1. The 3 types of Chemical Bonds: Ionic, Covalent, & Metallic
How do they bond?
Will gain, share or lose electrons
Why do they bond?
Gain stability
Full outer shell
Lower energy = more stable
Depend on what is bonding .

2. Electronegativity and Bond type
Bond type can be determined by the difference in electronegativity between the bonds involved.
Differences of:
≥ 2 = ionic
0.5 – 1.9 = polar covalent
≤ 0.4 = nonpolar covalent
Metals bonded to other metals are metallically bonded, regardless of the difference in electronegativity .

3. Metallic bonds
The atoms of metals are held together when the atom’s valence electrons float around the nuclei of the metals – the “sea of electrons”
Electrostatic forces keep everything together

4. Ionic Bonds:
Complete transfer of 1 or more electron from one atom to another (or another group)
one loses one or more e-, the other gains those e-
Atoms involved are
a metal and either a non-metal or a polyatomic ion
The cation and anion are attracted to each other by electrostatic attraction .

5. Covalent Bonds
2, 4, or 6 valence electrons that are shared between atoms
Molecules are formed when 2 or more atoms bond covalently.
We are going to name only simple covalent compounds that have 2 elements involved

6. Most bonds are somewhere in between ionic and covalent
Because not all atoms share e- equally
The conventions of naming assume absolute difference in bond types
Metals bonded to nonmetals or polyatomic ions are classified as having ionic bonds*
Materials made out of all non-metals are classified as having covalent bonds*
* semimetals are not a classification in naming; you need to treat the elements that are on the right of the line as non-metals, and those on the left as metals.
More on this when we talk about bonding

7. Covalent Bond Formation
Diatomic Molecules
Remember Br I N Cl H O F
These elements are more stable as a pair than as individual atoms
The arrangement of atoms in a covalent bond exist at an optimal distance between the nuclei.
Minimize electron-electron repulsion
Minimize nucleus-nucleus repulsion
Maximize nucleus-electron attraction
(*disrupt this balance and the bond can be broken!*)

8. Single Covalent Bond Formation
One (1) pair of electrons shared
Called the bonding pair
Can form up to 4 single bonds (at most)
a.k.a. sigma bond (σ)
Pair of shared e-’s in area centered between the 2 atoms
Where valence atomic orbitals overlap end to end to make a bonding orbital
s orbital overlaps another s orbital or p orbital
two p orbitals overlap

9. Lewis Structures
Use electron dot diagrams to show the arrangement of electrons in molecules
A line or pair of vertical dots between elements stand for a single bond.

10. Lewis Structures
Remember that all atoms want a full outer shell, so do Lewis structures
For a given Lewis structure, the number of electrons around the atoms must equal the total number of electrons individually assigned.
Ex: C has 4, H has 1, so CH4 must have 8 total

11. Sigma Bond Formation

12. Multiple Covalent Bond Formation
Multiple Covalent Bonds
More than one pair of electrons shared
Double bond: 2 pairs of e-’s are shared (4 e-’s total)
Triple bond: 3 pairs of e-’s are shared (6 e-’s total)
a.k.a. pi bond (π)
Shared e-’s in area above and below the line representing the bond
Parallel orbitals overlap end to end to make a bonding orbital
One sigma bond and at least one pi bond
double bond = one sigma bond and one pi bond
triple bond = one sigma bond and two pi bonds

13. Lewis Structures
(A): a line (or pair of vertical dots) between elements stand for a single bond.
(B) and (C): 2 or more lines (or pairs of vertical dots) between elements stand for a double and triple bonds.

14. Pi Bond Formation
C2H2
HC º CH
C2H4
H2C =CH2

15. Bond Energy Bond energy: the amount of energy needed to break a bond
Comes from the balance of the repulsive and attractive forces
Usually measured in KJ/ mol
a.k.a.bond-dissociation energy
Bond energy for each type of bond (Ex: C-C single, C=C double, C-O, C=O, H-H, C-N, C=N, etc.) is different
Why? Different atomic radii for each element, different electronegativity, bond order, bond length etc.)

