1. Chapter 10: Structure of Solids and Liquids
Chem 1110
Figures: Basic Chemistry 3rd Ed., Timberlake and Timberlake

2. Electron-Dot Symbols Electron-dot symbols:
Show the valence electrons of an atom
Electrons are arranged in s and p orbitals around the element symbol
Can mix orbitals and create “hybrid” orbitals with equal energy (like in Chapter 6)
sp, sp2, and sp3

3. Covalent Bond Revisited
Covalent Bond is a chemical bond that results from two nuclei attracting the same shared electron pair
Shared electrons in the bond
For Example: F2

4. Bonding and Lone Pair Electrons
Bonding Pair are shared to create the covalent bond Non-Bonding Pair or lone pair make up the octet

5. Bonding and Lone Pair Electrons
Bonding Pair are shared to create the covalent bond
Non-Bonding Pair or lone pair make up the octet H2
HCl
H2O

6. Rules for Drawing Electron Dot Formulas
Find the total number of valence electrons contributed by all atoms
Write chemical symbols in the order that they are bonded
Place one covalent bond between each atom
Determining the central atom is important
First atom is usually the central atom
EXCEPT when it is hydrogen (H)

7. Rules for Drawing Electron Dot Formulas
Add lone pairs to the bonded atoms first, then to central atom
Check the Octet Rule for ALL atoms
May need to make multiple bonds
For Example: SF2

8. Rules for Drawing Electron Dot Formulas
Find the total number of valence electrons contributed by all atoms
Write chemical symbols in the order that they are bonded
Place one covalent bond between each atom
Add lone pairs to the bonded atoms first, then to central atom
Check the Octet Rule for ALL atoms
May need to make multiple bonds

9. Electron Dot Formulas
EXAMPLES:
CO2
NH3
SiCl4

10. Exceptions to the Octet Rule
Boron is stable with a total of 6 electrons
Boron will make compounds with only three bonds to B
BCl3
Period 3 (and above) non-metals can form compounds with expanded octets
(> 8 valence electrons): P, S, Se

11. Multiple Bonds
Multiple bonds involve the sharing of more than one pair of electrons:

12. Bond Order Bond Order denotes the level of bonding in a molecule:
Single bond: Share one electron pair
Bond Order of 1
Double bond: Share two electron pairs
Bond Order of 2
Triple bond: Share three electron pairs
Bond Order of 3

13. Bond Length and Bond Energy
As we add additional bonds between atoms the atoms are drawn closer together:
Bond length decreases with more bonds:
Single > Double > Triple
As we add additional bonds between atoms the bonds get stronger:
Bond energy increases with more bonds:
Triple > Double > Single

14. Bond Length and Bond Energy
Practicing organic compounds:
Single bond (alkane): Ethane, C2H6
Double bond (alkene): Ethene, C2H4
Triple bond (alkyne): Ethyne (Acetylene), C2H2

15. Electron Dot Formulas
EXAMPLES: O2
HCN
CS2 CO

16. Resonance Structures SO2 SO3
Resonance structures are equivalent structures that show where multiple bonds may occur
Electrons are delocalized over entire molecule
Resonance changes bond order, bond length, and energy
Examples:
SO2
SO3

17. Learning Check
Draw the three resonance structures of carbonate, CO32-

18. Electron Dot Formulas
How can we draw electron dot structures for polyatomic ions?
NH4Br or Na2SO3
These compounds have both an ionic and covalent:
[NH4]+ [Br]-
2 [Na]+ [SO3]2-

19. Electron Dot Formulas
[NH4]+
[SO3]2-

20. Learning Check
Draw the electron dot formula for ClO3-

21. Learning Check
Write two resonance structures for nitrite.

22. Orbital Geometry and Molecular Shape
Electrons in bonds and lone pairs want to get as far apart from each other as possible
We have seen the tetrahedral shape of carbon
Causes the molecule to conform to a set of shapes

23. Determining Orbital Geometry
Electron pairs do NOT like to share space with other electron pairs
Electrons will move away as far as possible to obtain “personal space”
An electron pair can be in a bond
An electron pair can be a “lone pair” (non-bonding)

24. VSEPR Theory: Valence-Shell Electron Pair Repulsion Theory
Areas of electron density around an atom want to maximize their “personal space”
A lone pair of electrons is ONE area
Two electrons in a single bond are ONE area
Four electrons in a double bond are ONE area
Six electrons in a triple bond are ONE area

25. VSEPR Theory Determining Orbital Geometry:
Count ALL areas of electron density to determine the base geometry:
Two is linear (sp)
Three is trigonal planar (sp2)
Four is tetrahedral (sp3)

