole moment = 0
Bond dipoles
Overall dipole moment
The overall dipole moment of a moleculeis the sum of its bond dipoles. In CO2 thebond dipoles are equal in magnitude butexactly opposite each other. The overall dipole moment is zero.
In H2O the bond dipoles are also equal inmagnitude but do not exactly oppose eachother. The molecule has a nonzero overall dipole moment.
221
dqqk
F Coulomb’s lawm = Q r Dipole moment, m
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 315
......
Polar Bonds
H Cl
Polar
A molecule has a zero dipole moment because their dipoles cancel one another.
H HO
PolarF F
B
F
Nonpolar
HH
H
N
Polar
Polar Nonpolar
F FCl
F
F F
Xe
F FCl
ClC
Cl
Nonpolar Polar
Cl
HC
Cl
H
H
HF HCl HBr HI
Mark Wirtz, Edward Ehrat, David L. Cedeno*
How is the electron density distributed in these different molecules?
Based on your comparison of the electron density distributions, which molecule should have the most polar bond, and which one the least polar?
Arrange the molecules in increasing order of polarity.
CH3Cl CH2Cl2 CHCl3 CCl4
Mark Wirtz, Edward Ehrat, David L. Cedeno*
Describe how is the electron density distributed in these different molecules? Based on your comparison of the electron density distributions, which molecule(s) should be the most polar, and which one(s) the least polar?
Arrange the molecules in increasing order of polarity.
Benzene NO3-
Nitrobenzene
Mark Wirtz, Edward Ehrat, David L. Cedeno*
2s 2p (x, y, z) carbon
Mark Wirtz, Edward Ehrat, David L. Cedeno*
How does H2 form?
+ +
The nuclei repel
But they are attracted to electrons
They share the electrons
Hydrogen Bond Formation
0.74 A
- 436
0
H – H distance
Ene
rgy
(KJ/
mol
)
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 318
no interaction
increasedattraction
balanced attraction& repulsion
increasedrepulsion
Potential Energy Diagram - Attraction vs. Repulsion
(internuclear distance)
Covalent bonds
• Nonmetals hold onto their valence electrons.• They can’t give away electrons to bond.• Still want noble gas configuration.
• Get it by sharing valence electrons with each other.
• By sharing both atoms get to count the electrons toward noble gas configuration.
1s22s22p63s23p6…eight valence electrons (stable octet)
Covalent bonding
• Fluorine has seven valence electrons
F
•A second atom also has seven
F
By sharing electrons…both end with full orbitals
8 Valence electrons
8 Valence electrons
Single Covalent Bond
• A sharing of two valence electrons.• Only nonmetals and Hydrogen.• Different from an ionic bond because they
actually form molecules.• Two specific atoms are joined.• In an ionic solid you can’t tell which atom
the electrons moved from or to.
Sigma bonding orbitals
• From s orbitals on separate atoms
+ +
s orbital s orbital
+ ++ +
Sigma bondingmolecular orbital
Sigma bonding orbitals
• From p orbitals on separate atoms
p orbital p orbital
Sigma bondingmolecular orbital
Pi bonding orbitals
• P orbitals on separate atoms
Pi bondingmolecular orbital
Sigma and pi bonds
• All single bonds are sigma bonds• A double bond is one sigma and one pi
bond• A triple bond is one sigma and two pi
bonds.
Atomic Orbitals and Bonding• Bonds between atoms are formed by electron pairs in
overlapping atomic orbitals
1s 1s:E 1s• Example: H2 (H-H)
– Use 1s orbitals for bonding
• Example: H2O– From VSEPR: bent, 104.5°
angle between H atoms– Use two 2p orbitals for bonding?