16. Bond Energy Bond energy always changes as bonds break and form
Energy is released when bonds form
Energy is absorbed when bonds break
Bond energy relates to bond strength
High bond energy = strong bond
Low bond energy = weak bond
Total energy of a reaction depends on the energy needed to break bonds and reform new ones
Endothermic: energy is absorbed because a greater amount of energy is needed to break the existing bonds than is released making new ones
Exothermic: more energy is released during the making new bonds than is needed to break the existing bonds

17. Bonds and Bond Length
The distance between the nuclei at a position of maximum attraction
Influenced by atomic radius
Number of bonds (bond order)
Bond length affects bond strength
Shorter bond = stronger bond
Longer bond = weaker bond

18. Bonds and Bond Order Bond order tells the number of bonds
A single bond has a bond order of 1
Multiple bonds
double: Bond order of 2
triple: Bond order of 3
quadruple bonds do not occur (too hard to share 8 electrons; steric hinderance)

19. Bond Energy and Bond Order
Single bonds have lowest bond energy
Double bonds have higher bond energy
Triple bonds have highest bond energy

20. Bond energies of common bonds (kJ/ mol)
C-C 376
C=C 720
C º C 862
N º N 945

21. Atomic Radii and Bond Energy
Bond Bond Energy
H-F
H-Cl
H-Br
Cl-Cl
Br-Br
I-I

22. Bond Order and Bond Length
Bond Energy
Bond Length
C-C 1
348 kJ/mol
1.54 Å
C=C 2
614 kJ/mol
1.34 Å
CºC 3
839 kJ/mol
1.20 Å
N-N
163 kJ/mol
1.47 Å
N=N
418 kJ/mol
1.24 Å
NºN
941 kJ/mol
1.10 Å

23. Naming things: Is there a metal first?
If there is only one element present, name it. Atomic substances do not require “special” naming.
For anything with more than one element, remember that there is ONE MAIN THING to look for:
Is there a metal first?

24. So… Look to see if there is a metal first in the formula
Again, semimetals are not a classification in naming; you need to treat the elements that are on the right of the line as non-metals, and those on the left as metals.
If there are only metals, then metallic bonding
If first element is a metal, then the compound is an ionic compound
Nonmetals only signify a covalent compound

25. Naming rules for binary covalent compounds (2 elements)
Names are two words, with prefixes.
First element in the formula is always named first with no changes
Second element in the formula is named using its root and changing the ending to -ide.
Prefixes are used to indicate the number of atoms of in the compound.
The prefix “mono-” is never used with the 1st element
Ex. CO is carbon monoxide
If 2 vowels are next to each other after adding the prefix, one of the vowels is dropped
Ex. CO is not carbon monooxide

26. Covalent Naming Prefixes:
mono-
di-
tri-
tetra-
penta-
hexa-
hepta-
octa-
nona-
deca-
NUMBER

27. Common Names
A lot of chemicals have common names as well as the proper IUPAC name.
Chemicals that should always be named by common name and never named by the IUPAC method are:
H2O water, not dihydrogen monoxide
NH3 ammonia, not nitrogen trihydride

28. Acid Nomenclature Acids Compounds that form H+ in water.
Formulas usually begin with ‘H’.
In order to be an acid instead of a gaseous covalent compound, it must be aqueous
Meaning dissolved in water; symbolized by (aq)
Ternary acids are ALL aqueous
Two types:
Oxyacids
Non-oxyacids

29. Naming acids If the acid doesn’t have oxygen add the prefix hydro-
change the suffix -ide to -ic acid
HCl Hydrochloric acid
H2S Hydrosulfic acid
HCN Hydrocyanic acid

30. Naming acids If the formula has oxygen in it
write the name of the anion, but change
ate to -ic acid
ite to -ous acid
Watch out for sulfuric and sulfurous!
H2CrO4
HMnO4
HNO2
Chromic acid
Manganic acid
Nitrous acid

31. Acid Nomenclature Flowchart

32. Acid Nomenclature An easy way to remember which goes with which…
Binary 
Ternary
An easy way to remember which goes with which…
“In the cafeteria, you ATE something ICky”

33. Acid Nomenclature HBr (aq) H2CO3 (aq) H2CO2 (aq)  hydrobromic acid
2 elements, -ide
 hydrobromic acid
3 elements, -ate
 carbonic acid
3 elements, -ite
 carbonous acid

34. Acid
Name
HNO3
Nitric acid
HNO2
Nitrous acid
H2SO4
Sulfuric acid
H2SO3
Sulfurous acid
H3PO4
Phosphoric acid
HC2H3O2
Acetic Acid

35. Since you are sharing electrons, rather than giving them away/ picking them up, the charges are NOT relevant.
never

36. Use the Prefixes: PREFIX NUMBER 1 2 3 4 5 6 7 8 9 10 mono- di- tri-
tetra-
penta-
hexa-
hepta-
octa-
nona-
deca-
NUMBER

37. Formulas for acids Backwards from names.
If it has hydro- in the name it has no oxygen
Anion ends in -ide
No hydro, anion ends in -ate or -ite
Write anion and add enough H to balance the charges.

