2, 12, 14, 20, 22, 26-
32, 36, 38, 48-58,
62, 66-74

• Compounds and Molecules are held together by chemical bonds
• Three types of bonds
– Ionic
• Metals and non-metals
– Covalent
• Non-metal and Non-metal
– Metallic
• Between atoms of metals

• All atoms strive to have electronic configurations like the Noble Gases
• Eight electrons in the outermost shell, highest principle quantum number (n)
• Except H and He follow duet rule
– Want two electrons in outermost shell
• How do the atoms achieve an octet?

• Ionic Bonds
– Atoms take or give electrons from other atoms
• Covalent Bonds
– Atoms share electrons between themselves
• Metallic Bonds
– Sea of electrons

n=1 n=2 n=3 n=4 n=5 n=6 n=7
1
1
1
1
1
1
1
2 3 4 5 6 7
2
8
2
2
2
2
2
3
3
4
4
5
5
6
6
3
3
4
4
5
5
6
6
3 4 5 6
7
7
8
8
7
7
8
8
7 8

• Symbol represents the nucleus and all electrons except for those in the valence shell
• Give the Lewis Dot Structure for:
Na F O 2-
• Species with the same number of electrons are isoelectric
O 2F Ne Na + Mg 2+
How many electrons does each species have?

• GN Lewis developed the theory of covalent bonding
• Structures showing covalent bonds are called Lewis structures
• Each line represents a shared pair of electrons (2 electrons)
• Lone pairs of electrons are shown by a pair of dots

• Decide on atom connectivity and placement
– Hydrogen (never in the middle) is frequently bonded to oxygen
– Oxygen is rarely the central atom
– Oxygen will not bond to oxygen (except O
2
– Carbon will be the central atom
– Least electronegative atom is in the middle or O
3
)

• Count the total number of valence electrons
– An atom’s number of valence electrons is equal to its group number
• Determine the total number of shared electrons electrons needed – valence electrons present
• Connect the atoms with single bonds
– A single bond is one shared pair of electrons
• Use lone pairs and/or multiple bonds to give each atom an octet of electrons

• Draw Lewis Structures for:
• H
2
O
• HCl
• NH
3

• CO
2
• N
2

• Definition : Ions are atoms or groups of atoms with an electrical charge
• Cations : are positively charged, due to loss of electrons (Metals)
• Anions : are negatively charged, due to gain of electrons (Non-Metals)
• Number of electron’s gained or loss is due to atoms wanting Octet

• Na
• Ra
• Al
• Se
• O
• Cl
• F

• Ionic compounds are held together by ionic bonds , or the attraction of oppositely charged ions
• In the solid state, ionic compounds form crystalline lattices
– Cations are attracted to all the neighboring anions, not just one
– Thus, there are no discrete ionic “molecules”
Ball and Stick Model

• Most transition metals form more than 1 cation
+1 only Ag +
+2 only Zn 2+ , Cd 2+
+1 and +2 Hg
2
2+ , Cu +
Hg 2+ , Cu 2+
+2 and +3 Cr 2+ , Fe 2+ , Co 2+
Cr 3+ , Fe 3+ , Co 3+
+2 and +4 Sn 2+ , Pb 2+
Sn 4+ , Pb 4+

NO
3
nitrate
NO
2
nitrite
CO
3
2carbonate
HCO
3
-
Bicarbonate
Or hydrogen carbonate
SO
4
2sulfate
SO
3
2sulfite
NH
4
+ ammonium
OH hydroxide
PO
4
3phosphate
HPO
4
2monohydrogenphosphate
H
2
PO
4
dihydrogenphosphate
C
2
H
3
O
2
acetate

• The net charge on a formula unit must be zero
S
(+) charges =
S
(-) charges
• Since there are no ionic “molecules” the formula of an ionic compound is the simplest ratio of cation to anion that gives an electrically neutral combination
Al 3+ and O 2-
Ca 2+ and O 2-

• Write the formula for each of the following pairs of ions
• Na and Oxygen
• Mg and Fluorine
• Rb and Iodine

• Rules for naming compounds and molecules
• Anions
– Name the element, drop the ending leaving the root and add “ ide ”
• Element – root + “ide”
• Cl
• O
• N
• S
• I

1. Name the cation by naming the element
• If the cation is a transition metal you need to distinguish the charge using Roman
Numerals
• Fe 2+ is named Iron (II)
• Pb 4+ is named Lead (IV)
2. Name the anion
• Can be an elemental anion or polyatomic
3. Combine them as two words

K
2
O
Li
2
CO
3
K
2
SO
4
NaHCO
3
Cr
2
O
3

• What are the formulas of these compounds?
calcium sulfide iron (III) acetate
Chromium (III) sulfate

• Name each element
• Indicate how many of each element is present with a prefix multiplier
– Mono =1; di =2; tr i=3; tetra =4; penta =5; hexa =6; hepta = 7; octa = 8; nona = 9; deca
= 10
• Add the suffix “ide” to the last element
• The prefix multiplier mono is left off of the first element in the compound