E2s
2p
90° How do we explain the structure predicted by VSEPR
using atomic orbitals?2p
2p
1s
1s
Overlapping OrbitalsDraw orbital diagrams for F + F, H + O, Li + F
1s 2s 2p
1s2s2p
1s 2s 2p
1s
1s
F2
H2O
1s 2s 1s2s2p
LiF is ionic (metal + non-metal)
FLi
electron transfer
1+ 1-
e-
3p+
lithium atomLi
e-
loss of one valence
electron
e-
e-
lithium ionLi+
3p+e-
e-
9p+
fluorine atomF
e-
e-
e-
e-
e-
e-
e- e-
e-
gain of one valence
electron
fluoride ionF1-
10p+e-
e-
e-
e-e-
e-
e-
e-
e-
e-
e-
Formation of Cation
3p+
lithium atomLi
e-
loss of one valence
electron
e-
e-
lithium ionLi+
3p+e-
e-
e-
Formation of Anion
9p+
fluorine atomF
e-
e-
e-
e-
e-
e-
e- e-
e-
e-
gain of one valence
electron
fluoride ionF1-
10p+e-
e-
e-
e-e-
e-
e-
e-
e-
e-
Formation of Ionic Bond
fluoride ionF1-
9p+e-
e-
e-
e-e-
e-
e-
e-
e-
e-
lithium ionLi+
3p+e-
e-
Be
H
H
BeH2
s p
First, the formation of BeH2 using pure s and p orbitals.
The formation of BeH2 using hybridized orbitals.
atomic orbitals atomic orbitals
Be
s p
Be H
H
s p
atomic orbitals
hybrid orbitals
No overlap = no bond!
sp p
Be HH
All hybridized bonds have equal strength and have orbitals with identical energies.
BeH2Be
Be = 1s22s2
Hybrid OrbitalsGround-state Be atom
1s 2s 2p
1s 2s 2p
Be atom with one electron “promoted”
s
px py pz
sp
hybrid orbitals
Ene
rgy
hybridize
s orbital p orbital
two sp hybrid orbitals sp hybrid orbitals shown together(large lobes only)
1s sp 2p
Be atom of BeH2 orbital diagram
H HBe
n = 1
n = 2
Hybrid Orbitals
2s 2p
Ground-state B atom
s
px py pzEne
rgy
sp2 2p
B atom of BH3 orbital diagram
hybridize
s orbital
2s 2p
B atom with one electron “promoted”
sp2
hybrid orbitals
p orbitals sp2 hybrid orbitals shown together
(large lobes only)three sps hybrid orbitals
H
H
HB
Hybridization…the blending of orbitals
Valence bond theory is based on two assumptions:
1. The strength of a covalent bond is proportional to the amount of overlap between atomic orbitals; the greater
the overlap, the more stable the bond.
2. An atom can use different combinations of atomic orbitals to maximize the overlap of orbitals used by bonded
atoms.
We have studied electron configuration notation and the sharing of electrons in the formation of covalentbonds.
Methane is a simple natural gas. Its molecule has a carbon atom at the center with four hydrogen atoms covalently bonded around it.
Lets look at amolecule of methane, CH4.
What is the expected orbital notation of carbonin its ground state?
(Hint: How many unpaired electrons does this carbon atom have available for bonding?)
Can you see a problem with this?
Carbon ground state configuration
1s2s
2p
You should conclude that carbon only has TWO electrons available for bonding. That is not enough!
How does carbon overcome this problem so thatit may form four bonds?
The first thought that chemists had was that carbon promotes one of its 2s electrons…
…to the empty 2p orbital.
Carbon’s Empty Orbital
1s2s
2p
1s2s
2p
1s2s
2p
Non-hybridized orbital hybridized orbital
However, they quickly recognized a problem with such an arrangement…
Three of the carbon-hydrogen bonds would involvean electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom.
A Problem Arises
1s2s
2p
1s 1s1s 1s
Unequal bond energy
But what about the fourth bond…?
The fourth bond is between a 2s electron from thecarbon and the lone 1s hydrogen electron.
Such a bond would have slightly less energy than the other bonds in a methane molecule.
Unequal bond energy #2
1s2s
2p
1s 1s1s 1s
This bond would be slightly different in character than the other three bonds in methane.
This difference would be measurable to a chemistby determining the bond length and bond energy.
But is this what they observe?
The simple answer is, “No”.
Chemists have proposed an explanation – they call ithybridization.
Hybridization is the combining of two or more orbitalsof nearly equal energy within the same atom into orbitals of equal energy.
Measurements show that all four bonds in methane are equal. Thus, we need a new explanation for the bonding in methane.
Enter Hybridization
In the case of methane, they call the hybridization sp3, meaning that an s orbital is combined with threep orbitals to create four equal hybrid orbitals.
These new orbitals have slightly MORE energy thanthe 2s orbital…
… and slightly LESS energy than the 2p orbitals.
sp3 Hybrid Orbitals
s
px py pz
Carbon 1s22s22p2
Carbon could only make two bondsif no hybridization occurs. However,carbon can make four equivalent bonds.
sp3
hybrid orbitals
Ene
rgy
sp3
C atom of CH4 orbital diagram
B
A
BB
B
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 321
Hybridization of s and p Orbitals• The combination of an ns and an np orbital
gives rise to two equivalent sp hybrids oriented at 180º.
