1. Chapter 10 Molecular Structure

2. Chapter 10 Molecular Structure
10.1 Covalent Bonds
10.2 Hybridization of Atomic Orbitals
10.3 Valence Shell Electron-Pair Repulsion Theory (VSEPR)
10.5 Intermolecular Forces

3. Review of Chemical Bonds
What is a chemical bond?
the attractive force that holds atoms together in multi-atom elements or in compounds
Why does it come about?
Octet Rule: atoms combine to form bonds either by transferring electrons to form ionic bonds, or by sharing electrons in covalent bonds, until each atom is surrounded by eight valence electrons

4. Problem
How can an atom alter its electron configuration to obtain an octet (or duet) of electrons of a noble gas?
Solution
A metal a cation
A nonmetal an anion
ionic compounds
Sharing of electrons between nonmetals
molecular compounds

5. Types of Chemical Bonds:
Ionic Bonds: transferring electrons
Covalent Bonds: sharing electrons
Metallic Bonds: Valence electrons are detached from atoms, and spread in an 'electron sea' that "glues" the ions together

6. Ionic Bond
is a strong electrostatic attraction between a positive ion and a negative ion
electron is fully transferred from metal to nonmetal
non-directional, magnitude of bond equal is all directions
typically occur between a metal and a reactive non-metal

7. Covalent Bond Cooperative sharing of valence electrons
Covalent bonds are HIGHLY directional
typically occur between non-metal
have a relatively low melting and boiling point

8. 10.1 Covalent Bonds Key points:
Why and how do two atoms bond together ?
What’s the difference between them?
σbond andπbond
single bond, double bond and triple bond
polar and nonpolar covalent bond

9. 1. Describing Covalent Bonds
Definition: A shared pair of electrons between two atoms is known as a covalent bond.
Example: H2
no overlap; no attraction
74pm
Maximum attraction
repulsion
potential energy →
Bond dissociation energy
Bond length
Distance between nuclei →

10. Covalent Bond Theory :
Bonds are formed by atom sharing two electrons in overlapping atomic orbitals.
the number of bonds formed = the number of unpaired electrons usually NH3
Orbitals bond in the same axis to obtain maximum overlap. e.g.:HCl

11. Cl 1s22s22p6 3s23p5
H 1s1 x x z

12. 2. Bond Properties
Bond Energy: the energy required to break one mole of bonds in a gaseous species.
NH3(g)=NH2(g)+H(g) D1=435.3kJ·mol-1
NH2(g)=NH(g)+H(g) D2=397.5kJ·mol-1
NH(g)=N(g)+H(g) D3=338.9kJ·mol-1
NH3(g)=N (g)+3H(g) D=1171.5kJ·mol-1
EN-H=(D1+D2+D3)/3=390.5 kJ·mol-1→
The larger the bond energy, the stronger the chemical bond.

13. Bond Length: distance between nuclears of two bonded atoms
Bond lengths from x-ray analysis:
Bond Angle
143
122
113
104045‘

14. 3. Types of Bonds Ionic Bond: electrostatic attraction
Covalent Bond: sharing of electrons
Chemical Bond
Coordinate Covalent Bond: both electrons from one atom (CO)
Covalent Bond:
one electron from each atom
C 1s22s22p2
O 1s22s22p4
Coordinate Covalent Bond
Single Bond
Double Bond
Triple Bond
Bond order
σbond
πbond
orbital overlap
Polar Covalent Bond
Nonpolar Covalent Bond
Bond Polarity

15. Sigma () bonds sigma bond and Pi bond
Sigma (σ) bond = end-to-end overlap
☆ head-on overlap = end-to-end overlap
☆ s-s , s-px , px-px *
s + s sigma overlap
s + p sigma overlap
p + p sigma overlap
Sigma () bonds

16. Pi () bond Pi (π) bond = side-by-side overlap ☆ sideways overlap
☆ pz-pz , py-py *
Pi () bond +
p orbital
p orbital
 bond

17. σ键 π键

18. Note here, that the extent of overlap between the orbitals is less in a pie bond as compared to a sigma bond. Thus, the force of attraction between the shared electron pair and the nuclei is comparatively less than that in a sigma bond. Therefore, a pie bond is always weaker than a sigma bond,and it can only coexist with sigma bond.
A single covalent bond is a sigma bond.
A double covalent bond is made up of one sigma and one pie bond.
A triple covalent bond is made up of one sigma and two pie bonds.

