1. What types of elements combine to form ionic and covalent compounds?
2. What factors determine the shape of covalent molecules?
3. What is a polar molecule?
4. Can you identify the Lewis Dot
Structures for ionic and covalent compounds?
5. How do you determine the type of bonding using electronegativity?

• Ion -atom with a charge
• Cation -Positive ion, lost electrons
• Anion -Negative ion, gained electrons
• Oxidation number -the charge that represents the number of electrons lost or gained
• (New) Polyatomic ion -more than one element attached to the charge.
23
Na

1
11

• Bonds – the attractive force between atoms or ions in a compound.
• A bond depends upon:
• The electron configuration (involves valence electrons)
• Electronegativity
• Why do elements bond?
• To achieve a stable electron configuration (8 electrons; noble gas configuration)
» Octet rule – atoms lose, gain or share electrons to achieve a stable configuration of eight valence electrons
• To achieve the lowest possible energy state
(lowest potential energy)

FAMILY SINGLE ELEMENT BONDED ELEMENT
IA X X
IIA X X
IIIA X X
IVA X X
VA through VIIIA the bonded element structure is the same as the single element

• Ionic – complete transfer of electrons
(loses or gains); example: NaCl
• Nonpolar covalent (also called pure covalent) – equal sharing of electrons between atoms; example: O
2
• Polar covalent – unequal sharing of electrons between atoms; example: H
2
O
• Electronegativity (EN) – indicates how strongly an atom wants to gain an electron (Decreases down, Increase Across)

• Occurs when electrons are completely transferred from one atom to another.
• The strongest type of bond.
• Formed between a metal and a nonmetal.

http://www.beyondbooks.com/psc92/3b.asp

• Lose e- to form same stable configuration as Noble Gas in preceding period ; forms ions.
Na Na + + 1e -
1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6
Same Configuration as: Ne atom
1s 2 2s 2 2p 6

• Gain e- to form same configuration as Noble Gas at end of the same period; forms ion.
Cl + 1e Cl 1-
1s 2 2s 2 2p 5 3s 2 3p 5 1s 2 2s 2 2p 6 3s 2 3p 6
Same Configuration as: Ar
1s 2 2s 2 2p 6 3s 2 3p 6

• Make sure you understand valence electrons and electron configurations
• Draw Dot structures of valence electrons
• Know your oxidation numbers
• Make sure that positive and negative charges add up to zero!
http://wine1.sb.fsu.edu/chm1045/notes/Bonding/Ionic/Bond02.htm

 Equation showing an ionic bond
 Metals give up e-
 Non-metals gain e-
 Must get 8e- total
 Na• + F  Na +1 F -1
 Na has no dots and F has all 8 dots
 You must show the charges ( will equal 0 )
 One Na +1 ion and one F -1 ion form a formula unit (NaF)

• The simplest ratio of the ions represented in an ionic compound is called a formula unit
• We use this because no single particle of an ionic compound exists
• Total # of e- gained by nonmetals atoms = total # of e- lost by the metal,so overall charge = 0

• Examples:
KBr Potassium bromide (1:1)
MgCl
2
Magnesium chloride (1:2)
• Practice:
Sodium phosphide, three sodium ions for every phosphide ion
Na
3
P

Practice by writing the equations for the following ions:
1. Magnesium & Sulfur
2. Aluminum & Oxygen
3. Magnesium & Iodine
4. Copper & Bromine (note: Copper can be +1 or +2)

• The result of sharing valence electrons.
The shared electrons are part of the complete outer shell of both atoms.
– Occurs when elements are close together on the periodic table
– Between nonmetallic elements
• Molecule -formed when two or more atoms bond covalently.

• Nonpolar covalent (also called pure covalent) – equal sharing of electrons between atoms; example:
O
2
• Polar covalent – unequal sharing of electrons between atoms; example: H
2
O

• Can exist as gases, liquids, or solids depending on molecular mass or polarity
• Usually have lower MP and BP than ionic compounds
• Do not usually dissolve in water
• Do not conduct electricity


NONPOLAR COVALENT
 e- are equally shared

No difference in electronegativity

All diatomic molecules are nonpolar covalent

H
2
,I
2
,O
2
,Cl
2
,N
2
,Br
2
,F
2

POLAR COVALENT

 unequal sharing

E- spends more time with the more electronegative atom

Difference in EN = 0.1-1.7
 regions of partial charges
 also known as a dipole (two poles) d d
Partial positive
H Cl
Partial negative pulled more by chlorine