26. Orbital Geometry Linear (sp)

27. Orbital Geometry Trigonal Planar (sp2)

28. Orbital Geometry Tetrahedral (sp3)

29. Orbital Geometry

30. Linear and Trigonal Planar Molecules

31. Tetrahedral Molecules

32. Molecular Shape is Altered by Lone Pairs
Two different “bent” structures which differ in their bond angle: 120° vs °
Lone pairs take up more room and squeeze bond angles together:
Trigonal planar: 120°
Trigonal planar with lone pair: < 120°
Tetrahedral: 109.5°
Tetrahedral with lone pair: < 109.5°

33. Orbital Geometry and Molecular Shape
REMEMBER:
Electrons repel each other
Electrons in bonds have a fixed location
Electrons in lone pair take up more space
Two atoms always define a line.

34. Molecular Shape Let’s look at two seemingly similar molecules:
CO2 and H2O
We first have to draw the molecule to determine their shapes:

35. Molecular Shape We can base molecular shape on orbital geometry:
Linear: CO2
Trigonal Planar: BF3, CH2O
Tetrahedral: CH4 , NH3, H2O

36. Learning Check
Determine the shape of the N2O molecule:

37. Electron Dot Formulas
Practice drawing electron dot structures and identify the proper molecular shape for:
AlH3
NF3
SiCl4
HCN CO

38. Electronegativity WHY????
In the preceding structures we had sharing of electrons (covalent bonds) and charged species (transfer of electrons, ionic bonds):
How do we know what bond type we have?
General Rule of Thumb:
A metal with a non-metal forms an ionic bond
A non-metal binding with a non-metal forms a covalent bond
WHY????

39. Electronegativity
The type of bond interaction is determined by differences in electronegativity:
Electronegativity is the ability of an atom to attract shared electrons in a bond towards itself
Time to revisit Periodic Trends…

40. Electronegativity Fluorine is the MOST electronegative
Increasing Electronegativity
Increasing Electronegativity
Fluorine is the MOST electronegative
Cs/Fr are the LEAST electronegative

41. Electronegativity Values
We can assign values for electronegativity:

42. Learning Check
Place the following in order of increasing electronegativity: O, K, and C

43. Determining Bond Type
Difference in Electronegativity (ΔEN) allows us to determine the type of bond formed:
Difference > → ionic bond
Difference = → pure covalent bond
Difference between 0.0 and 0.4
Non-polar covalent
Difference between 0.4 and 1.8
Polar covalent bond

44. Common Pure Covalent Bond Species
O2 Cl2 Br2 I2 H2 N2 S8 P4

45. Bond Polarity
Bonds which have an unequal sharing of electrons are said to be Polar:
The electron cloud is pulled toward the more electronegative atom
This sets up a separation of charges: dipole
The more electronegative element attracts the cloud more strongly - partially negative
The less electronegative element attracts the cloud less - partially positive

46. Bond Polarity

47. Bond Polarity Blue: low electron density Red: high electron density
Green: non-polar

48. A different way of looking at polarity:
Bond Polarity
A different way of looking at polarity:

49. EXAMPLES Molecule Δ EN and Bond Type NH3 O2 NaCl SO2
0.9 → polar covalent 0 → pure covalent 2.3 → ionic 1.0 → polar covalent

50. Electronegativity and Bond Types

51. Predicting Bond Types

52. Learning Check 1) K and N 2) N and O 3) Cl and Cl 4) H and Cl
Use electronegativity differences to classify each of the following bonds as nonpolar covalent (NP), polar covalent (P), or ionic (I):
A bond between:
1) K and N
2) N and O
3) Cl and Cl
4) H and Cl

53. Drawing Bond Dipoles
K – N
N – O
Cl – Cl

54. Molecular Shape and Polarity
Dipoles can add across a molecule
Polar molecules – have an overall dipole
Non-polar molecules – no overall dipole
To determine whether a molecule is polar or non-polar, we must determine the 3-D shape of the molecule

55. Molecular Polarity Additive dipoles across a molecule make it polar:
If dipoles cancel, no net dipole and the molecule is non-polar:

56. Molecular Polarity
EXAMPLES: CO2 NH3 CF4

57. Learning Check Identify each of the following molecules as
(P) polar or (NP) nonpolar:
PBr3
HBr
Br2
D. SiBr4

58. Learning Check
Draw the Lewis Structure for SeCl2 and determine its shape and molecular polarity:

59. Forces in Matter
Intermolecular Forces: are attractive forces between molecules:
Weaker than bonds (covalent or ionic)
Help to determine the state of matter:
solid, liquid or gas
Intermolecular forces organize matter and are opposed by motion of molecules which disorganize matter

