We will be going over information for Exam 2. Talking a lot about naming of compounds and learning electron domain geometries with molecular geometries.

1. CHEM TEST 2
By Madison Archer

2. Identify types of bonds in a compound
Covalent bonds and their characteristics.
Ionic bonds and characteristics
Know formulas, names and charges of polyatomic ions.
Names and formulas of ionic compounds.
Name molecular compounds and acids from their
formulas, and v.v.
Given a chem formula, convert grams to moles to
numbers of particles, and v.v.
“Dilution problems”
Unit 3 Overview

3. Types of Bonds
Ionic
Metal+Nonmetal
-Electrons are transferred between atoms
Covalent
Nonmetal+Nonmetal
-Electrons are shared between atoms
Metallic
Metal+Metal
-Electrons are “pooled” between atoms
BREAKING bonds requires (absorbs) energy- Endothermic
FORMING bonds releases energy- Exothermic

4. Bond formed between metals and non-metals
Exception: Due to it’s positive charge, the NH4
+ ion acts
like a metal. So, the bond between NH4
+ and “X” will be ionic,
even if no metal is present.
-When place in water, ionic bonds tend to break into their
separate ions. Two things will happen
1. The more electronegative atom will take the
electrons with it, forming an anion.
2. The less electronegative atom will now have
lost its electrons, forming a cation.
Ionic Bonding

5. Ionic Bonding
-High melting points (>300°)
-Hard, Brittle
-Not conductive, UNLESS in
a liquid state or dissolved in
water.
Formation of an ionic compound is
accounted for by ionization energy,
electron affinity, and lattice energy.
Ionic compounds are stable due to
the electrostatic attraction between
its positive and negative ions. The
measure of this attraction is called the
lattice energy.

6. Ionic Bonding Explained
To understand ionic bonding, it helps to compare ionization
energies and electron affinities.
Ionization Energy
Metals tend to lose
electrons, such as group 1-2
elements, due to their low
ionization energies.
Electron Affinity
Non-metals tend to gain
electrons, such as fluorine,
due to their high electron
affinities.
In reactions involving both metals and non-metals, electrons are
transferred from the metal to the non-metal. This causes the
metal to oxidize and the non-metal to be reduced. Example on
next slide 

7. 1. Na atom loses electron to form the Na+ ion
(due to its low ionization energy)
2. Cl molecule gains electron to form the Cl- ions
(due to its high electron affinity)
When doing this, energy is either gained or released, depending on what is
happening. The energies needed for this are shown for Na and Cl below.
Do you see a problem??
Ionic Bond: Na+ + Cl- = NaCl
Low Ionization Energy Na → Na+ + e- +496 kJ/mol (endothermic)
High Electron Affinity Cl + e- → Cl- -349 kJ/mol (exothermic)
Looking at this you can do the math and see that forming NaCl requires
an addition of at LEAST +147 kJ/mol of energy to turn the atoms into
reactive ions, so what’s happening?

8. Since the atoms are moving towards each other due to
their attraction, they come together in a lattice formation. But
guess what, the formation of a salt lattice ALSO takes energy!
When forming NaCl, we know we still need +147 kJ/mol, but once
the crystal lattice forms, it releases a whooping
-788 kJ/mol of energy, causing a highly exothermic release of
-641 kJ/mol of energy.
We know the formation of a bond is exothermic, so
where is the extra energy being released from?
What’s happening then??

9. The magnitude of the lattice energy depends directly on the size
of charges and inversely on the distance between the ions
meaning…
Lattice energy = charge of the ion
Lattice energy = size of ions
(also depends on the arrangement of the ions in the salt lattice)
Energy is gained when ions break apart the lattice (Endothermic)
Energy is released when ions assemble into a lattice
(Exothermic)
Lattice Energy Explained

10. We know that F in the most electronegative atom and Cs is near the
bottom of group-1, therefore making CsF the most ionic.
Then we know CsI has a much greater difference in electronegativity that
NF, so that makes the answer b.
Real Test Question