• Two Hydrogen Atoms (H
2
) http://web.jjay.cuny.edu/~acarpi/NSC/5-bonds.htm

• Unequal sharing of electrons
• Have poles (dipoles) – regions that are positive & regions that are negative
• Electrons are pulling toward more electronegative element
• Symbols:
δ + δ : show regions of partial charge
: arrow points to more electronegative element

• Electrons are equally shared
• No difference in electronegativity
(ex. diatomic molecules)
• Also, can be due to shape of molecule
– Electrons pulled equally in all directions, polar effect cancels
(ex. I Be I )

• Look at table of EN values and subtract the values for the 2 atoms involved in the bond
• take the absolute value
• If EN difference is…
0 – 0.4  nonpolar covalent bond
0.5 – 1.6  polar covalent bond
>1.7  ionic bond

0 = NONPOLAR COVALENT
>0 TO 1.6 = POLAR COVALENT
1.7 TO 3.4 = IONIC
POLAR
COVALENT
IONIC
0 1.7 3.4
NONPOLAR

• Use VSEPR (valence shell electron pair repulsion) rules:
1) Draw the Lewis dot structure for the molecule
2) Identify the central atom
3) Count total # of electron pairs around the central atom
4) Count # of bonding pairs of electrons around the central atom
5) Count # of lone pairs of electrons around the central atom
6) Look at summary chart, identify shape

# of e- pairs around central atom
# of bonding pairs of e-
# of lone pairs of e-
Name
2 2 0 linear
3 3 0 trigonal planar
4 4 0
4
4
4
3
2
1
1
2
3 tetrahedral trigonal pyramidal angular
(bent) linear
Shape

Determine the shape:
1. BCl
3
2. CH
4

• TO DETERMINE THE POLARITY OF A
MOLECULE , not a bond , you must know the type of bond and the shape.
• POLAR MOLECULES must meet 2 criteria:
• Must have a polar covalent bond (EN difference between 0.5 and 1.6) AND
• Must have an asymmetrical shape: trigonal pyramidal, angular, or 2 element linear
• If both criteria are not met , it is not a polar molecule, it is either a nonpolar molecule or an ionic compound.

Label the following MOLECULES as polar or nonpolar.
1. NH
3
2. CH
4
3. HCl

• Atoms of some elements attain a noble-gas configuration by sharing more than one pair of electrons between two atoms
• When writing structural formulas a line can represent a pair of shared electrons

• Double Bond: share 2 pairs of e-
O=O
• Triple Bond: share 3 pairs of e-
N N

• Beryllium (2 valence e-, full with 4 valence e-) BeI
2
• Aluminum (3 valence e-, full with 6 valence e-) AlCl
3
• Boron (3 valence e-, full with 6 valence e-) BH
3

• Odd number of valence e-, cannot form an octet around each atom
– NO
2
• Compounds can form with fewer than
8 e- (rare)
– BH
3
• Expanded octet – can hold more than
8 valence e- due to empty d orbital
– S and P most common elements
• SF
6 or PCl
5

• Some atoms bond so they have more than 8e in the outer level
• Occur only around a central nonmetallic atom from period 3 or higher
PCl
5
SF
6 http://www.wou.edu/las/physci/poston/ch222/VSEPR-Geometries.htm

PCl
3
Cl
Cl
Cl
Cl
PCl
5
Cl
Cl
Cl

Cl
SCl
2
SCl
6
Cl Cl
Cl
Cl
Cl Cl

Property
Electrons are:
Difference in
Electronegativity:
Bond between:
State at room temperature
Particle Name:
Melting Point:
Conducts Electricity?
Dissolves in water?
Flammable
Forms a more stable configuration?
Examples:
Ionic Bonds
Transferred
≥1.8
A metal and nonmetal
Solid
Formula Unit
High
Yes
Yes
No
Yes
NaCl, MgI
2
Covalent Bonds
Shared
0-nonpolar
>0 to 1.7-polar
2 nonmetals
Usually gas
(can be solid or liquid)
Molecule
Low
No
No (usually)
Yes
Yes
NH
3
, CHCl
3

• Metals often form lattices in the solid state similar to ionic crystals
• Even though metal atoms have at least one valence ethey do not share or lose electrons
• Electron sea model – all the metal atoms in a metallic solid contribute their valence eto form a “sea” of e-

• Positive ions in a “sea of electrons,” belongs to the crystal as a whole
• Force applied (left to right) highlighted cation unchanged.
• Explains ease of deformation of metals http://cwx.prenhall.com/petrucci/medialib/media_portfolio/12.html