60. States of Matter Solid Attractive Forces >> Disruptive Forces
Liquid
Attractive Forces ≈ Disruptive Forces
Gas
Attractive Forces << Disruptive Forces

61. Intermolecular Forces
Three primary types of forces that help to hold covalent molecules together:
London Dispersion Forces
Dipole – Dipole Interactions
Hydrogen Bonding

62. Intermolecular Forces
London Dispersion Forces occur in non-polar molecules by forming a temporary dipole:
Temporary dipoles induce dipoles in nearby molecules
Lasts for milliseconds!
ALL matter has London Dispersion Forces
Let’s consider H2 (non-polar molecule, H–H)

64. Intermolecular Forces
London Dispersion Forces
Increase with increasing size
Increase with increased number of electrons
Examples: CH4, CF4, CCl4, CBr4
CBr4 will have the strongest London dispersion forces as it is the largest
CH4 will have the weakest London dispersion forces

65. Intermolecular Forces
London Dispersion Forces
Increase with increasing size
Increase with increased number of electrons
Electrons on larger atoms are further from the nucleus therefore easier to induce temporary dipole
Examples: CH4, CF4, CCl4, CBr4
CBr4 will have the strongest London dispersion forces as electrons on Br are further from nucleus

66. Intermolecular Forces
Dipole-Dipole Interactions occur between polar molecules where δ+ end attracts the δ- end of another molecule
There is a permanent dipole in polar molecules: attraction of molecules
Occur only between polar molecules
Consider FCl:
Fluorine is more electroneagtive (δ-)

68. Intermolecular Forces
Dipole-Dipole Interactions
Stronger than London dispersion forces
Strength of interaction increases with increasing size of the permanent dipole
For Example:
FCl (ΔEN= 1.0), FBr (ΔEN= 1.2), FI (ΔEN= 1.5)
FI will have the strongest dipole–dipole interaction
FCl will have the weakest dipole–dipole forces

69. Intermolecular Forces
Hydrogen bonds are a special type of dipole-dipole force when we have a VERY large difference in electronegativity
To have a Hydrogen Bond:
Need to have H
Need to have F, O, or N
H must be bonded to F, O, or N
At least one lone pair on the F, O, or N

70. Intermolecular Forces
Hydrogen bond: The hydrogen on one molecule is attracted to the lone pair on another molecule
The Hydrogen bond is the STRONGEST intermolecular force!
Let’s consider H2O!

71. Hydrogen Bonds in Water

72. Hydrogen Bonds

73. Hydrogen Bonds Predicting strength of H-bonds: CH3OCH3 vs. C2H5OH
CH3NH2 vs. CH3NHCH3 vs. (CH3)3N

74. Learning Check
Identify the major type of attractive force in each of the following substances:
A. NCl3
B. H2O
C. Br2
D. KCl
E. NH3

75. Equilibrium
When the forces acting on matter are in balance, we have Equilibrium
Or, Equilibrium occurs when two or more opposing forces balance each other

76. Equilibrium: Balance of Forces
Equilibrium occurs when two or more opposing forces balance each other, and is a dynamic process:
Represented by the symbol: ⇄
For Example (Changes of State):
Solid ⇄ Liquid Liquid ⇄ Gas
Melting point Boiling point

77. Changes of State

78. Intermolecular Forces
Strength of intermolecular forces increase as previously described:
London < Dipole-Dipole < Hydrogen Bond
As the strength of intermolecular forces increases, the properties of molecules change:
Want to “stick together” more!
Polar molecules tend to be liquids or solids at room temperature
Requires increased kinetic energy (disruptive forces) to separate them!

79. Intermolecular Forces
Strength of intermolecular forces increase as previously described:
London < Dipole-Dipole < Hydrogen Bond
As the strength of intermolecular forces increases, they change the properties of molecules:
Melting Point Increases
Boiling Point Increases
Vapor Pressure Decreases

80. Learning Check
Identify the compound in each pair that has the higher melting point. Explain your choice:
NCl3 or NH3
HBr or Br2
C. KCl or HCl

81. Balance of Forces Let’s consider H2O:
Hydrogen bonds are the strongest force
Requires lots of energy to break apart intermolecular interactions
Consequently, water has an increased melting point, decreased vapor pressure, and increased boiling point

82. Liquid Water