1. Ch. 4: Chemical Bonding and Molecular Geometry
Dr. Namphol Sinkaset
Chem 200: General Chemistry I

2. I. Chapter Outline Introduction Ionic Bonding Covalent Bonding
Chemical Nomenclature
Lewis Dot Symbols and Structures
Formal Charges and Resonance
Molecular Structure and Polarity

3. I. Introduction
When elements form compounds, the original properties of the elements are lost.

4. I. To Lower Potential Energy
A chemical bond is the force that holds atoms together in a compound.
But why would atoms want to join with other atoms?
It all comes back to positive-negative attractions between particles in the atom which lead to lower PE!

5. I. Two Main Ways to Lower PE

6. I. Metal + Nonmetal = Ionic

7. I. Nonmetal + Nonmetal = Covalent
Instead of transferring e-’s, covalent bonding occurs via sharing of e-’s.
Attraction to two nuclei lowers PE.

8. II. Ionic Bonding
In ionic bonding, metal transfers e- to the nonmetal.
Metals have low IE’s; nonmetals have high EA’s.
Resulting ions attracted by +/- charge.
Number of cations and anions must lead to a neutral compound.
Charge on one becomes subscript on other.

9. II. Not Just One Ionic Bond

10. II. Cation Electronic Structure
Cations are often isoelectronic with noble gases.
Lower groups have partial loss of e-’s due to inert pair effect – stability of paired s e-’s.
Transition metals lose higher n e-’s first!

11. II. Anion Electronic Structure
Most monatomic ions form when valence s and p orbitals are “filled.”
Gaining e-’s makes anions isoelectronic with the next noble gas.
e.g. Se2-: [Ar] 3d104s24p6 = Kr
Recall that charges for main-group monatomic ions can be deduced from location on periodic table.

12. II. Properties of Ionic Compounds
Ionic compounds have some distinguishing characteristics:
Have crystalline structure
Are rigid and brittle
Have high melting and boiling points
Conduct electricity when molten or when dissolved in water

13. II. Ionic Compounds Melt at High Temperatures
Why are such high temperatures needed?

14. II. Electrical Conductivity

15. III. Covalent Bonding
In covalent bonding, nonmetals share some (or all) of their valence electrons.
They share because the atoms involved have identical or similar IE’s and EA’s.
The formation of the bond can be explained energetically.

16. III. Formation of H2

17. III. Fully Equal Sharing
If atoms involved in a covalent bond are the same, then e-’s will be shared equally.
Called a pure covalent bond.
Examples include all the diatomics like H2, Cl2, Br2, etc.

18. III. Unequal Sharing
If atoms in a covalent bond are different, then e-’s aren’t shared equally.
The e-’s will be more attracted to one of the atoms resulting in a polar covalent bond.

19. III. Determining Bond Polarity
Is there a way to predict bond polarity?
Linus Pauling derived electronegativity values from looking at energies needed to break different types of bonds.

20. III. Electronegativity
To determine which atom in a polar covalent bond holds the partial negative charge, we use electronegativity.
Electronegativity is the relative ability of a bonded atom to attract shared e-.
It can be thought of as how greedy an atom is for e- when it is sharing them.
Note: electronegativity ≠ electron affinity.

21. III. EN Values

22. III. Using ΔEN
Differences in electronegativity can be used to determine the bond type.

23. III. Properties of Covalent Compounds
Covalent compounds have some distinguishing characteristics:
Have low melting and boiling points relative to ionic compounds
Tend to be insoluble in water
Are poor electrical conductors in any state

24. IV. Chemical Nomenclature
Naming inorganic compounds is systematic.
Key is identify what type of compound you are naming.
Variables include:
Ionic or covalent?
Single or multiple charge states?
Polyatomic ion(s) involved?
Acid?

25. IV. Ionic Nomenclature
Ionic compounds are named systematically, broken into two groups.

26. IV. Type I Compounds
Type I compounds are ionics that have a metal from Groups 1 or 2 and a nonmetal from Groups
Examples:
NaCl = sodium chloride
MgBr2 = magnesium bromide
K2S = potassium sulfide

27. IV. Type I Compounds
To get a formula from a name, remember that a compound must be neutral.
Ion charges can be found by locating the element on the periodic table.
“The charge on one becomes the subscript of the other.”