• Combination of an ns and two or three np orbitals produces three equivalent sp2 hybrids or four equivalent sp3 hybrids.
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Hybridization of s and p Orbitals
• Both promotion and hybridization require an input of energy; the overall process of forming a compound with hybrid orbitals will be energetically favorable only if the amount of energy released by the formation of covalent bonds is greater than the amount of energy used to form the hybrid orbitals.
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Hybridization Involving d Orbitals
3s 3p 3d 3s 3p 3d
promote
five sp3d orbitals 3dF
F
FP
F
F
A Be
Be
Be
Ba
Ba
Trigonal bipyramidal
hybridize
degenerateorbitals
(all EQUAL)
unhybridized P atomP = [Ne]3s23p3
vacant d orbitals
Pure atomicorbitals of
central atom
Hybridizationof the central
atom
Numberof hybridorbitals
Shape of hybridorbitals
s,p
s,p,p
s,p,p,p
s,p,p,p,d
s,p,p,p,d,d
sp
sp2
sp3
sp3d
sp3d2
2
3
4
5
6
Linear
Trigonal Planar
Tetrahedral
TrigonalBipyramidal
Octahedral
Hybridization Animation, by Raymond Chang
Hybridization Animation, by Raymond Chang
Bonding
• Single bonds– Overlap of bonding orbitals on bond axis– Termed “sigma” or σ bonds
• Double bonds– Sharing of electrons between 2 p orbitals
perpendicular to the bonding atoms– Termed “pi” or π bonds
2p 2p
Bond Axis of σ bond
One π bond
Multiple Bonds
2s 2p 2s 2p sp2 2p
promote hybridize
C C
H
H H
H
C2H4, ethene
one s bond and one p bond
H
H
CC
H
H s
ss
s
s
H
H
CC
H
H
Two lobes ofone p bond
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 325-326
Multiple Bonds
2s 2p 2s 2p sp2 2p
promote hybridize
C C
H
H H
HC2H4, ethene
one s bond and one p bond
H
H
CC
H
H s
ss
s
s
H
H
CC
H
H
Two lobes ofone p bond
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 325-326
C C
H
H
sp2
sp2
sp2
H
H
sp2
sp2
sp2
p p
p p
p bond
Internuclear axis
p p
s bonds
H
CHC
H
C
H
C
H C
H
C
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329
C6H6 = benzene
2p atomic orbitals
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329
s bonds and p bonds
H
C
HCH
C
H
C
H C
H
C
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329
s bondsH
CH
C
HC
H
C
H
C
H
C
HC
H
C
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329
s bondsH
C
H
C
HC
H
C
H
C
H
C
HC
H
C
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 329
N NO
O
O
O
dinitrogen tetraoxide
N2O4
NO
ON
O
O
hn 2 NO2nitrogen dioxide
(free radical)
colorless red-brown
1s 1s
s1s
s*1s
1s 1s
s1s
s*1s
H2 molecule
He2 molecule
H atom H atom
He atom He atom
Ene
rgy
Ene
rgy
Energy-level diagram for (a) the H2 molecule and (b) the hypothetical He2 molecule
(a)
(b)
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 332
Bond Order
Bond order = ½ (# or bonding electrons - # of antibonding electrons)
• A bond order of 1 represents a single bond,
• A bond order of 2 represents a double bond,
• A bond order of 3 represents a triple bond.
• A bond order of 0 means no bond exists.
Because MO theory also treats molecules with an odd number of electrons, Bond orders of 1/2 , 3/2 , or 5/2 are possible.
1s2 1s2
s1s
s*1s
2s1 2s1
s2s
s*2s
Energy-level diagram for the Li2 molecule
Li = 1s22s1
Li Li
Li2
Ene
rgy
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 334
2s 2s
s2s
s*2s
2p 2p
s2p
s*2p
p2p
p*2p
Energy-level diagram for molecular orbitals of second-row homonuclear diatomic molecules.