38. Molecular Structures

39. Molecular Structures
Lewis structures: electron dot diagrams show the arrangement of electrons
Ball and stick and space filling
Colored spheres stand for the elements bonded
Typically O is red, H is white, C is black, N is blue, halides are green
Structural formula
Use letter symbols and bonds
Can be predicted from the Lewis structure

40. Drawing Lewis Structures
Add the valence electrons of all the elements bonding (add/remove any for polyatomic ions).
Identify the central atom (usually the one with the lowest EN, highest molecular mass and closest to the center of the periodic table).
2b. Place the central atom in the center of the molecule and add all other atoms around it.
Place one bond (two electrons) between each pair of atoms.

41. Drawing Lewis Structures
4. Determine the number of e- remaining. (Subtract the # of e-’s used in the bonds from the e-’s counted in #1)
4b. Place lone pair e-’s around all the atoms surrounding the central atom completing their octets. Any remaining pairs go on the center atom.
5. If the center atom does not have an octet around, convert one or two of the lone pairs on the outer atoms into double or triple bonds.

42. Determining formal charge
Determining formal charge
Formal charge can be determined by:
Normal number of electrons in outer shell -
[(1/2 the number of bonded electrons) + lone electrons]
_____________________________________
= formal charge

43. Formal Charge, continued
Example: N in NH4+
FC =5- [(1/2 of 8)+ 0]= +1
H in NH4 +
FC =1- [(1/2 of 2)+ 0]= 0
Overall, the formal charge on NH4+ is + , so we write NH4+ or as [NH4]+
This bracketed version is typically used, and is more precise for reasons we have yet to get into, but we will.

44. Formal charge and stability
Formal charge and stability
The most “happy” molecules tend to have no formal charges
However, molecules may be “happy” if they have no NET charge on them
if there is 1+ and 1-, so a net of +1 + (-1)=0
Structures that are the best have a minimal formal charge and a full octet (valence shell) around each atom

45. Resonance Structures
Have the same alignment of atoms, but different bonding (electrons ONLY are moved, both in bonds and lone pairs)

46. Resonance structures of BF3
Remember that none of these is a real picture of BF3, but the real picture is a hybrid of all of these. (BO=1.3 for all B-F bonds)
Some of these are better pictures of what really happens than others: the better ones are those with the least formal charge.

47. AGAIN: Formal charge and stability
AGAIN: Formal charge and stability
The most “happy” molecules tend to have no formal charges
However, molecules may be “happy” if they have no NET charge on them
if there is 1+ and 1-, so a net of +1 + (-1)=0
Resonance structures that are better structures have a minimal formal charge and a full octet (valence shell) around each atom

48. Resonance, Formal Charge and Good Lewis Structures
Those that are not as good, but fulfill the octet rule are considered to be minor contributors
They can exist, but not as much of the overall picture looks like them
If it doesn’t fulfill the octet rule and it isn’t a KNOWN exception*, and/ or it has crazy formal charges, toss it. It doesn’t work.
(* exceptions: next slide)

49. Words About Exceptions:
They happen. A fair amount of the time with CERTAIN ELEMENTS- not with everything.
Hence, exceptions to the general rules, but rules for those elements are now different.

50. Words About Exceptions:
Q: Why do chemists do this to me?
A: Because they hate you.
Q: Really?
A: No, they don’t. The properties of the atoms set the rules, based upon their electron configurations and other properties discussed in class (IE, EA, AR, shielding, Zeff….) Chemists do not make up rules to annoy you; they make up general rules to fit MOST situations; elements sometimes don’t fit the constraints.
Q: Will the exceptions mess me up?
A: Not if you can follow all of the other rules.

51. Deficient* atoms: *(less than a full shell)
Boron: USUALLY keeps 6 electrons in its outer shell. It just does, because of being small and being a nonmetal.
Basically, every element below is a metal, and gives up electrons to bond ionically.
But B is so small it wants more. So it bonds covalently to get them. But it can’t go to a full 8- it’s too many.

52. Coordinate Covalent Bonds:
Forms when one atom donates both of the electrons to be shared (instead of one e- from each element).
Atoms that start with lone pair e-’s often form these bonds.