28. IV. Type I Compounds
e.g. What are the formulas for sodium nitride, calcium chloride, potassium sulfide, and magnesium oxide?

29. IV. Transition Metals
Transition metals are found in the “Valley,” Groups 3-12, of the periodic table.
Transition metal cations often can carry different charges, e.g. Fe2+ and Fe3+.
Thus, a name like “iron chloride” is ambiguous.

30. IV. Type II Compounds
Type II compounds are ionics that have a transition metal (Groups 3-12) and a nonmetal (Groups 14-17).
Examples:
FeCl2 = iron(II) chloride
FeCl3 = iron(III) chloride

31. IV. Type II Compounds
e.g. Give the correct name or formula for the compounds below.
MnO2
copper(II) chloride
AuCl3
molybdenum(VI) fluoride
Hg2Cl2

32. IV. Type II Compounds
An archaic naming system uses common names for transition metal cations of different charge.
Higher charge given –ic suffix
Lower charge given –ous suffix
FeCl3 = ferric chloride
FeCl2 = ferrous chloride

33. IV. Additional Complications
To make naming ionic compounds harder, sometimes polyatomic ions are involved.
polyatomic ion: an ion composed of two or more atoms

34. IV. Common Polyatomic Ions
Name
Formula
ammonium
NH4+
hydrogen sulfate
HSO4-
permanganate
MnO4-
sulfite
SO32-
peroxide
O22-
hydrogen sulfite
HSO3-
hydroxide
OH-
phosphate
PO43-
acetate
CH3COO-
hydrogen phosphate
HPO42-
cyanide
CN-
dihydrogen phosphate
H2PO4-
azide
N3-
perchlorate
ClO4-
carbonate
CO32-
chlorate
ClO3-
bicarbonate
HCO3-
chlorite
ClO2-
nitrate
NO3-
hypochlorite
ClO-
nitrite
NO2-
chromate
CrO42-
sulfate
SO42-
dichromate
Cr2O72-

35. IV. Oxyanion Families
Oxyanions are anions that contain oxygen and another element.
There are families of oxyanions, and they have a systematic naming system.
Have either two- or four-member families.
e.g. NO2- and NO3-
e.g. ClO-, ClO2-, ClO3-, and ClO4-

36. IV. Two-Member Families
For a two-member family, oxoanion with fewer O atoms is given the “–ite” suffix while the one with more O atoms is given the “–ate” suffix.
e.g. NO2- = nitrite and NO3- = nitrate

37. IV. Four-Member Families
For the four-member families, the prefixes “hypo-” and “per-” are used to indicate fewer or more oxygen atoms.
e.g. the chlorine oxoanions
ClO- = hypochlorite
ClO2- = chlorite
ClO3- = chlorate
ClO4- = perchlorate

38. IV. Hydrated Ionic Compounds
Ionics with trapped waters are are called hydrates.
Greek prefixes are used to indicate #’s of trapped waters.
e.g. cobalt(II) chloride hexahydrate.

39. IV. Naming Practice
e.g. Give names or formulas for the following compounds.
Na2CO3
magnesium hydroxide
CuSO4·5H2O
CoPO4
nickel(II) sulfate
NaClO2

40. IV. Covalent Nomenclature
For covalent compounds, many different compounds can exist from the same two elements.
e.g. NO, NO2, N2O, N2O3, N2O4, N2O5!
Therefore, we need a systematic naming method.

41. IV. Type III Compounds
Type III compounds are covalent (nonmetal bonded to nonmetal).
Naming rules:
Element w/ lower group # is named 1st using the normal element name EXCEPT when halogens are bonded to oxygen.
If elements are in the same group, lower element named first.
Second element is named using its root and the “-ide” suffix.
#’s of atoms indicated with Greek prefixes EXCEPT when there is only one atom of the first element.

42. IV. Greek Prefixes

43. IV. Type III Compounds Some examples: ClO2 = chlorine dioxide
N2O5 = dinitrogen pentoxide
S2Cl2 = disulfur dichloride
SeF6 = selenium hexafluoride

44. IV. Acids
There are many definitions of acids and bases, but we will use the Arrhenius definitions for now.
Acids are molecular compounds that produce H+ ions in aqueous solution.
We will look at the chemistry of acids in a later chapter.