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 337
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 338
s2s
s*2s
s2p
p2pEnergy of molecular orbitals
O2, F2, Ne2 B2, C2, N2
Increasing 2s – 2p interaction
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 338
s2s
*s 2s
p2p
s2p
*p 2p
*s 2p
Large 2s – 2p interaction Small 2s – 2p interaction
s2s
*s 2s
s2p
p2p
*p 2p
*s 2p
C2 N2B2 F2 Ne2O2
Bond order
Bond enthalpy (kJ/mol)
Bond length (angstrom)
Magnetic behavior
1 2 3 2 1 0
290 620 941 495 155 -----
1.59 1.31 1.10 1.21 1.43 -----
Paramagnetic Diamagnetic Diamagnetic Paramagnetic Diamagnetic _____
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 339
s2s
s*2s
p2px
p*2px
p2py
p*2py
s2p
s*2p
C2
Arrange the atomic and molecular orbitals in order of increasing energy.How many orbitals are per molecule?
Can you distinguish the bonding from the antibonding MOs?
Mark Wirtz, Edward Ehrat, David L. Cedeno*
Magnetic Properties of a Sample
PARAMAGNETISM – molecules with one or more unpaired electrons are attracted into a magnetic field. (appears to weigh MORE in a magnetic field)
DIAMAGNETISM – substances with no unpaired electrons are weakly repelled from
a magnetic field. (appears to weigh LESS in a magnetic field)
Experiment for determining the magnetic properties of a sample
The sample is first weighed in the absence of a magnetic field.
When a field is applied, a diamagneticsample tends to move out of the fieldand appears to have a lower mass.
A paramagnetic sample is drawninto the field and thus appears to gain mass.
Paramagnetism is a much stronger effect than is diamagnetism.
sample N SN SN S
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 339
Experiment for determining the magnetic properties of a sample
N S N S
The sample is first weighed in the absence of a magnetic field.
When a field is applied, a diamagneticsample tends to move out of the fieldand appears to have a lower mass.
A paramagnetic sample is drawninto the field and thus appears to gain mass.
Paramagnetism is a much stronger effect than is diamagnetism.
sample
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 339
lone Pair single bond double bond triple bond
Electron Domains
Determine the shape of the BCl3 molecule:
B
Cl
ClCl
: :
::
:: :
:
:
There are 3 electron domains aboutthe central atom: no lone pairs andthree single bonds. Three electrondomains arrange themselves in atrigonal plane, with 120o angles.We predict a trigonal planar geometry.
:
B
Cl
ClCl
: :
::
:: :
:
Electron-domain geometry: trigonal planar
Molecular geometry (shape):
trigonal planar
One s orbital
Two p orbitals
Three sp2 hybrid orbitals
sp2 hybrid orbitalsshown together(large lobes only)
Hybridize
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Ammonia, NH3Ammonia, NH3
Ammonia, NH3Ammonia, NH3
Triangular pyramidalTriangular pyramidal
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Introduction toBonding
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Chemical bond — the force that holds atoms together in a chemical compound
Covalent bonding — electrons are shared between atoms in a molecule or polyatomic ion
Ionic bonding — positively and negatively charged ions are held together by electrostatic forces
Ionic compounds — dissolve in water to form aqueous solutions that conduct electricity
Covalent compounds — dissolve to form solutions that do not conduct electricity
Vocabulary Chemical Bond
–attractive force between atoms or ions that binds them together as a unit
–bonds form in order to…• decrease potential energy (PE)• increase stability
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Vocabulary
CHEMICAL FORMULA
molecularformula
formulaunit
IONIC COVALENT
CO2NaClCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Vocabulary
COMPOUND
ternarycompound
binarycompound
2 elementsmore than 2
elements
NaNO3NaClCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Vocabulary
ION
polyatomicIon
monatomicIon
1 atom 2 or more atoms
NO3-Na+
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
IONIC COVALENT
Bond Formation
Type of Structure
Solubility in Water
Electrical Conductivity
OtherProperties
e- are transferred from metal to nonmetal
high
yes (solution or liquid)
yes
e- are shared between two nonmetals
low
no
usually not
MeltingPoint
crystal lattice true molecules
Types of Bonds
Physical State
solid liquid or gas
odorousCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
“electron sea”
METALLICBond Formation
Type of Structure
Solubility in Water
Electrical Conductivity
OtherProperties
MeltingPoint
Types of Bonds
Physical State
e- are delocalized among metal atoms
very high
yes (any form)
no
malleable, ductile, lustrous
solid
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Lattice Energies in Ionic Solids
Ionic compounds
1. Usually rigid, brittle, crystalline substances with flat surfaces that intersect at characteristic angles
2. Not easily deformed
3. Melt at relatively high temperatures
4. Properties result from the regular arrangement of the ions in the crystalline lattice and from the strong electrostatic attractive forces between ions with opposite charges
Metallic Bonding - “Electron Sea”
Types of Bonds
Bond Polarity Most bonds are
a blend of ionic and covalent characteristics.