53. Expansion of the octet
P, S, halogens and noble gases (yeah, I know- we’ll explain why later on) heavier than Br (Br, I, At, Kr, Xe, Rn) [In other words…any nonmetal in the 3rd period or greater (up to Rn) excluding Cl and Ar]
Typically 5 or 6 PAIRS of electrons, and these are added as lone pairs, not as bonded pairs.
Add lone pairs on the central atom only until the number of electrons needed is reached.
Ex: PCl5, I3, SF4, XeF3

54. An excerpt from a web site on expansion:
 The concept of the Expanded Octet occurs in any system that has an atom with more than four electron pairs attached to it.
Most commonly, atoms will expand their octets to contain a total of five or six electron pairs, in total. In theory, it is possible to expand beyond those number.
The large amounts of negative charge concentrated in small volumes of space prevent those larger expanded octets from forming.
When an atom expands its octet, it does so by making use of empty d orbitals that are available in the valence level of the atom doing the expanding.
Atoms that do not have empty valence level d orbitals will not be able to expand their octets. The atom that expands its octet in a structure will usually be located in the center of the structure and the system will not use any multiple bonds in attaching atoms to the central atom.
The process of expanding octets is strictly a last resort on the part of atoms.

55. More theft: But this is well stated at http://www. towson
“Expanded octet” refers to the Lewis structures where the central atom ends up with more than an octet, such as in PCl5 or XeF4. In drawing the Lewis structure for PCl5 , there is a total of 40 valence electrons to put in (5 + 5x7 = 40). One can easily see that if the central atom, P, is to be joined to five Cl atoms, P would have 10 electrons instead of the octet. (Remember that one does NOT string out the elements to look thus: P-Cl-Cl-Cl-Cl-Cl.  No !!)
Clearly there is a violation of the “Octet Rule”.
How do we know to allow this violation? It’s simply this: PCl5 exists. Our rules have to be revised to accommodate observations. If PCl5 exists, then this violation must be permitted, since there is no other way to explain it. Again, chemists don’t hate you. Nature just does things that we didn’t plan on when we use the easiest explanations.

56. Exceptions continued…
A word of caution: Does this mean that we can violate the Octet Rule any time we wish? No! We can violate the Octet Rule only when there is no other way to explain it.
Limitations to the “expanded octet”:
This cannot occur with elements that do not contain d-orbitals! Which elements would that be? Elements smaller than neon (atomic number 10) have electrons only in the first two main energy levels (n = 1 and n = 2), and those energy levels do not contain d-orbitals.

57. More on exceptions…
Second of all, expanded octets generally occur when there are too many electrons to fit in. Since double bonds and triple bonds occur only when there are insufficient number of electrons, you would not normally apply the “expanded octet” to central atoms with multiple bonds! In other words, the extra electrons must be added as lone pairs and NOT as double or triple bonds!

58. Summary of limitations of applying the “expanded octet”:
The central atom with an expanded octet
MUST have an atomic number larger than 10 (beyond Ne)
(this means it has available d orbitals to extend into)
Extra electrons should be first placed on the outside atoms to fulfill their octets.
After that, there are still extra electrons, start with placing them as lone pairs on the central atom.
If the central atom does not have a positive formal charge, do not go any further. You have the correct Lewis structure.
Only if the central atom has a positive charge should you move a lone pair from the outside atoms to share (to neutralize the formal positive charge)
Do NOT indiscriminately double or triple bond!

59. Let’s make XeF4
Step One: Count the total number of valence electrons. 8 from Xe+ 4 F x (7e- per F) = 36e-, or 18 pairs.
Step Two: The first element is usually the central atom, and then you cluster the other atoms around it.
Step Three: The 4 covalent bonds shown above account for 4 pairs. As you put lone pairs onto the surrounding F, you would account for 12 more pairs, giving you a total of 16 pairs. Where are you going to put two more pairs? The only place would be on the central atom (on Xe) as lone pairs. Xe would therefore have 4 atoms and a lone pair (AX4E2).

60. Xenon now has twelve electrons instead of the octet
Xenon now has twelve electrons instead of the octet! This is called an “expanded octet”, expanding beyond the octet.
Notice we do NOT put a double bond between the Xe and one of the F! A mistake that students often make. In the structure shown above, Xe has a formal charge of zero, so there is no reason to do any more to the structure.

61. This structure is incorrect because Xe now has a formal charge of 2+ (6 electrons instead of 8) and the two F with double bonds each has a formal charge of −1 (8 electrons instead of 7). Compared to the structure above which has no formal charges, this is certainly not preferred.
In addition, in terms of the Octet Rule, F has exceed the Octet and it is not an element than can exceed the Octet Rule (atomic number les than 10).