45. IV. Acid Nomenclature
There are two categories of acids that have different naming rules.
Binary acids
Oxyacids
You can easily recognize acids because their formula has H as the first element.

46. IV. Naming Binary Acids Binary acids contain a nonmetal anion.
HCl = hydrochloric acid
HBr = hydrobromic acid
H2Se = hydroselenic acid
HI = hydroiodic acid

47. IV. Naming Oxyacids Oxyacids contain an oxyanion. Set 1 Set 2
HNO3 = nitric acid
H2SO4 = sulfuric acid
HClO3 = chloric acid
HClO4 = perchloric acid
Set 2
HNO2 = nitrous acid
HClO2 = chlorous acid
HClO = hypochlorous acid
H2SO3 = sulfurous acid

48. IV. Naming Practice
e.g. Give the correct formula or name of the compounds below.
CoCl3
dichlorine heptaoxide
HBr(aq)
SrO
magnesium hydroxide
carbon tetrachloride
MgSO4·7H2O
sodium hydride
HIO3
V2O5
Ru(ClO4)3
NI3
titanium(IV) oxide
N2F2

49. V. Lewis Dot Symbols
Valence e-’s are the most important e-’s in bonding.
Lewis dot symbols are a way to depict the valence e-’s of atoms.
Lewis dot symbols have two parts:
element symbol: represents nucleus and core e-
dots around symbol: represent valence e-’s

50. V. Lewis Dot Symbols
The number of valence e- is given by the element’s group number!!

51. V. Depicting Ionic Bonds

52. V. The Octet Rule
Lewis dot structures are more often used for covalent bonding.
Main group elements tend to form enough bonds to have 8 valence e-’s.
Known as the octet rule (duet for H and He).
Using Lewis dot symbols, it becomes like a puzzle.

53. V. Single Bonds
Single bonds are the basic connections that hold covalent compounds together.
Single bonds form when one pair of e-’s are shared.
Note that lone pairs are e- pairs that don’t participate in bonding.

54. V. Double Bonds
When more e-’s need to be shared to reach an octet, a double bond is possible.

55. V. Triple Bonds
When even more e-’s need to be shared, a triple bond is possible.
e.g. molecular nitrogen, :N≡N:
As the bond order increases, the bond gets stronger and shorter.

56. V. Steps for Drawing Lewis Structures
Determine total # of valence e-.
Place atom w/ lower Group # (lower electronegativity) as the central atom.
Attach other atoms to central atom with single bonds.
Fill octet of outer atoms. (Why?)
Count # of e- used so far. Place remaining e- on central atom in pairs.
If necessary, form higher order bonds to satisfy octet rule of central atom.
Allow expanded octet for central atoms from Period 3 or higher (n ≥ 3).

57. V. Lewis Structure Practice
Draw correct Lewis structures for NF3, CO2, SeCl2, PI5, IF2-, IF6+, and H2CO.

58. V. Exceptions to the Octet Rule
We’ve already discussed expanded valence (hypervalent molecule) cases, but there are other exceptions as well.
Compounds w/ odd # of e-’s: free radicals. Examples include NO and NO2.
Incomplete octets: e- deficient atoms like Be and B, e.g. BeCl2 and BF3.
Expanded octets – when d orbitals are used to accommodate more than an octet.

59. VI. Multiple Valid Lewis Structures
Sometimes more than one Lewis structure can be drawn for the same molecule.
For example, ozone (O3).

60. VI. Resonance Structures
Resonance structures are also known as resonance forms.
A resonance structure is one of two or more Lewis structures that have the same skeletal structure (atoms in same place), but different electron arrangements.

61. VI. Resonance Hybrid
Neither resonance form is a true picture of the molecule.
The molecule exists as a resonance hybrid, which is an average of all resonance forms.
In a resonance hybrid, e- are delocalized over the entire molecule.

62. VI. Sample Problem
Draw the resonance structures of the carbonate anion.

63. VI. Important Resonance Forms
If all resonance forms have the same surrounding atoms, then each contributes equally to the resonance hybrid.
If this is not the case, then one or more resonance forms will dominate the resonance hybrid.
How can we determine which forms will dominate?