Difference in electronegativity determines bond type.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Ionic
Polar-covalent
Nonpolar-covalent
3.3
1.7
0.3
0
100%
50%
5%
0%
Diff
eren
ce in
ele
ctro
nega
tiviti
es Percentage ionic character
Types of Chemical Bonds
Copyright © 2006 Pearson Education Inc., publishing as Benjamin Cummings
Bond Polarity
Electronegativity–Attraction an atom has for a shared
pair of electrons.–higher e-neg atom -
– lower e-neg atom +
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Ionic bonding: Li + ClIonic bonding (stealing/transfer of electrons) can be represented in three different ways
Li + Cl [Li]+[Cl]–
1e-
3p+
4n02e- 17p+
18n08e-8e-2e 3p+
4n02e-1e- 17p+
18n07e- 8e- 2e-
Li Cl [ Cl ]–[Li]+
lithium atom chlorine atom lithium ion chlorine ionchloride ion
Ionic bonding: Mg + O
Mg + O [Mg]2+[O]2–
12p+
12n02e- 8e- 2e-
1e-
[ O ]2–[Mg]2+
6e- 2e-
8n0 8p+
1e-
8e- 2e-
8n0 8p+ 12p+
12n02e- 8e-
OMg
F
4.0
Ar
--
Kr
3.0
Xe
2.6
Rn
2.4
Bond Polarity Electronegativity Trend
– Increases up and to the right.
Be
1.5
Al
1.5
Si
1.8
Ti
1.5
V
1.6
Cr
1.6
Mn
1.5
Fe
1.8
Co
1.8
Ni
1.8
Cu
1.9
Zn
1.7
Ga
1.6
Ge
1.8
Nb
1.6
Mo
1.8
Tc
1.9
Ag
1.9
Cd
1.7
In
1.7
Sn
1.8
Sb
1.9
Ta
1.5
W
1.7
Re
1.9
Hg
1.9
Tl
1.8
Pb
1.8
Bi
1.9
N
3.0
O
3.5
F
4.0
Cl
3.0
C
2.5
S
2.5
Br
2.8
I
2.5
Na
0.9
K
0.8
Rb
0.8
Cs
0.7
Ba
0.9
Fr
0.7
Ra
0.9
H
2.1
B
2.0
P
2.1
As
2.0
Se
2.4
Ru
2.2
Rh
2.2
Pd
2.2
Te
2.1
Os
2.2
Ir
2.2
Pt
2.2
Au
2.4
Po
2.0
At
2.2
Actinides: 1.3 - 1.5
Li
1.0
Ca
1.0
Sc
1.3
Sr
1.0
Y
1.2
Zr
1.4
Hf
1.3
Mg
1.2
La
1.1
Ac
1.1
Lanthanides: 1.1 - 1.3
*
*y
y
Ne
--
He
--
Bond Polarity Electronegativity Trend
– Increases up and to the right.
1
2
3
4
5
6
1
2
3
4
5
6
7
1A
2A
3B 4B 5B 6B 7B 1B 2B
3A 4A 5A 6A 7A
8A
8B
Nonpolar Covalent Bond–electrons are shared equally–symmetrical electron density–usually identical atoms
Bond Polarity
+ -
Bond Polarity Polar Covalent Bond
–electrons are shared unequally–asymmetrical e- density– results in partial charges (dipole)
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Nonpolar
Polar
Ionic
Bond Polarity
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Bond PolarityExamples:
Cl2
HCl
NaCl
3.0 - 3.0 = 0.0 Nonpolar
3.0 - 2.1 = 0.9 Polar
3.0 - 0.9 = 2.1 Ionic
Ionic
Polar-covalent
Nonpolar-covalent
3.3
1.7
0.3
0
100%
50%
5%
0%
Diff
eren
ce in
ele
ctro
nega
tiviti
es Percentage ionic character
Mg
Write the electron dot diagram for
Na Mg C O F Ne He
1s22s22p63s1
1s22s22p63s2
1s22s22p2
1s22s22p4
1s22s22p5
1s22s22p6
1s2
Na
C
O
F
He
Ne
Ionic Bonding
Na Cl
transfer of electron
+ -
NaCl
Ca +2
P -3Ca
+2
P
All the electrons must be accounted for!