19. C—C C C C C Bond Order and Multiple Bonds
☆ single bond => 1 pairs shared
☆ double bond => 2 pairs shared
☆ triple bond => 3 pairs shared N2
Relationship Between Order, Length, and Energy
Bond order = length (pm) = energy (kJ/mol) =
Conclusion: Bond lengths SHORTEN as bond order INCREASES. A double bond is STRONGER than a single bond.
C—C C C C C

20. 10.1 Covalent Bonds Polar Covalent Bonds and Nonpolar Covalent Bonds
Problem
The electrons in the covalent bond are shared equally?
Solution
The electronegativity (EN) differences of the elements are used to determine the extent of this unsharing.

21. Nonpolar Bond
Polar Bond
Ionic Bond
diff. EN = 0
diff. EN > 0
diff. EN > 1.7
Electrons shared equally.(atoms with same electronegativity)
Electrons shared unequally.(atoms with different electronegativity)
Electron transferred.
H2:
HCl:
NaCl:
Notice: bond polarity≠molecular polarity

22. 10.2 Hybridization of Atomic Orbitals
Questions:
Why can they hybrid?
What kinds of orbitals are used for hybridization?
Why do they need to hybrid?
How to predict the shapes of molecules?

23. Problem
Experiment shows that the methane molecule (CH4) is tetrahedral, with four equivalent bonds.
Solution
1s22s12p3

24. mutually perpendicular
spherically symmetrical

25. “atomic orbitals” vs. “hybrid orbitals”
☆ on the same atom (central atom only )
☆ in the same principal energy level or, occasionally, in adjacent energy level
hybrid orbitals
☆ the shapes and directional properties of new hybrid orbitals ≠the orbitals used in constructing them
☆ the number of hybrid orbitals formed = number of atomic orbitals used

26. 10.2 Hybridization of Atomic Orbitals
After the paint is mixed, there are still the same numbers of cups of paint, but the color has changed. The mixing of orbitals is analogous. sp3

27. 10.2 Hybridization of Atomic Orbitals
The directional characteristics of the new hybrid orbitals will be decided by the type and number of atomic orbitals using in the mixing.
Table 3 Three Types of s-p Hybrid
Atomic orbitals
Hybrid Orbitals
Geometric Arrangement
one s, one p
two sp
Linear*
one s, two p
three sp2
Trigonal planar*
one s, three p
four sp3
Tetrahedral

28. 10.2 Hybridization of Atomic Orbitalssp

29. 10.2 Hybridization of Atomic Orbitals
sp2

30. 10.2 Hybridization of Atomic Orbitals
sp3

31. 10.2 Hybridization of Atomic Orbitals
Needed to form 4 sigma bonds
The mutual repulsion among the electron pairs will orient them toward the apices of a tetrahedron.

32. 1s22s12p3

33. 1. An Outline of Hybrid Orbital Theory
Why can they hybrid?
We can mathematically mix two or more of these wave functions that describe the electron, and produce an equal number of wave functions that have different shapes and orientation. However, the number of new hybrid orbitals will be the same as the number of orbitals used to form them.

34. What kind of orbitals are used for hybridization?
Hybrids are formed among orbitals lying in the same principal energy level or, occasionally, in adjacent energy levels.
Example ns np
( n-1) d ns np or ns np nd

35. 10.2 Hybridization of Atomic Orbitals
Why do they need to hybrid?
Each sp3 hybrid orbital has a large lobe pointing in one direction and a small lobe pointing in the opposite direction. The enlarged lobe of the resulting orbital can give more favorable overlap with the orbital of another atom and thus form a stronger bond than can either the p or s orbitals alone.
Mix ＋ _ ＋ _ ＋ _
Mix

36. 10.2 Hybridization of Atomic Orbitals
Why do they form a geometry ?
The mutual repulsion among the electron pairs will orient hybrid orbitals toward the apices of a geometry.
Hybrid Orbitals
Number of Orbitals
Geometric Arrangement sp 2
Linear
sp2 3
Trigonal
sp3 4
Tetrahedral
sp3d 5
Trigonal bipyramidal
sp3d2 6
Octahedral

37. 10.2 Hybridization of Atomic Orbitals
2. Application- Beryllium Chloride, BeCl2 : sp
Use sp hybrid orbitals
Two electron pairs produce a linear arrangement. Cl—Be—Cl

38. 10.2 Hybridization of Atomic Orbitals
BF3 : sp2
Use sp2 hybrid orbitals
Three electron pairs produce a triangular planar arrangement .