11. (Non-metal + Non-metal)
Electrons are equally shared, but why?
The two positively charged nucleii repel each other and the two
negatively charged electrons repel each other, BUT the nucleii and
electrons of the opposite atoms are attracted to each other.
Attractive forces MUST exceed the repulsive forces for ions of bond
Formation of a covalent bond is due to non-metals:
1. high ionization energies
(making it difficult to remove their valence electrons)
2. high electron affinities
(causing the non-metals to also be attracted to each other)
Covalent Bonding

12. Bond Order: the number of electron pairs being shared by a given
pair of atoms.
Single bond: bond order=1
Double bond: bond order=2
Triple bond: bond order=3
Bond Energy (BE): the energy needed to overcome the attraction
between the nuclei and the shared electrons, in other words; the
amount of energy it would take to break the bond
Bond Length: the distance between the nuclei of the bonded
atoms.
Covalent Bonding Terms

13. Lewis Theory: More electrons two atoms share, the shorter the bond.
Triple bond=shortest Double bond=middle Single bond=longest
Bond lengths are affected by: 1. bond order, duhhh
2. size of atoms
3. bond energy
Bond energy/strength: amount of energy need to break the bond
1. The more electrons two atoms share, the stronger the bond.
Triple bond > Double bond > Single bond
2. The shorter the bond means the stronger the bond.
Br—F (237 kJ) > Br—Cl (218 kJ) > Br—Br (193 kJ)
ALWAYS take into account the atomic radius trend: bonds get stronger going
up and right on the periodic table
Covalent Bonding

14. Covalent Bonding Properties
Intermolecular Attractions: attractions occurring between molecules,
these are weak forces.
Covalent Bonds are significantly stronger than molecular attractions, as
bonds can undergo phase transitions and not break.
Covalent bonding physical properties:
-Low melting/boiling points
-Hardness (resist scratches) and Brittleness (resist breaking)
-NOT conductive, UNLESS it is a molecular acid dissolved in water

15. Obviously a compound containing only non-metals must have all covalent bonds!
Questions involving bond types are usually pretty basic, so know your bonds! Don’t
miss the points they’re trying to give you.
Real Test Question

16. Octet Rule: Atoms tend to gain, lose, or share electrons until they are
surrounded by eight valence electrons
Duet Rule: Atoms share electrons to achieve an octet or duet in covalent
bonds
Each atom counts both electrons in the bond to achieve an octet or duet.
Each H+ in H2would have a 1s2
electron configuration
Octet and Duet Rule
H

17. H, Li, Be, and B attain an electron configuration like that of He.
-He can have ONLY two valence electrons, a duet
-Li loses its one valence electron
-H shares or gains one electron
-Be loses two electrons to become Be2+
-B loses three electrons to become B3+
(commonly shares its three electrons in covalent bonds, giving it six valence electrons)
Exceptions to Octet Rule: He

18. Each dot represents an electron, to represent an atom you must
add one dot at a time to each of the four sides, then begin pairing.
Atoms and Ions interact through valence electrons to achieve
noble gas configurations, of octets and duets.
Lewis Dot Method
Sodium loses an electron and chlorine
gains an electron. The ions are
attracted to each other to form an
ionic bond.
Hydrogen and chlorine share
electrons to form a covalent
bond

19. Review naming monoatomic ions:
Metals: element + “ion” Non-metals: element + “-ide ion”
Na+ =sodium ion O2- =oxide ion
Polyatomic Ions: charged, covalently bonded compounds made of two
or more atoms that are considered a single unit.
Most are oxyanions, anions containing oxygen+another element.
MEMORIZE THESE: common polyatomic ions asked about on exam
-Acetate -Carbonate -Hydroxide -Cyanide
C2H3O2
- CO3
2- OH- CN-
-Phosphate -Ammonium -Borate -Peroxide
PO4
3- NH4
+ BO3
- O2
2-
Naming Polyatomic Ions