+2
Ionic Bonding
Ca -3
Ionic Bonding
Ca3P2
Formula Unit
Ca2+
Ca2+
Ca2+
P3-
P 3-
Ca2+
P3- Ca2+
P3- Ca2+
Metals are Malleable Hammered into shape (bend). Ductile - drawn into wires. Electrons allow atoms to slide by.
+ + + ++ + + +
+ + + +
Ionic solids are brittle
+ - + -+- +-
+ - + -+- +-
Force + - + -
Strong repulsion breaks crystal apart.
How does H2 form?
+ +
The nuclei repel
But they are attracted to electrons
They share the electrons
Hydrogen Bond Formation
0.74 A
- 436
0
H – H distance
Ene
rgy
(KJ/
mol
)
Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 318
no interaction
increasedattraction
balanced attraction& repulsion
increasedrepulsion
Potential Energy Diagram - Attraction vs. Repulsion
(internuclear distance)
Covalent bonds Nonmetals hold onto their valence
electrons. They can’t give away electrons to bond. Still want noble gas configuration. Get it by sharing valence electrons with
each other. By sharing both atoms get to count the
electrons toward noble gas configuration.
Covalent bonding
F F
Fluorine has seven valence electrons A second F atom also has seven By sharing electrons Both end with full orbitals (stable octets)
8 Valence electrons
8 Valence electrons
Single Covalent Bond A sharing of two valence electrons. Only nonmetals and Hydrogen. Different from an ionic bond because
they actually form molecules. Two specific atoms are joined. In an ionic solid you can’t tell which
atom the electrons moved from or to.
How to show how they formed It’s like a jigsaw puzzle. I have to tell you what the final formula
is. You put the pieces together to end up
with the right formula. For example - show how water is
formed with covalent bonds.
Water
H
O
Each hydrogen has 1 valence electron
Each hydrogen wants 1 more
The oxygen has 6 valence electrons
The oxygen wants 2 more
They share to make each other happy
Water Put the pieces together The first hydrogen is happy The oxygen still wants one more
H O
Water The second hydrogen attaches Every atom has full energy levels A pair of electrons is a single bond
H OH H
H O
Lewis Structures
1) Count up total number of valence electrons
2) Connect all atoms with single bonds
- “multiple” atoms usually on outside - “single” atoms usually in center;
C always in center,H always on outside.
3) Complete octets on exterior atoms (not H, though)4) Check - valence electrons math with Step 1 - all atoms (except H) have an octet; if not, try multiple bonds - any extra electrons? Put on central atom
Multiple Bonds Sometimes atoms share more than one
pair of valence electrons. A double bond is when atoms share two
pair (4) of electrons. A triple bond is when atoms share three
pair (6) of electrons.
Carbon dioxide CO2 - Carbon is central
atom ( I have to tell you)
Carbon has 4 valence electrons
Wants 4 more Oxygen has 6 valence
electrons Wants 2 more
O
C
Carbon dioxide Attaching 1 oxygen leaves the oxygen 1
short and the carbon 3 short
OC
Carbon dioxide Attaching the second oxygen leaves
both oxygen 1 short and the carbon 2 short
OCO
The only solution is to share more Requires two double bonds Each atom gets to count all the atoms in
the bond8 valence electrons
8 valence electrons
8 valence electrons
Carbon dioxide
OCO
Formation of Multiple Covalent Bonds
By combining more than one unpaired electron at a time, a double bond is formed.
Both oxygen atoms end up with eight valence electrons.
Ox x
x
xOx
x
x
x
x
xO
x
x O
How to draw them Add up all the valence electrons. Count up the total number of electrons
to make all atoms happy. Subtract. Divide by 2 Tells you how many bonds - draw them. Fill in the rest of the valence electrons
to fill atoms up.