62. What is VSEPR?
Valence
Shell
Electron
Pair
Repulsion
Theory

63. Why? The shape of molecules influences their characteristics:
Why?
The shape of molecules influences their characteristics:
Things like polarity which influence things like
boiling point, melting point, which dictate their nature (solid, liquid or gas at room temperature)
Minimizes the repulsion of shared and unshared electron pairs around the central atom.

64. Hybridization
Hybridization: process by which atomic orbitals mix and form new, identical hybrid orbitals.
Come from the valence e-’s
Each hybrid orbital contains one e- that can share with another atom.
The number of atomic orbitals that mix = # of e- pairs (in bonds or lone e-’s)
In CH4: C has 2 s and p e-’s and forms bonds so has 4 hybrid orbitals = sp3

65. Hybridization Examples
In CH4:
C has 2 s and 2 p e-’s
1 of the s e-’s moves into the open p orbital
forms 4 bonds
has 4 hybrid orbitals = sp3

66. Hybridization Examples
In BF3:
B has 2 s and 1 p e-
1 of the s e-’s moves into the open p orbital and 1 p orbital remains unoccupied
forms 3 bonds so has 3 hybrid orbitals = sp2

67. Hybridization Lone pairs occupy hybrid orbitals BeCl2 H2O
Lone pairs occupy hybrid orbitals
BeCl2 vs. H2O
BeCl2 H2O

68. The parent geometries: all others come from these

69. Steric Number The number of “things” sprouting off of an atom
These can be either
Bonds
Of any order (1, 2, or 3) Or
Lone pairs of electrons

70. Steric Number Examples
Ex #1: CH4
There are 4 H’s branching off , so the steric number is 4
SN=4
Ex #2: H2O
SN= 4
Explain why
Ex #3: CO2
SN= 2

71. General Formulas
All molecules with a shared general formula have a shared geometry
we use them to help note shape
Formulas are typically written with A’s, X’s, and E’s
The letters stand for:
A= the central atom
X *= the number of atoms attached to the central atom
E= the number of lone pairs of electrons attached to the central atom
*Some sources use A’s, B’s, and E’s

72. General Formula Examples
Ex #1: CH4
AX4
Ex #2: H2O
AX2E2
Ex #3: CO2
AX2

73. Linear Trigonal planar
AX2
AX3
Tetrahedral
AX4

74. Trigonal Pyramidal Bent Tetrahedral parent shape
Tetrahedral parent shape
1 lone pair of electrons
AX3E
Tetrahedral parent shape
2 lone pair of electrons
AX2E2

75. Trigonal bipyramidal
AX5

76. Seesaw a.k.a. Teeter-totter
Trigonal bipyramidal parent shape
1 lone pair of electrons
AX4E

77. T-shaped Trigonal bipyramidal parent shape 2 lone pair of electrons
AX3E2

78. Linear Trigonal bipyramidal parent shape 3 lone pair of electrons
AX2E3

79. Octahedral
AX6

80. Square pyramidal
Octahedral parent shape
1 lone pair of electrons
AX5E

81. Square planar
Octahedral parent shape
2 lone pair of electrons
AX4E2

83. Sweet drill and practice web site
Given generic shapes to ID:
Given molecules to draw out:
Basic:
Advanced:

84. Electronegativity and Bond Type
Determine the absolute value between the electronegativities of the atoms involved
If it is , ≥ 2 then the bond is ionic
If it is 0.5 ≥1.9 then the bond is polar covalent
Electrons are shared unequally between the atoms
resulting in partial charges (dipoles)
If it is ≤ 0.4, then the bond is nonpolar covalent
Electrons are shared equally between the atoms of the bond
no partial charges

85. Examples: A bond from H-O H has an electronegativity of 2.1
O has an electronegativity of 3.5
The difference is 1.4, so the bond is polar
The more electronegative element has a slightly negative charge (here, O)
The less electronegative element has a slightly positive charge (here, H)

86. Electronegativity and Bond Polarity
K-F
Ionic
ED = 3.2
Cl-F
Covalent
Polar
ED = 1.0
F-F
NonPolar
ED = 0.0

87. When determining polarity it is important to look at the dipole moments- do they cancel out?

90. Bond Type and Properties
Solubility: ability to dissolve
Depends on the type of bond
”like dissolves like”
Polar compounds dissolve in polar and ionic compounds (vinegar in water)
Nonpolar compounds dissolve in nonpolar compounds (oils in rubbing alcohol)
Low melting and boiling points
Soft molecular solids
Covalent network solids
Atoms connected by a network of covalent bonds
Brittle but very hard
Ex. Quartz and diamond