Laboratory 06
MOLECULAR GEOMETRY AND POLARITY
Background- Lewis structure
Diagrams that show the bonding between atoms of a
molecule and the lone pairs of electrons that may
exist in the molecule
A Lewis structure can be drawn for any covalently
bonded molecule, as well as coordination
compounds.
Construction of Lewis Structures
Two Rules
1. Total # of valence electrons – the total number
of valence electrons must be accounted for, no
extras, none missing.
2. Octet Rule – every atom should have an octet
(8) electrons associated with it.
Hydrogen should only have 2 (a duet).
Determining the number of valence
electrons
Full d-orbitals do not count as valence electrons.
They belong to the inner shell.
For example:
Pb
[Xe]4f145d106s26p2
This is four (4) valence electrons.
The 5d is part of the inner shell (n=5) which is full.
The total number of available valence electrons is just the
sum of the number of valence electrons that each atom
possesses (ignoring d-orbital electrons)
For H2O,
The total number of valence electrons = 2 x 1 (each H is 1s1) +
6 (O is 2s22p4)
=8
For CO2
The total number of valence electrons = 4 (C is 2s22p2) +
2 * 6 (O is 2s22p4)
= 16
Central Atom
In a molecule, there are only 2 types of atoms:
1.
2.
“central” – bonded to more than one other atom.
“terminal” – bonded to only one other atom.
Almost always the least electronegative atom is the central
atom.
For example, in ClO2, the Cl is the central atom;
in SF5 the S is the central atom.
Hydrogen never is the central atom. It forms only one bond,
so it must generally be in the outer layer of atoms.
You can have more than one central atom in a molecule.
Bonds
Bonds are pairs of shared electrons.
Each bond has 2 electrons in it.
You can have multiple bonds between the same 2 atoms.
For example:
C-O
C=O
C O
Each of the lines represents 1 bond with 2 electrons in it.
Lewis Dot Structure
Each electron is represented by a dot in the
structure
.
:Cl:
¨
That symbol with the dots indicate a chlorine atom
with 7 valence electrons.
Drawing Lewis Dot Structures
1. Determine the total number of valence electrons.
2. Determine which atom is the “central” atom.
3. Stick everything to the central atom using a single
bond.
4. Fill the octet of every atom by adding dots.
5. Verify the total number of valence electrons in the
structure.
6. Add or subtract electrons to the structure by
making/breaking bonds to get the correct # of
valence electrons.
7. Check the “formal charge” of each atom.
Formal Charge of an Atom
Formal charge = number of valence electrons – number of
bonds – number of non-bonding electrons.
Dot structure for H2O
1. Total number of valence electrons:
6 + (2 x 1) =8
2. Central Atom – O
3. Stick all terminal atoms to the central atom using a single
bond.
Dot structure for H2O
H–O-H
Dot structure for H2O
..
H–O–H
¨
FC (H) = 1-1-0 = 0
FC (O) = 6 – 2 – 4 = 0
That is a total of 8 valence electrons used: each
bond is 2, and there are 2 non-bonding pairs.
Expanded Octets
Example PCl5:
.. ..
:Cl: :Cl:
..
..
:Cl – P - Cl :
¨ | ¨
: Cl:
¨
Total valence e- = 40
FC(P) = 5 – 5 – 0 =0
FC (Cl) = 7 – 1 – 6 = 0
Background - Covalent Bonds
• The simplest covalent bond is that in H2
– the single electrons from each atom combine to
form an electron pair
– the shared pair functions in two ways
simultaneously; it is shared by the two atoms and
fills the valence shell of each atom
• The number of shared pairs
– one shared pair forms a single bond
– two shared pairs form a double bond
– three shared pairs form a triple bond
Background – Polarity
Polar and Nonpolar Covalent Bonds
• Although all covalent bonds involve sharing of electrons, they
differ widely in the degree of sharing
• We divide covalent bonds into
– nonpolar covalent bonds
– polar covalent bonds
D i fference in
El ectron eg ati vity
Betw een Bo nded Ato ms
Less than 0.5
0.5 to 1.9
Greater than 1.9
Typ e of Bond
N on pol ar cov alent
Pol ar co valent
Io ns f orm
An example of a polar covalent bond is that of H-Cl
– the difference in electronegativity between Cl and
H is 3.0 - 2.1 = 0.9
– Polarity can be shown by using the symbols d+
and d-, or by using an arrow with the arrowhead
pointing toward the negative end and a plus sign
on the tail of the arrow at the positive end
d+ dH Cl
H
Cl
Polar bonds and polar molecules
direction
of dip ole
moment
N
O
H
H
Water
 = 1.85D
H
H
Ammonia
 = 1.47D
H
direction
of dip ole
moment
Laboratory 07
QUALITATIVE ANALYSIS : TESTING THE
SOLUBILITY RULES
Background : Ionic Compounds
1. Most ionic compounds are also called salts.
2.Most ionic compounds exist as solids and many dissolve to
form aqueous solutions.