Examples NH3
N - has 5 valence electrons wants 8
H - has 1 valence electrons wants 2
NH3 has 5+3(1) = 8
NH3 wants 8+3(2) = 14
(14-8)/2= 3 bonds 4 atoms with 3 bonds
N
H
N HHH
Examples Draw in the bonds All 8 electrons are accounted for Everything is full
Examples HCN C is central atom N - has 5 valence electrons wants 8 C - has 4 valence electrons wants 8 H - has 1 valence electrons wants 2
HCN has 5 + 4 + 1 = 10
HCN wants 8 + 8 + 2 = 18
(18 - 10) / 2= 4 bonds 3 atoms with 4 bonds -will require
multiple bonds - not to H
HCN Put in single bonds Need 2 more bonds Must go between C and N
NH C
HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons - 2 more to add
NH C
HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons - 2 more to add Must go on N to fill octet
NH C
Another way of indicating bonds
Often use a line to indicate a bond Called a structural formula Each line is 2 valence electrons
H HO =H HO
Structural Examples
H C N
C OH
H
C has 8 electrons because each line is 2 electrons
Ditto for N
Ditto for C here Ditto for O
Coordinate Covalent Bond When one atom donates both electrons
in a covalent bond. Carbon monoxide CO
OC
Coordinate Covalent Bond When one atom donates both electrons
in a covalent bond. Carbon monoxide CO
OC
Coordinate Covalent Bond When one atom donates both electrons
in a covalent bond. Carbon monoxide CO
OC
How do we know if Have to draw the diagram and see what
happens. Often happens with polyatomic ions and
acids.
Resonance When more than one dot diagram with
the same connections are possible.
NO2-
Which one is it? Does it go back and forth. It is a mixture of both, like a mule.
NO3-
VSEPR Valence Shell Electron Pair Repulsion. Predicts three dimensional geometry of
molecules. Name tells you the theory. Valence shell - outside electrons. Electron Pair repulsion - electron pairs
try to get as far away as possible. Can determine the angles of bonds.
VSEPR Based on the number of pairs of
valence electrons both bonded and unbonded.
Unbonded pair are called lone pair.
CH4 - draw the structural formula
Has 4 + 4(1) = 8 wants 8 + 4(2) = 16 (16-8)/2 = 4 bonds
VSEPR
Single bonds fill all atoms.
There are 4 pairs of electrons pushing away.
The furthest they can get away is 109.5º.
C HH
H
H
4 atoms bonded Basic shape is
tetrahedral. A pyramid with a
triangular base. Same shape for
everything with 4 pairs. CH H
H
H
109.5º
3 bonded - 1 lone pair
N HH
H
NH HH
<109.5º
Still basic tetrahedral but you can’t see the electron pair.
Shape is calledtrigonal pyramidal.
2 bonded - 2 lone pair
OH
H
O HH
<109.5º
Still basic tetrahedral but you can’t see the 2 lone pair.
Shape is calledbent.
3 atoms no lone pair
CH
HO
The farthest you can the electron pair apart is 120º
3 atoms no lone pair
CH
HO
The farthest you can the electron pair apart is 120º.
Shape is flat and called trigonal planar.
C
H
H O
120º
2 atoms no lone pair With three atoms the farthest they can
get apart is 180º. Shape called linear.
C OO180º
Hybrid Orbitals
Combines bonding with geometry
Hybridization The mixing of several atomic orbitals to form the
same number of hybrid orbitals. All the hybrid orbitals that form are the same
(degenerate = equal energy). sp3 - one s and three p orbitals mix to form four
sp3 orbitals. sp2 - one s and two p orbitals mix to form three sp2
orbitals leaving one p orbital. sp - one s and one p orbitals mix to form four sp
orbitals leaving two p orbitals.
Hybridization We blend the s and p-orbitals of the
valence electrons and end up with the tetrahedral geometry. We combine one s orbital and three p-orbitals.
sp3 hybridization has tetrahedral geometry.
sp3 geometry
109.5º
This leads to tetrahedral shape.
Every molecule with a total of 4 atoms
and lone pair is sp3 hybridized.
Gives us trigonal pyramidal and bent shapes also.
How we get to hybridization
We know the geometry from experiment. We know the orbitals of the atom hybridizing
atomic orbitals can explain the geometry. So if the geometry requires a tetrahedral
shape, it is sp3 hybridized. –This includes bent and trigonal
pyramidal molecules because one of the sp3 lobes holds the lone pair.
sp2 hybridization
C2H4
double bond acts as one pair trigonal planar Have to end up with three blended orbitals
–use one s and two p orbitals to make three sp2 orbitals.