39. 10.2 Hybridization of Atomic Orbitals
sp3:CH4 NH3 H2O
CH4
NH3
H2O
equivalent hybridization
nonequivalent hybridization
104045‘

40. C2H6
C2H4
What would be the name of the hybrid orbitals created by joining 1 s type, 3 p type and 2 d type orbitals?
a. spd b. sp3d2 c. s3p2d
C2H2

41. 10.3 Valence Shell Electron-Pair Repulsion Theory
(VSEPR)
Key Points:
What’s the difference between the arrangement of electron pairs and the molecular geometry ?
How to predict the shapes of molecules?
PbCl2 Pb: 5d106s26p2 Pb2+→ sp？linear？
HgCl2 Hg: 5d106s2 Hg2+→ sp？linear？

42. 10.3 Valence Shell Electron-Pair Repulsion Theory
1. VSEPR model
The arrangement of electron pairs describes the arrangement of all electron pairs, shared or unshared, around a central atom.
The molecular geometry of a molecule is the geometry described by the bonded atoms and does not include the unshared pairs of electrons.
The minimum repulsion occurs when the electron pairs are as far apart as possible.

43. 10.3 Valence Shell Electron-Pair Repulsion Theory
Table 1 Arrangement of electron pairs about an atom
Number of electron Pairs
Arrangement of Electron Pairs 2
Linear 3
Trigonal planar 4
Tetrahedral 5
Trigonal
bipyramidal 6
Octahedral

44. 10.3 Valence Shell Electron-Pair Repulsion Theory
If more than one structure is possible, the most stable structure can often be determined using the following rules:
Bond angles: 90º>120º>180º
Lone pairs (LP): LP-LP>LP-BP>BP-BP
Bond order：triple >double >single bond
Electronagetivity of the donor atom：Xsmall > Xlarge
Select the stable structure which has the smallest number of lone pair-lone pair repulsions (usually 90o).

45. 10.3 Valence Shell Electron-Pair Repulsion Theory
2. How to predict molecular structures?
Determine the number of electron pairs (both bonded pairs and lone pairs) around the central atom.
the number of electron-pairs=1/2 (the number of valence electrons on the central atom + the number of electrons furnished by donor atom)
the number of valence electrons on the central atom ＝group number
hydrogen and the halogens (Group VIIA) donate one electron each for sharing，

46. 10.3 Valence Shell Electron-Pair Repulsion Theory
Elements in the oxygen group (Group VIA) are considered to donate no electrons.
Positive ion: subtract the positive charge on the ion.
Negative ion: add the negative charge on the ion.
In the case of an odd number of electrons, treat the extra electron (one-half of an electron-pair) as if it were an electron-pair (lone pair).
Count a multiple bond as one pair.

47. 10.3 Valence Shell Electron-Pair Repulsion Theory
Identify the molecular geometry
No. of e- pairs
Arrangement of pairs
Molecular geometry of ABn
2（sp）
Linear
AB2
3（sp2）
Trigonal planar
AB AB2
4（sp3）
Tetrahedral
AB AB AB2 5
（sp3d ）
Trigonal
bipyramidal
AB AB AB AB2 6
(sp3d2)
Octahedral
AB AB AB4

48. 10.3 Valence Shell Electron-Pair Repulsion Theory
3. Applying VSEPR Theory
Using the VSEPR theory we could predict the geometry of an AXn molecule or ion .
Using the VSEPR theory we could predict bond angles and molecular geometry.
Using the VSEPR theory we could predict whether the molecule is polar or nonpolar.
Using the VSEPR theory we could describe the hybrid orbitals used by the central atom.