20. Naming Oxyanions
Two Ions in Series
-the one with more oxygen atoms
has the ending –ate
-the one with fewer ends in -ite.
NO3
- nitrate SO4
2- sulfate
NO2
- nitrite SO3
-2 sulfite
Four Ions in Series
-the one with one more oxygen atom than
the oxyanion ending in –ate, now also begins
with –per
-the one with one less oxygen atom than the
oxyanion ending is –ite, now begins with –hypo
BrO4
- perbromate
BrO3
- bromate
BrO2
- bromite
BrO- hypobromite
When a series of oxyanions contains different numbers of oxygen atoms, they
are named according to the number of oxygen atoms in the ion.
IO4
- periodate
IO3
- iodate
IO2
- iodite
IO-
hypoiodite

21. Naming Oxyanions with H+
(Anions derived by adding an H+ to an oxyanion)
MEMORIZE BOTH VERSIONS
Current Version
-add the prefix, “hydrogen” or “dihydrogen”
CO3
2- carbonate ion = HCO3
- hydrogen carbonate ion
PO4
3- phosphate ion = HPO4
2- hydrogen phosphate ion
H2PO4
- dihydrogen phosphate ion
Older Method
-used the prefix “bi”, instead of “hydrogen”
HCO3
- bicarbonate ion
HPO4
2- biphosphate ion

22. Naming Binary Ionic Compounds
(1 metal+1 non-metal)
1. metals forming only one ion Formula: metal+non-metal+ide
KCl potassium chloride Na2S sodium sulfide
2. metals forming more than one ion. Some transition metals can form more than
one charge state, so we use roman numerals to show the charge of the metal.
Current Formula: metal+(charge in roman numerals)+non-metal+“-ide”
Naming CrBr3 : 1. Look if it has a net charge, no net charge.
2. Find the charge of non-metal. Br has a -1 charge and it has 3 of them.
3. Calculate. Cr charge+ (-3)=0 Cr= +3 chromium (III) bromide
CuO copper (II) oxide Cu2O copper (I) oxide
Older Formula: –ic indicates higher charge/ –ous indicates lower charge
CuO cupric oxide Cu2O cuprous oxide

23. Ionic compounds containing polyatomic ions: naming is similar to other ionic
compounds, except the name of the polyatomic ion is used whenever it occurs.
NaNO2 sodium nitrite Na3PO4 sodium phosphate
Hydrates: ionic compounds containing a specific number of water molecules
associated with each formula unit (use the listed prefixes)
1=mono 2=di 3=tri 4=tetra 5=penta
6=hexa 7=hepta 8=octa 9=nona 10=deca
Formula: MgSO4  7H2O  magnesium sulfate heptahydrate
Formula: CoCl2  6H2O  cobalt (II) chloride hexahydrate
Naming more Ionic Compounds

24. Writing Formulas for Ionic Compounds
Steps to writing the formula:
1. Write the symbol for the metal + its charge
2. Write the symbol for the non-metal + its charge
3. Adjust subscript on each to balance the overall charge
4. Check that the sum of the charges equal each other
Al3+ O2- Ca2+ O2-
Al2O3 CaO

25. Steps: 1) Write the name of the element with the smallest group number first
(If the elements are in the same group, the element with largest row number goes
first)
2) Find the prefix needed for each non-metal, prefixes are the same as
the ones used in naming hydrates
Formula: (prefix)+ (1st non-metal)+(prefix)+( 2nd non-metal)
nitrogen trichloride: NCl3 -3 Cl, so add the “-tri”
diphophorous pentoxide: P2O5 -2 P, so add “-di” -5 O, so add “-pent”
Naming Molecular Compounds
(two or more non-metals)
PREFIXES
mono=1 di=2 tri=3 tetr(a)=4 pent(a)=5
hex(a)=6 hept(a)=7 oct(a)=8 non(a)=9

26. Acids: molecular compounds that release H+ atoms when dissolved in water
-compounds are NOT considered acids unless they dissolve in water
TYPES:
-Binary Acids: made up of a hydrogen atom and a non-metal
Formula: hydro+name of non-metal+“-ic”+acid
HBr hydrobromic acid HCl hydrochloric acid
-Oxyacids: made up of a hydrogen atom and a polyatomic ion
oxyanions ending with –ate, turn to –ic plus acid
oxyanions ending with –ite, turn to –ous plus acid
-ate  -ic -ite  -ous
ClO3
- chlorate  chloric acid ClO2
- chlorite  chlorous acid
SO4 sulfate  sulfuric acid NO2
- nitrite  nitrous acid
Naming Acids