Example : AgCl insoluble in water but AgNO3 is soluble
3. An ionic compound is made up of a metal and a nonmetal;
metals are located on the left side of the periodic table and
nonmetals are on the right side.
4. The cation (positive ion) is written first followed by the anion
(negative ion).
Nomenclature of binary ionic
compounds
Symbol
Anion
Symbol
Anion
Name
Br
Br-
Bromide
Cl
Cl-
Chloride
F
F-
Fluoride
H
H-
Hydride
I
I-
Iodide
N
N-3
Nitride
O
O-2
Oxide
P
P-3
Phosphide
S
S-2
Sulfide
NaCl
H2S
NaF
BaCl2
Mg3N2
K2O
Sodium chloride
Hydrogen sulfide
Sodium fluoride
Barium chloride
Magnesium nitride
Potassium oxide
Polyatomic anions
NO3- = nitrate
NO2- = nitrite
SO4 2 - = sulfate
SO32- = sulfite
PO43- = phosphate
CO32- = carbonate
HCO3- = hydrogen carbonate or bicarbonate
OH- = hydroxide
CN- = cyanide
C2H3O2- = acetate
C2O42- = oxalate
NaHCO3
= sodium hydrogen carbonate or sodium
bicarbonate
K2SO3
= potassium sulfite
MgSO4
= magnesium sulfate
KCN
= potassium cyanide
H2PO4
= hydrogen phosphate
Ca(OH)2
= calcium hydroxide
NH4NO3
= ammonium nitrate
Zn(NO3)2
= zinc nitrate
Li3PO4
= lithium phosphate
HNO3
= hydrogen nitrate
Solubility Rules for Ionic Compounds
in Water
Soluble Ionic Compounds
Insoluble Ionic Compounds
1. All common compounds of Group 1A 1. All common metal hydroxides are
ions (Li+, Na+, K+, etc.) and ammonium insoluble, except those of Group 1A and
ion (NH4+) are soluble.
the larger members of Group 2A
(beginning with Ca2+).
2. All common nitrates (NO3-), acetates 2. All common carbonates (CO32-) and
(CH3COO-), and most perchlorates phosphates (PO43-) are insoluble, except
(ClO4-) are soluble.
those of Group 1A and NH4+.
3. All common chlorides (Cl-), bromides 3. All common sulfides are insoluble
(Br-), and iodides (I-) are soluble, except except those of Group 1A, Group 2A,
those of Ag+, Pb2+, Cu+, and Hg2+. All and NH4+.
common fluorides (F-)are soluble,
except those of Pb2+ and Group 2A.
4. All common sulfates (SO42-) are
soluble, except those of Ca2+, Ba2+, Ag+,
and Pb2+.
Precipitation reactions
General form: Solution A + Solution B → Insoluble Solid C + Solution D.
In a precipitation reaction two solutions are mixed together to produce an
insoluble solid which is called the precipitate.
This type of reaction is also called a double displacement reaction
Lead nitrate(aq) + Potassium iodide(aq) → Lead iodide(s) + potassium nitrate(aq)
Pb(NO3)2(aq) + 2KI(aq) → PbI2(s) + KNO3(aq)
Barium chloride(aq) + Sodium sulfate(aq) → Barium sulfate(s) + Sodium chloride(aq)
BaCl2(aq) + Na2SO4(aq) → BaSO4(s) + NaCl(aq)
Copper sulfate(aq) + Sodium hydroxide(aq) → Copper hydroxide(s) + Sodium sulfate(aq)
CuSO4(aq) + 2NaOH(aq) → Cu(OH)2(s) + Na2SO4(aq)
Mercury(II) nitrate(aq) + Potassium iodide(aq) → Mercury iodide(s) + Potassium nitrate
Hg(NO3)2(aq) + 2KI(aq) → HgI2(s) + KNO3(aq)
Silver nitrate(aq) + sodium chloride(aq) →Silver chloride(s) + sodium nirate(aq)
AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)