Homework for Monday
• Finish reading Chapter 3(sections 6-8)
• Chap. 3: 26,28,32,34,42,46
Outline
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Types of Chemical Bonding
Lewis Dot Symbols
Naming Ions and Ionic Compounds
Lewis Structures for Covalent Bonding
First a word on the Periodic Table
Three Common Covalent
Compounds
Types of Chemical Bonding
• Covalent
– Sharing of electrons
– Example
• Table sugar(sucrose)
• Ionic
– One atom has more
than its share of
electrons
– Example
• Table salt (NaCl)
Ammonia
NH3
Water
H2O
Methane
CH4
Lewis Structures and Ionic
Compounds
• Atoms are combined so that their “valence
shell is filled”
– Either by loss or gain of electrons
• Generally follows the octet rule:
Na+ (95 pm)
Na+
F- (136 pm)
– “Whenever possible the valence electrons are
distributed so that eight electrons surround
each main group element (except Hydrogen)”
Ionic bond
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Naming Simple Compounds
• Simplest type are binary ionic (Type I)
– Metals from Groups I, II and III and those that only
form a single cation (Ag and Zn)
– Non-metals from right side of Periodic Table
– Examples include NaCl, KBr, LiF, AlCl3
– Charge Neutrality must be maintained!
• Cation name first, anion second
• Anion: the suffix -ide replaces the last syllable
of its name
• KBr = potassium bromide
• GaN = gallium nitride
Ionic Compounds With
Polyatomic Ions
• Memorize table 3.5 in text
• Most common polyatomic cations are
ammonium (NH4+) and
hydronium(H3O+)
Important Polyatomic Anions
NO2- nitrite
NO3- nitrate
SO32- sulfite
SO42- sulfate
HSO3- hydrogen sulfite
OH- hydroxide
O22- peroxide
CN- cyanide,
PO43- phosphate
HPO42- hydrogen
phosphate
– H2PO4- dihydrogen
phosphate
–
–
–
–
–
–
–
–
–
–
Examples of Compounds with
Polyatomic Ions
• Ba(NO3)2
• KMnO4
• H2O2
• (NH4)2Cr2O7
• LiClO4
• NaClO
• barium nitrate
• potassium
permanganate
• hydrogen peroxide
• ammonium
dichromate
• lithium perchlorate
• sodium hypochlorite
–
–
–
–
–
–
–
–
–
–
–
–
CO32- carbonate
HCO3- hydrogen carbonate
ClO- hypochlorite
ClO2- chlorite
ClO3- chlorate
ClO4- perchlorate
HCO3- hydrogen carbonate
C2H3O2- acetate
MnO42- manganate
MnO4- permanganate
Cr2O72- dichromate
CrO42- chromate
Naming Acids
• No oxygen in acid
– Anions include
• Halides (F-, Cl-, Br-, I-)
• Cyanide (CN-)
• Sulfide (S2-)
• Names contain hydro
–
–
–
–
–
–
Hydrochloric (HCl)
Hydrofluoric (HF)
Hydrobromic (HBr)
Hydroiodic (HI)
Hydrocyanic (HCN)
Hydrosulfuric (H2S)
• Oxygen in acid
– Anions include
•
•
•
•
•
•
Nitrate
Nitrite
Sulfite
Sulfate
Phosphate
Acetate
• Names include
•
•
•
•
•
•
Nitric (HNO3)
Nitrous (HNO2)
Sulfurous (H2SO3)
Sulfuric (H2SO4)
Phosphoric (H3PO4)
Acetic
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Lewis Structures and Covalent
Compounds
• Atoms are combined so that their “valence
shell is filled”
• Generally follows the octet rule:
– “Whenever possible the valence electrons are
distributed so that eight electrons surround
each main group element (except Hydrogen)”
Non-Ionic Compounds???
Again – tend to form Noble gas configurations
But now Sharing valence shell e-’s
Examples
1) F2: F
7 valence e-’s, needs 1 e-, but can’t be transfer, must share!!
F
F
Both have Ne configuration
2) O2: 6 valence e-’s, needs 2
O
valence
electron
represented
by a dot
O
Again, both have Ne configuration
Demonstration
Polyatomics
OF2
F O F
OH2
H O H
Ne configurations, each F share 2 e-’s
O shares 4 e-’s
Now Ne & He configurations,
H shares 2 e-’s, O shares 4 e-’s
Multiple Bonding
• Sometimes it is necessary for elements to share
more than one pair of electrons to complete an
octet.
• We can represent this with Lewis structures by
putting two or three electron pairs between the
atoms.
• This is a good model because it helps to define
the molecular geometry
• Multiple bonds are shorter and stronger than
single bonds.
What unit is 10 -10m?
The Angstrom
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Calculating Formal Charges
• The formal charge on an atom is given by
the following equation:
F=V-L-S/2
where F is the formal charge
V is the atom’s number of valence electrons
L is the atom’s number of lone pair electrons
S is the number of electrons shared in
chemical bonds
Formal Charges
• When some molecules are formed, the constituent atoms end with
either a greater or fewer number of electrons due to the formation of
multiple bonds
• Carbon monoxide is a classical and simple example of this situation
– There are a total of 10 valence electrons available
– In order to satisfy the octet rule for both atoms three pairs of electrons
need to be shared
-1 +1
:C:::O:
Although formal charges have no real meaning for covalent compounds
(it does not make them ionic), it does give us a sense of where chemical
interactions might occur-wait until lab next semester!
When writing possible structures, the most likely structure is the one
that has the least separation of charge (fewer formal charges)
Building Lewis Structures for More
Complex Molecules
• Draw out molecular skeleton
• Add up the total number of valence electrons available (A)
• Determine the number of electrons needed (N) by each atom to
complete its octet
• The number of electrons that has to be shared(S) is S=N-A
• Assign one pair of electrons to each of the bonds that needs to be
formed
• If any shared electrons remain, they should be used in forming
double or triple bonds
• Assign the remaining electrons as lone pairs to the atoms giving
octets to all atoms except hydrogen
• Determine the formal charge on each atom and write it next to the
atom
• Make sure that the sum of the formal charges is equal to the overall
charge on the molecule or molecular ion
Examples
• thiosulfate ion (S2O32-)
• phosphorus trichloride (PCl3)
• nitrate ion (NO3-)
4
Problems
“So What Do I Do If The Rules Don’t Work?”
• Remember that some elements are electron deficient (Be, B)
• “The model says that I need to share four pairs of electrons, but
there are five atoms that I need to form bonds to!”
• Change the rule to fit your molecule
– Usually involves expanded octets or odd numbers of electrons
• Most of the time Lewis structures provide possible models for more
complex molecules, but you have to be careful
• Keep charge separation in the molecule to a minimum
• This only provides a “simple” model for chemical bonding and there
are often many exceptions to the rules
– When a model has several resonance structures, the real structure
probably falls somewhere in the middle
• phosphorus pentachloride (PCl5)
• nitric oxide (NO)
• [CH3]• BeCl2
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