27. Formula weight (FW): Mass of one ionic unit or one molecule
-the sum of the atomic masses of each atom in its chemical formula, using the
atomic weights and its chemical formula, we can calculate formula weight
Example: Find the formula weight of sulfuric acid, H2SO4
2 H atoms = 2 × 1.00794 amu = 2.01588 amu
1 S atom = 1 × 32.065 amu = 32.065 amu
4 O atoms = 4 × 15.994 amu = 63.9976 amu
98.0784 amu
Formula Weight
+

28. Formula unit: name of the chemical formula of an ionic substance, shows
number of ions rather than molecules (amu)
Molar Mass: mass of one mole (g/mol)
Molecular weight (MW)/molecular mass: the sum of the weights of atoms
present in molecules
Okay, things are going to get a bit confusing now! Don’t freak out though,
once you get it, it’ll be like second nature. You got this!
Terms to know

29. How many moles of Fe2(SO4)3 are in 92.4 grams of Fe2(SO4)3?
Step 1: Calculate molar mass 2 Fe atoms: 2×55.845 = 111.69
3 S atoms: 3×32.066 = 96.198
12 O atoms: 12×15.999 = 191.993
molar mass: 399.88 g Fe2(SO4)3 = 1 mole Fe2(SO4)3 399.88 g
Step 2: Calculate moles using the molar mass of the compound
-we are given grams (92.4), so we need to go “up” the stair stepper and divide
92.4 ÷ 399.88 = 0.231 mol Fe2(SO4)3
+

30. Step 1: Calculate molar mass. Too much to type so….. MM=342.14 g Al2(SO4)3
Step 2: Change GRAMS of the compound to MOLES of the compound
25 ÷ 342.14 = ???? .073 moles of Al2(SO4)3
Step 3: Change MOLES of the compound to ATOMS of the compound
.073× (6.022×1023)= 4.4×1022 moles of Sulfur
Step 4: Change ATOMS of the compound to ATOMS of the element
4.4×1022 atoms of Al2(SO4)3 × 3 atoms of S = 1.32×1023 atoms of S
How many atoms of S are in a 25 g sample of Al2(SO4)3?

31. Concentration: The amount of solute in a total volume of solution
Molarity: moles of a solute per liter of solution, do NOT get this confused
with moles!
Molarity (M)=
Note: volume must be in liters, but on exams they will usually give
this to you in mL, don’t forget to convert!
Molarity of Solutions
mol (amount of solute)
V (volume of solution)

32. Step 1: Find molar mass! Too much typing so….C6H12O6 molar mass=180 g/mol
Step 2: Convert mg to g. 120 mg = .12 g
Step 3: Convert grams to moles of glucose
.12÷ 180=6.67×10-4 moles
Step 4: Plug in numbers! (don’t overlook deciliter)
M= = 6.67 × 10-3 M
Real Test Question
moles
liter
6.67×10-4 moles glu
0.1 L

33. Step 1: convert the equation! M=mol/L  mol=(M)(L)
Step 2: Plug in numbers. (1.56 M)(2.85 L)=4.45 moles
Step 3: The problem asked for mass, so change moles to grams
(MUST find formula weight of CaCl2…. 111 g)
(4.45)(111)= 493 grams
Real Test Question

34. Dilution problems involve two solutions: a more concentrated one,
higher molarity, and a diluted one, lower molarity.
The moles of the solute will not change but the Molarity will so
watch out!
(M1V1)=(M2V2)
Dilutions of Concentrations

35. Step 1: Plug in numbers. M1V1=M2V2 (.25)(x)=(.1)(350)
Step 2: Calculate. V1=140 mL
These are not hard questions so don’t let them trip you up over
tiny conversions or anything.
Real Test Question
How much of a 0.250 M is needed to make 350 mL of 0.100 M solution?