49. 10.3 Valence Shell Electron-Pair Repulsion Theory
Central atom: Cl 3s23p5
LP–LP
LP–BP
BP–BP
Using the VSEPR theory we could predict the geometry of an AXn molecule or ion .
Predict the geometry of the following molecules or ions.
BF3 PCl5 H2O H3O+ NH3 NH4+ NO2 NO2- SF ClF3 XeF2 IF XeF4 SO O3

50. 10.3 Valence Shell Electron-Pair Repulsion Theory
Procedure
☆Determine the number of electron pairs.
※a multiple ≌ one pair ≌ an unpaired electron
☆ Arrange the electron pairs.
☆ Identify the molecular geometry.
with lone pairs: the arrangement of electron pairs ≠ the molecular geometry

51. 10.3 Valence Shell Electron-Pair Repulsion Theory
Using the VSEPR theory we could predict bond angles and molecular geometry.
CH4 : NH3 : H2O:
No lone pair Tetrahedral
One lone pair Trigonal pyramidal
Two lone pairs Bent
Notice: When lone pairs are present, the
bond angles are smaller than the perfect angles.

52. 10.3 Valence Shell Electron-Pair Repulsion Theory
Using the VSEPR theory we could predict whether the molecule is polar or nonpolar.
Problem
How dose bond polarity relate to molecular polarity?
Solution
☆ For diatomic molecules, if the bond is polar then the molecule is polar.
☆ For all other molecules, we need to know the molecular shape to be able to predict the polarity.

53. 10.3 Valence Shell Electron-Pair Repulsion Theory
If two people of exactly equal strength pull on the box in exactly opposite directions, their efforts cancel and there is no movement.
Direction of movement

54. 10.3 Valence Shell Electron-Pair Repulsion Theory
nonpolar molecule
polar molecule
1. Polar or nonpolar covalent bond
Polar covalent bond
2. has a perfect geometry (no lone pairs )
does not have a perfect geometry (with lone pairs )
3. dipole moments (μ) = 0
μ ≠ 0
4. Example:
SO2 :
CO2 :
Linear
Bent
polar bonds nonpolar molecule
polar bonds polar molecule

55. 10.3 Valence Shell Electron-Pair Repulsion Theory
Conclusion
☆ For diatomic molecules, if the bond is polar then the molecule is polar.
☆ For all other molecules, we need to know the molecular shape to be able to predict the polarity.
※ Molecules with lone pairs of electrons on a central atom are generally polar molecules.
※ Nonpolar molecules that contain polar bonds are observed when there are no lone pairs on the central atom, and all of the atoms bonded to the central atom are identical.

56. 10.3 Valence Shell Electron-Pair Repulsion Theory
Use the VSEPR model to obtain the arrangement of electron pairs.
From the arrangement of the electron pairs, deduce the type of hybrid orbitals.
Arrangement of electron pairs
Hybrid Orbitals
Linear
Sp(2)
Trigonal planar
Sp2(3)
Tetrahedral
Sp3(4)
Trigonal bipyramidal
Sp3d(5)
Octahedral
Sp3d2(6)

57. 10.5 Types of Intermolecular Forces
Key Points:
What kinds of attractions exist between molecules?
What’s the difference between the intermolecular forces and the intramolecular forces ?

58. 10.5 Types of Intermolecular Forces
Intramolecular forces hold atoms together in molecules.
Ionic Covalent Metallic
Cation-anion Nuclei-shared e- pair Cations-delocalized e-
Intermolecular forces are those between molecules.
H bond
Dipole-dipole
Dipole-induced dipole
Dispersion (London)
Polar bond to H-dipole charge (high EN of N, O, F)
Dipole charges
Dipole charge-polarizable e- cloud
Polarizable e- clouds

59. 10.5 Types of Intermolecular Forces
2. Explaining Liquid Properties
Many of the physical properties of liquids (and certain solids) can be explained in terms of intermolecular forces.
play a key role in determining the conditions under which a substance is a solid, a liquid, or a gas
affect physical properties, such as the melting points of solids and the boiling points of liquids.

60. 10.5 Types of Intermolecular Forces
3. Dipole Moment
A polar molecule and dipole moment
A polar molecule
has permanent dipole
Polar molecules are aligned in an electric field

61. 10.5 Types of Intermolecular Forces
Dipole moment 
= qx
, dipole moment
q, charge (or fractional charges)
x, distance between charges
a quantitative measure of the degree of charge separation in a molecule.
expresses a molecule’s polarity. Molecules with large dipole moments are highly polar. Nonpolar molecules have a zero dipole moment.

62. 10.5 Types of Intermolecular Forces
Instantaneous And Induced Dipoles
In describing electronic structures we speak of electron charge density or the probability that an electron is in a certain region at a given time.
One probability is that at some particular instant―purely by chance―electrons are concentrated in one region of an atom or molecule. This displacement of electrons causes a normally nonpolar species to become momentarily polar.
After this, electrons in a neighboring atom or molecule may be displaced to also produce a dipole, called an induced dipole.