• IBr
• NI
3
• N
2
O
4

• Sulfur dioxide
• Diphosphorous pentoxide
• Carbon tetrachloride

Molecular Compounds: Common Names
• These compounds have common (non-systematic names)
– Water (H
2
O)
• Dihydrogen monoxide
– Ammonia (NH
3
)
• Nitrogen trihydride
– Methane (CH
4
)
• Carbon tetrahydride
– Nitrous oxide (N
2
O)
• Dinitrogen monoxide
– Hydrazine (N
2
H
4
)
• Dinitrogen tetrahydride

• Acids are compounds that can donate an hydrogen ion (H + ion)
• Acids fall into two categories
– Binary Acids HX
– Oxoacids HXO n
• Polyatomic anions

• Most binary acids result from dissolving the corresponding molecular compound in water
• Binary acids are named as hydro ( stem name of X ) ic acid
HCl
(g)
HCl
(aq)
HCN
(g)
HCN
(aq)

• Oxoacids are named based on the oxoanion
• “Ate” anion => ic acid
• “Ite” anion => ous acid

Polyatomic Anion
CO
3
2(carbonate anion)
NO
2
(nitrite anion)
NO
3
(nitrate anion)
PO
4
3(phosphate anion)
SO
3
2(sulfite anion)
SO
4
2(sulfate anion)
H
2
CO
3
(carbonic acid)
HNO
2
(nitrous acid)
HNO
3
(nitric acid)
H
3
PO
4
(phosphoric acid)
H
2
SO
3
(sulfurous acid)
H
2
SO
4
(sulfuric acid)

• We can write formulas
• We can name compounds and molecules
• We can draw Lewis Structures
– But what do these molecules look like?

• VSEPR: Valence Shell Electron Pair
Repulsion
• Like charges repel and want to be as far apart as possible
• Therefore a given combination of electrons will form into a specific shape

1. Draw the Lewis Structure
2. Assign the central atom (A)
3. Determine the number (n) of atoms bonded to (A) designate them (X n
)
4. Determine the number of lone pairs on
(A) designate them (E m
)
5. Put together the AX n
E m notation

X + E = 2
X + E = 3
X + E = 4
X + E = 5
X + E = 6

• What is the geometry of
• CO
2
• BF
3
• H
2
O
• NH
3

• Linus Pauling developed the electronegativity scale
• Electronegativity is a measure of an atom’s affinity for electrons
• Fluorine is the most electronegative element (EN=4.0)
• The closer an atom is to fluorine, the more electronegative it is

• If two atoms of identical electronegativity are bonded together, the bond is non-polar
• If two atoms of different electronegativity are bonded together, the bond is polar , and the electrons spend more time around the more electronegative atom
– This creates partial charges
• The greater the difference in EN between two atoms, the more polar the bond
– The limiting example of this is the ionic bond


EN
0.0
0.1 – 0.4
0.5 – 1.4
1.5 – 3.2
Type of Bonding
Pure covalent bond
(equal sharing of e ‘s)
Non-polar covalent bond
(almost equal attraction for shared e pairs)
Polar covalent bond
(unequal sharing of e ’s)
Ionic bond (e transfer)

• The bond in hydrogen is
• The bond in hydrogen chloride is

• Bond dipoles are vectors
• The vectoral sum of the bond dipoles gives the molecular dipole
• Based on the shape of the molecules you can predict if the dipoles will cancel each other or if they will create a dipole moment
• If a dipole moment exists then the molecule is said to be polar
• If no dipole moment exists then the molecule is said to be non-polar

• Is carbon dioxide polar or non-polar?
• Is water polar or non-polar?
• Is boron trifluoride polar or non-polar?

• These are attractive forces between molecules or atoms or ions
• Immensely important
– These forces hold DNA molecules in a helix and and are the mechanism for DNA transcription

• This is the attraction between the opposite
(partial) charges of polar molecules d +
H Cl d d +
H Cl d -

• This is generally stronger than dipolar attractions
• Hydrogen bonding occurs between a hydrogen atom and O, N or F.
• For H-bonding to happen the H must be directly bonded to a O, N or F.
O—H O—H
H This is an attraction not really a bond
H

• Also called Van der Waal’s forces, these are created by instantaneous dipoles
• London forces are much weaker than either dipole-dipole or H-bonding
• London forces get stronger with larger atoms/molecules

London Forces Between Helium Atoms d-
He d+ d-
He d+ d+

• This is the attraction between an ionic charge and a polar molecule
• This attraction allows ionic solids to dissolve in water
• The strength of this force varies widely and depends on the magnitude of the dipole moment of the polar species and the size of the ion

A Sodium Ion and a Chloride Ion
Hydrated by Water Molecules

• More intermolecular forces mean:
– Higher boiling and melting points
– More viscous liquids
• IM Forces also affect solubility
– ‘like dissolves like’

SnH
4
, CH
4
, GeH
4
, SiH
4
HBr, HI, HCl, HF

H
2
O
H
2
Se
H
2
S
H
2
Te

• Is carbon dioxide soluble in water? Explain

• Are ionic compounds more soluble in water or in gasoline (a non-polar solvent)?
Explain