64. VI. Formal Charge
formal charge: the charge an atom would have if bonding e- were shared equally
formal charge = (# valence e-) – (# lone pair e-) – (½ # bonding e-)

65. VI. Formal Charges in O3
We calculate formal charge for each atom in the molecule.
For oxygen atom A (left structure), there are 6 valence e-, 4 lone pair e-, and 4 bonding e-. The formal charge for this O atom is 0.
NOTE: sum of all formal charges must equal the overall charge of the molecule!

66. VI. Using Formal Charges
Formal charges help us decide the most important resonance forms (or most likely structure):
Structure w/ all f.c.’s equal to 0 preferable.
Structure with smallest nonzero f.c.’s preferable.
Structure with adjacent f.c.’s of 0 or of opposite sign preferable.
Structure w/ negative f.c.’s on higher EN elements preferable.

67. VI. Sample Problem
Find the dominant resonance structures for the sulfate anion.

68. VI. Sample Problem
Draw the three possible Lewis structures for CO2 and use formal charges to show which structure is most likely.

69. VII. Chemistry Happens in 3D!
To understand how molecules react and interact with each other, we need to know how they look in 3 dimensions.
Important details include bond angle and bond length.

70. VII. VSEPR Theory
From a correct Lewis structure, we can get to the 3D shape using this theory.
VSEPR stands for valence shell electron pair repulsion.
The theory is based on the idea that e- pairs want to get as far away from each other as possible!

71. VII. VSEPR Categories
There are 5 electron pair geometries from which all molecular structures derive.

72. VII. VSEPR Categories

73. VII. Drawing w/ Perspective
We use the conventions below to depict a 3D object on a 2D surface.

74. VII. Determining 3D Shape
The 5 electron pair geometries (EG) are a starting point.
To determine the molecular structure (MG), we consider the # of atoms and the # of e- pairs that are associated w/ the central atom.
All the possibilities for molecular structure can be listed in a classification chart.

75. VII. Linear/Trigonal Planar Geometries
First, we have the linear and trigonal planar EG’s. EG
Bonds
Lone Pairs MG
Linear 2
linear
Trigonal planar 3
trigonal planar 1
bent

76. VII. Tetrahedral GeometriesEG
Bonds
Lone Pairs MG
Tetrahedral 4
tetrahedral 3 1
pyramidal 2
bent
linear

77. VII. Trigonal Bipyramidal GeometriesEG
Bonds
Lone Pairs MG
Trigonal Bipyramidal 5
trigonal bipyramidal 4 1
see-saw 3 2
T-shaped
linear

78. VII. Octahedral GeometriesEG
Bonds
Lone Pairs MG
Octahedral 6
octahedral 5 1
square pyramidal 4 2
square planar 3
T-shaped
linear

79. VII. Summary Chart

80. VII. Steps to Determine Molecular Structure
Draw Lewis structure.
Count # of bonds (double/triple count as one) and lone pair e-’s on the central atom.
Select electronic pair geometry.
Place e-’s and atoms that lead to most stable arrangement (minimize e- repulsions).
Determine molecular geometry.

81. VII. Trig Bipy is Special
In other EG’s, all positions are equivalent.
In trig bipy, lone pairs always choose to go equatorial first.
Why?

82. VII. Lone Pairs Take Up Space
Lone pair e-’s don’t have another nucleus to “anchor” them.

83. VII. Distortion of Angles
Lone pair e-’s take up a lot of room, and they distort the optimum angles seen in the EG’s.

84. VII. Some Practice
Draw the molecular geometries for SF4, BeCl2, ClO2-, TeF5-, ClF3, and NF3.

85. VII. Larger Molecules

86. VII. Molecular Polarity
Individual bonds tend to be polar, but that doesn’t mean that a molecule will be polar overall.
To determine molecular polarity, you need to consider the 3D shape and see if there’s a net dipole moment.
Net dipole moments arise when bond dipole moments (vectors) don’t cancel.

87. VII. Bond and Net Dipole Moments

88. VII. Sample Problem
Determine the molecular structure of IF2- and state whether it is polar or nonpolar.

89. VII. Polarity and Properties
Polarity is the result of a compound’s composition and structure.
Knowing that a compound is polar/nonpolar allows us to explain some of its properties.

90. VII. Alignment in Electric Field

91. VII. “Like Dissolves Like”