36. -identify polarity
-draw lewis structures of compounds
-identify resonance structures
-write formulas and names
-determine geometric arrangement of molecules
-find polarity of molecules
-use valence bond theory to find hybrid orbitals and modle
sigma and pi bonds
Unit 4 Overview

37. Polarity Explained
In Covalent Bonds, atoms form compounds by sharing electrons.
-if electrons are unequally shared, it is called a polar covalent bond
-if electrons are equally shared, it is called a non-polar covalent bond
Polar covalent bonds have diploe moments where one side is more positive
while the other is more negative
BIG difference between a molecule having polar bonds and molecule being polar

38. Bonding Continuum
The electronegativity difference (ΔEN), is determined by subtracting one
EN value from the other:
F2 ΔEN = 4.0 - 4.0= 0 ΔEN <0.4, nonpolar
HCl ΔEN = 3.0 - 2.1= 0.9 ΔEN < 2.0, polar
HF ΔEN = 4.0 - 2.1= 1.9 ΔEN ≥ 2.0, ionic
LiF ΔEN = 4.0 - 1.0= 3.0

39. The polarity of the molecule affects its properties.
A molecule can have polar bonds and still not be polar due to the
arrangement of the polar bonds around the central atoms.
It will NOT be polar if: the bonded atoms are arranged symmetrically, and
the bonded atoms are identical or have similar electronegativity's
Polarity Explained
Example: CO2
CO2 is nonpolar- the dipole moments
cancel each other out due to the
symmetry of the molecule
Example: H2O
H2O is polar- the lone pairs of electrons
on the oxygen atom cause an
asymmetrical distribution of the dipole
moments

40. Step 1: Sum the valence electrons from all atoms, including ion charge
Step 2: Write symbols for the atoms to show how atoms are attached, and
connect them with a single bond
Step 3: Complete the octets around all the atoms bonded to the central atom
Step 4: Place any left over electrons on the central atoms
Step 5: If there are not enough electrons to give the central atoms an octet,
try forming multiple bonds
Steps to Drawing Lewis Structures

41. 1. Molecules and polyatomic ions containing an odd number of electrons.
EX: NO BOTH are acceptable
2. Molecules and polyatomic ions in which an atom has fewer than an
octet of valence electrons.
3. Molecules and polyatomic ions in which an atom has more than an
octet of valence electrons.
The octet rule will fail in three situations:

42. Explanation of Failures
Fewer than eight electrons Example: BF3
If boron were given a filled octet, a negative charge would be put on boron and a
positive charge on fluorine, which is not an accurate distribution of electrons in BF3
Structures that put a double bond between boron and fluorine are less important
(less likely) than the one that leaves boron with only 6 valence electrons.
If filling the octet of the central atom results in a negative charge on the central
atom and a positive charge on the more electronegative outer atom, don’t fill the
octet of the central atom.

43. More than eight electrons Example: PO4
HINT: When the central atom is on the 3rd row or below and expanding its
octet eliminates some formal charges, this structure will be preferable.
There is a Lewis structure that only has 8 electrons around the central phosphorous,
BUT and better structure puts a double bond between the phosphorus and one of the
oxygen atoms, eliminating the charge on the phosphorus and the charge on one of the
oxygen atoms.
Explanation of Failures

44. Occur when there is more than one Lewis structure for a molecule that
differs only in the position of electrons. The actual molecule is a combination
of all the possible resonance forms.
A molecule exhibits resonance if it has identical outer atoms with non-identical
bonds OR if you can draw 2 or more Lewis structures
Rules:
1. Resonance structures must have the same connectivity (only electron positions
can change).
2. Resonance structures must have the same number of electrons.
3. Second row elements have a maximum of eight electrons (bonding and
nonbonding). The third row can have expanded octet
4. Formal charges must total the same.
Resonance Structures

45. OZONE: O3
Resonance Structure Examples
CARBONATE: CO3
2-
NITRATE: NO3
-

46. Evaluate the Formal Charge: the number of valence electrons in an atom,
minus the number of electrons assigned to the atom
Rules for the “BEST” Structure:
1. The sum of all formal charges in a neutral molecule must be zero.
2. The sum of all formal charges in an ion must equal the charge of the ion.
3. Smaller formal charges on individual atoms are better than large ones.
4. When formal charge cannot be avoided, negative formal charge should
reside on the most electronegative atom.
Resonance Structures
How do you know which structure is the “best” or “most correct”??