– leaves one p orbital perpendicular
Where is the P orbital? Perpendicular The overlap of
orbitals makes a sigma bond
(s bond)
Two types of Bonds Sigma bonds from overlap of orbitals
between the atoms Pi bond (p bond) above and below atoms Between adjacent p orbitals. The two bonds of
a double bond
CCH
H
H
H
H
H
HB
sp2 hybridization
when three things come off atom trigonal planar 120º one p bond
B
B
A
B
trigonal planar
hybridize
s orbital
p orbitals three sps hybrid orbitals
What about two when two things come off one s orbital and one p orbital hybridize linear
sp hybridization end up with two lobes 180º
apart. p orbitals are at right
angles makes room for two p
bonds and two sigma bonds.
a triple bond or two double bonds
CO2
C can make two s and two p O can make one s and one p
CO O
N2
N2
Polar Bonds When the atoms in a bond are the
same, the electrons are shared equally. This is a nonpolar covalent bond. When two different atoms are
connected, the atoms may not be shared equally.
This is a polar covalent bond. How do we measure how strong the
atoms pull on electrons?
Electronegativity
A measure of how strongly the atoms attract electrons in a bond.
The bigger the electronegativity difference the more polar the bond.
0.0 - 0.5 Covalent nonpolar
0.5 - 1.0 Covalent moderately polar
1.0 -2.0 Covalent polar
>2.0 Ionic
How to show a bond is polar Isn’t a whole charge just a partial charge + d means a partially positive - d means a partially negative
The Cl pulls harder on the electrons The electrons spend more time near the Cl
H Cl+d -d
Polar Molecules
Molecules with ‘ends’
Polar Molecules Molecules with a positive and a
negative end Requires two things to be true
The molecule must contain polar bonds This can be determined from
differences in electronegativity.
Symmetry can not cancel out the effects of the polar bonds. Must determine geometry first.
Is it polar?
......
H Cl
Polar
H HO
PolarF F
B
F
Nonpolar
HH
H
N
Polar
Polar Nonpolar
F FCl
F
F F
Xe
F F
Nonpolar Polar
Cl
ClC
Cl
Cl
HC
Cl
H
H
XeF4
CCl4 CH3Cl
HCl H2O BF3 NH3
Bond Dissociation Energy The energy required to break a bond C - H + 393 kJ C + H We get the Bond dissociation energy
back when the atoms are put back together
If we add up the BDE of the reactants and subtract the BDE of the products we can determine the energy of the reaction (DH)
Find the energy change for the reaction
CH4 + 2O2 CO2 + 2H2O
For the reactants we need to break 4 C-H bonds at 393 kJ/mol and 2 O=O bonds at 495 kJ/mol= 2562 kJ/mol
For the products we form 2 C=O at 736 kJ/mol and 4 O-H bonds at 464 kJ/mol
= 3328 kJ/mol reactants - products = 2562-3328 = -766kJ
Intermolecular Forces
What holds molecules
to each other?
Intermolecular Forces They are what make solid and liquid
molecular compounds possible. The weakest are called van derWaal’s
forces - there are two kinds Dispersion forces Dipole Interactions
–depend on the number of electrons –more electrons stronger forces–bigger molecules
Depend on the number of electrons More electrons stronger forces Bigger molecules more electrons
fluorine (F2) is a gas
bromine (Br2) is a liquid
iodine (I2) is a solid
Dipole interactions
Dipole interactions
Occur when polar molecules are attracted to each other.
Slightly stronger than dispersion forces. Opposites attract but not completely
hooked like in ionic solids.
Dipole interactions Occur when polar molecules are
attracted to each other. Slightly stronger than dispersion forces. Opposites attract but not completely
hooked like in ionic solids.
H Fd+ d-
H Fd+ d-
Dipole Interactionsd
+ d
-
d+ d-
d+ d -
d+ d-
d+ d -
d+
d-
d + d
-d+
d-
Hydrogen bonding Are the attractive force caused by
hydrogen bonded to F, O, or N. F, O, and N are very electronegative so
it is a very strong dipole. The hydrogen partially share with the
lone pair in the molecule next to it. The strongest of the intermolecular
forces.
Hydrogen Bonding
HH
Od+
d-
d+
H HOd
+
d-
d+
Hydrogen bonding
HH
O H HO
HH
O
H
H
OH
HO
H
HO HH
O
Resources - Bonding
Objectives
Episode 8 – Chemical Bonds
Episode 9 – Molecular Architecture