63. 10.5 Types of Intermolecular Forces
Dipole-Dipole Forces
Dipole- Induced Dipole Forces
London Forces
Hydrogen Bonding
Van Der Waals Forces

64. 10.5 Types of Intermolecular Forces
Dipole-Dipole Forces
A polar molecule has two "poles". (Permanent Dipole )
Molecules with a permanent dipole can align themselves so that the negative end of one molecule is attracted to the positive end of another. These attractions are known as dipole-dipole attractions.

65. 10.5 Types of Intermolecular Forces
Dipole-Dipole Forces
Notice
only in permanence
somewhat stronger than London forces because permanent
occurs in all polar molecules

66. 10.5 Types of Intermolecular Forces
Dipole-Induced Dipole Interactions
Dipole H2O and nonpolar O2:
The dipole of H2O induces a dipole in O2 by be distorting the O2 electron cloud – Polarizability
Interactions of polarizable molecules are called induced dipole interactions
Notice occurs in polar and nonpolar molecules

67. 10.5 Types of Intermolecular Forces
Induced dipole - Induced dipole Interactions (London forces )
Two nonpolar atoms or molecules has no dipole moment .
Momentary attractions and repulsions between nuclei and electrons in neighboring molecules lead to induced dipoles.
The attraction between these induced dipoles then spreads from molecule to molecule, thus providing an attractive force between molecules.

68. 10.5 Types of Intermolecular Forces
Even nonpolar molecules and uncombined atoms have attractive forces between them.
Notice
all particles have London forces
London forces increase with molecular weight.
only force between noble gases and nonpolar compounds.
about 1/1000 as strong as a covalent bond

69. Dipole- Induced Dipole Forces
Notice
All particles have London forces. London forces increase with molecular weight.
about 1/1000 as strong as a covalent bond
Dipole-Dipole Forces
Dipole- Induced Dipole Forces
London Forces
polar-polar √
polar-nonpolar
nonpolar-nonpolar

70. 10.5 Types of Intermolecular Forces
Hydrogen Bonding
Problem Examination of the boiling points of the groups VA, VIA, and VIIA reveals that the first compound in each of these series has an unexpectedly high boiling point.

71. A hydrogen atom bonded to an electronegative atom appears to be special. The electrons in the O-H bond are drawn to the O atom, leaving the dense positive charge of the hydrogen nucleus exposed. It’s the strong attraction of this exposed nucleus for the lone pair on an adjacent molecule that accounts for the strong attraction. A similar mechanism explains the attractions in HF and NH3.

72. Condition
X—H…Y(X,Y=F、O、N)
X→ Xlarge、rsmall
Y→ Xlarge、rsmall ，lone pair electron

73. 10.5 Types of Intermolecular Forces
Notice
Rare, strong form of dipole-dipole interaction.
Require H bonded to highly electronegative atom (N, O or Halogen): one of the following three structures must be present. H-N O-H F-H
One molecule has a hydrogen atom attached by a covalent bond to an atom of nitrogen, oxygen, or fluorine. The other molecule has a nitrogen, oxygen, or fluorine atom present that possesses one or more lone pairs.
strongest of four intermolecular forces
are HIGHLY directional

74. 10.5 Types of Intermolecular Forces
Types of Hydrogen Bonds
intermolecular hydrogen bonds
intramolecular hydrogen bonds
Increase boiling points
Decrease boiling points

75. 10.5 Types of Intermolecular Forces
5. The Double Helix of DNA is held together by hydrogen bonding
In the three-dimensional Watson-Crick structure, two polynucleotide DNA strands wind around each other to form a double helix. Hydrogen bonds form between specific base pairs. Adenine is hydrogen bonded to thymine, and guanine is hydrogen bonded to cytosine to form complementary base pairs (A on one strand with T on the other strand, or C with G).
adenine腺嘌呤 thymine胸腺嘧啶
guanine鸟嘌呤cytosine胞嘧啶

76. 10.5 Types of Intermolecular Forces
white = hydrogen blue = nitrogen black = carbon red = oxygen

77. 10.5 Types of Intermolecular Forces
Two hydrogen bonds occur between every adenine and thymine pair; three between each guanine and cytosine.