47. Formal Charge of NCS
N: 5e-–(4+2)e- = -1
C: 4e-–(4)e- = 0
S: 6e-–(4+2)e- = 0
N: 5e-–(6+1)e- = -2
C: 4e-–(4)e- = 0
S: 6e-–(2+3)e- = +1
N: 5e-–(2+3)e- = 0
C: 4e-–(4)e- = 0
S: 6e-–(6+1)e- = -1
The BEST structure would be the one outline in red.
It has the smallest formal charges on individual atoms, AND the
negative formal charge is on the most electronegative atom!

48. For the most part, you must draw out the Lewis structures to
see if a molecule exhibits resonance –but some can be ruled
out without drawing the Lewis, saving time.
Here’s how:
-You can never resonate into hydrogen or halogens. So if those are the
only outer atoms, no resonance.
-Molecules with just 2 atoms can’t have resonance
-Since H’s and halogens can’t resonate, molecules where all but 2 atoms
are H’s or halogen can’t have resonance.
Tips for Resonance Structures

49. Real Test Question
Which of the following is the best Lewis structure a resonance structure?
A. SO3 B. CO C. CH4 D. PF5 E. NH3
1. We learned you can NEVER resonate with hydrogen or halogens, so
that eliminates C, D, and E
2. Molecules with just 2 atoms can’t have resonance, eliminating B
3. Now to look at the structure for SO3

50. C often bonds to carbons, hydrogen, nitrogen, oxygen, and halides.
Rules to forming bonds in organic structures:
-Carbon has four bonds and zero lone pairs.*
-Hydrogen has one bond and no lone pairs.
-Nitrogen has three bonds and one lone pair.
-Oxygen has two bonds and two lone pairs.
-Halogens has one bond and three lone pairs.
Exceptions: cyanide ion and carbon monoxide have a lone pair of electrons
Organic Molecules

51. Organic Compounds
(Functional Groups to Memorize)
Alcohols:
Ethers:
Esters:
Amines:
Carboxylic acids:
R—OH
R—O—R
R—C—OR
R—N—R
R—C—OH
O
R
O

52. Alkanes: only single carbon-carbon bonds
Formula: CnH2n+2
Alkenes: contain one or more double carbon-carbon bond(s)
Formula: CnH2n
Alkynes: contain one or more triple carbon-carbon bond(s)
Formula: CnH2n-2
Aromatic: carbon atoms connect in a planar ring joined by alternating
single and double bonds between carbon atoms.
Hydrocarbon Basics

53. Molecular formulas show the number and type of atoms in the molecule, but
they do not show how the atoms are attached
Butane: C4H10
Structural formulas show the bonding pattern in the molecule
Condensed structural formulas indicate the atoms attached to each carbon.
CH3CH2CH2CH3
Hydrocarbon Basics

54. Saturated
Contain the maximum number of
hydrogen atoms possible on the
carbon (no double bonds)
-Alkanes
-Cyclic Alkanes
Saturated VS Unsaturated
Hydrocarbons
Unsaturated
Contain fewer hydrogen atoms than
alkanes having the same number of
carbon atoms
-Alkenes
-Alkynes
-Aromatic hydrocarbons
-Cyclic alkenes

55. Naming Hydrocarbons
There are three essential components of a
name:
(prefix)+(base)+(suffix)
Prefix: Indicates position, number, and type of
branches or functional group
Base: indicates the length of the longest carbon
chain or ring
Suffix: indicates the type of functional group
(-ane, -ene, -yne)

56. Naming Chart for Hydrocarbons
Alkanes Alkenes Alkynes
N= prefix CnH2n+2 CnH2n CnH2n-2
1 -meth CH4 methane CH2 methene CH methyne
2 -eth C2H6 ethane C2H4 ethene C2H2 ethyne
3 -pro C3H8 propane C3H6 propene C3H4 propyne
4 -but C4H10 butane C4H8 butene C4H6 butyne
5 -pent C5H12 pentane C5H10 pentene C5H8 pentyne
6 -hex C6H14 hexane C6H12 hexene C6H10 hexyne
7 -hept C7H16 heptane C7H14 heptene C7H12 heptyne
8 -oct C8H18 octane C8H16 octene C8H14 octyne
9 -non C9H20 nonane C9H18 nonene C9H16 nonyne
10 -dec C10H22 decane C10H20 decene C10H18 decyne

57. Electron Domain Geometries
This arrangement counts both the
atoms AND the lone pairs as one
electron domain on the central atom
VSEPR Model
Molecular Geometries
This arrangement ONLY accounts for the
position of the atoms, not the lone pairs.
Within each electron domain geometry, there
are more than one molecular geometry
ABnE
A= central atom
B=bonding atom
n=number of B atoms bonding to A
E=number of lone pairs
BrF3 would be written as: AB3E2

58. VSEPR Models
AB2 Linear AB2 Linear
AB3 Trigonal planar AB3 Trigonal planar
AB2E Bent
AB4 Tetrahedral AB4 Tetrahedral
AB3E Trigonal pyramidal
AB2E2 Bent
AB5 Trigonal bipyramidal AB5 Trigonal bipyramidal
AB4E Seesaw
AB3E2 T-shaped
AB2E3 Linear
AB6 Octahedral AB6 Octahedral
AB5E Square pyramidal
AB4E2 Square planar
E.D. Geometry Molecular Geometry

59. VSEPR Table
#
of
ED
E Domain
Geometry
#
Bond
atoms
# lone
pairs
Molecular
Geometry
Bond
Angles
Polarity Example Shapes
2 Linear 2 o Linear 180 nonpolar
3 Trigonal
planar
3 0 Trigonal
planar
120 nonpolar
2 1 Bent <120 polar
4 Tetrahedral 4 0 Tetrahedral 109.5 nonpolar
3 1 Trigonal
pyramidal
<109.5 polar
2 2 Bent <109.5 polar

60. # of
ED
E Domain
Geometry
Bond
Pairs
# Lone
Paris
Molecular
Geometry
Bond
Angles
Polarity Example
5 Trigonal
bipyramidal
5 0 Trigonal
bipyramidal
120
90 axial
nonpolar
4 1 Seesaw <120
<90
axial
polar
3 2 T-shaped <90 polar
2 3 Linear 180 nonpolar
6 Octahedral 6 0 Octahedral 90 nonpolar
5 1 Square
pyramidal
<90 polar
4 2 Square planar 90 nonpolar

61. Predicting Polarity using Shape
ALWAYS POLAR SHAPES: bent, trigonal pyramidal, seesaw,
T-shaped, square pyramidal
If all bonding atoms are the SAME: linear, trigonal planar,
tetrahedral, trigonal bipyramidal, and octahedral molecular shapes
will be nonpolar
For a molecule to be polar it must:
1. Have polar bonds (EN difference/dipole moments)
2. Have an asymmetrical shape

62. Valence Bond Theory describes the formation of covalent bonds due
to the overlap of half filled valence atomic orbitals in two different
atoms.
Two types of overlapping orbitals: sigma (σ) and pi (π) bonds
Sigma bonds occur when orbitals of two electrons overlap head-to-head
Pi bonds occur when two orbitals overlap when they are parallel
Single bonds = one sigma bond
Double bonds = one sigma and one pi bond
Triple bonds = one sigma and two pi bonds
Valence Bond Theory

63. Hybrid Orbitals/Hybridization
ED
Geometry
# of atoms +
lone pairs
Hybridization Examples
Linear 2 sp CO2 BCl2
Trigonal
Planar
3 sp2 NO2
- BCl3
Tetrahedral 4 sp3 H2Te CBr4 AsCl3
Trigonal
bipyramidal
5 sp3d AsCl5 SF5 ICl3
Octahedral 6 sp3d2 IF5 SF6