Chapter 5
Chemical Bonds, Nomenclature,
Lewis Structure and Molecular
Shapes
Homework & Quizzes – Chapter 5
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Text Homework (not turned in): pages 147 – 151.
Problems: 1, 6 – 8, 17, 27, 35, 39, 41 – 66, 68, 69,
72, 86, 88(b&c), 107, 109, 110, 112, 113.
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Quiz: Do the graded quiz in Blackboard.
I. Chemical Bonds
A. Introduction (summary of chapter)
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Atoms can combine to produce new larger units called
molecules or compounds.
Each molecule has a unique name (two rules to learn).
Molecules held together by chemical bonds (two types).
Bonds result from either transfer of valence electrons
(Ionic Bonds) or from sharing of valence electrons –
Covalent Bonds.
Valence electrons rearrange to mimic closest Group
VIIIA (18) structure.
Molecules resulting from covalent bonding will have
predictable shapes.
I. Chemical Bonds
B. Ionic Bonds
- Metals (except H) loose electrons, form cations; Nonmetals gain electrons
to form anions. Both strive for e- configuration of nearest inert gas.
- The resulting opposite ions attract in a ratio which produces a neutral unit.
Reduce formula to simplest ratio.
-
Ionic Bond Definition: bond formed by electrostatic attraction between
anions (-) and cations (+).
-
Write formula with + element first; do not show charges; Final compound is
neutral.
-
Generality: any compound formed from metallic and nonmetallic elements
is ionic.
I. Chemical Bonds

B. Ionic Bonds
Know: Metals combine with nonmetals & form ionic
bonds by losing or gaining electrons to mimic closest
Inert Gas (VIIIA).
IA
- Na, K, Li, etc become +1 ions:
IIA - Ca, Mg, etc become +2 ions:
IIIA
- Al, Ga become +3 ions:
VA
- N, P become -3 ions:
VIA
- O, S become -2 ions:
VIIA
- F, Cl, Br, I become -1 ions:
Na+
Ca+2
Al+3
N-3
O-2
F-1
Inert Gas e- Configurations
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Opposite ions attract in a ratio so that the product is neutral.
I. Chemical Bonds
B. Ionic Bonds Example
Do not show charges in final formula. NaCl NOT Na+Cl-
I. Chemical Bonds
B. Ionic Bonds
Example
I. Chemical Bonds
B. Ionic Bonds - Examples
Na+ + Na+
Ca+2
Mg+2
Al+3 + Al+3
+
+
+
+
O-2
F- + FS-2
O-2 + O-2 + O-2
Na2O
CaF2
MgS
Al2O3
Give the formulas for the following: Na & Br
Ca & O
Ba & I
Li & O Al & F Mg & N
Many transition metals form ionic bonds & can have several
charges such as Fe+2 = Iron (II); Fe+3 = Iron (III); Cu+2 =
Copper (II); Cu+1 = Copper (I)
I. Chemical Bonds
C. Electron Dot (Lewis) Structures
- A Lewis electron dot structure is a symbol in which the
valence electrons are shown as dots.
- Examples: Na.
Na+
Mg:
H:1- (Called Hydride)
:C:
Ca2+
:Si:
- How many valence electrons (dots) would
N3O2F- or Ne have? What about
8
8
8
8
Mg+2?
0
II. Covalent Bonds
A. Introduction
EN = electronegativity
- Definition of a covalent bond: A bond formed by the
sharing of two electrons.
-
-
-
When two atoms of similar EN combine, neither has
the “pull” to take electrons away & a sharing of
electrons results.
This occurs when NONMETALS, including H, combine
with NONMETALS.
Example: H. + H. ---) H—H
= H2
- The atoms share valence electrons to get stable group
VIIIA e- configurations.
II. Covalent Bonds
A. Introduction
-
Covalent bond = sharing of 2 electrons.
-
2 shared electrons with
4 shared electrons with
6 shared electrons with
-
-
-
(Single Bond).
(Double Bond).
(Triple Bond).
We frequently show the structure as a Lewis Structure covalent bonds with lines and nonbonding valence
electrons as dots.
- Note: Group IVA usually forms 4 bonds; VA three bonds;
VIA two bonds; and VIIA (along with H) one bond.
II. Covalent Bonds
H. +
F::: ---) H
H. + O
+ .H
---)
:N + N: ---) :N
B. Examples
F:::
H
O
H
N:
:::Cl. + .O. + .Cl::: ---) :::Cl
O
Cl:::
::O: + :C: + :O:: -----) ::O = C = O::
II. Covalent Bonds
C. Lewis Structures 1. Rules for
Drawing Lewis Structures
1. Calculate the total # of valence electrons; take into account charge if the
sample is an ion.
2. Place atom that forms most bonds at center (Closest to Group IVA & Lowest
if in same group). If there is a charge, then add or subtract the appropriate
number of electrons on the central atom.
3. Arrange other atoms around central atom & allow sharing so that each atom
has stable electron configuration. Show bonding pairs as dashes &
nonbonding valence e- as dots.
4. Double check: a) each atom has a stable electron configuration & b) have the
same total number of valence electrons as in step 1.
II. Covalent Bonds
C. Lewis Structures
2. Examples
HI
H2O
NH4+
H2O2
CH4
SO2
AlCl4-
NO2-
CN-
Bonding Summary
Two General Bonding Types
1.
Ionic: Compound containing metallic element.
Atoms lose/gain e to look like nearest inert
gas. Add together ions such that neutralize
charge.
Ia
+1
IIa
+2
IIIa
+3
Va
-3
VIa
-2
VIIa
-1
2. Covalent: Compound containing nonmetals.
Atoms obtain inert gas configuration by
sharing valence electrons. :
::
:::
II. Covalent Bonds – Organic Compounds
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Can write organic structures several ways.
Example – Butane (Note the five ways of presenting)
Note: Carbon always has four bonds.
C4H10
CH3CH2CH2CH3
H
H H
H
H–C–C–C–C–H
H
H H
H
CH3-CH2-CH2-CH3
II. Covalent Bonds – Organic Compounds

Cyclic Organics: Example of Cyclopropane
H
H
C
H
C C
H
H
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Aromatics
H
Contain Benzene, C6H6
H
H
H
C
C
C
C
H
C
C
H
C6H6
H
II. Covalent Bonds – Organic Compounds
O
O
C OH
C OH
OH
O C CH3
O
Salicylic Acid
C7H6O3
MW = 138g
Acetylsalicylic Acid
C9H8O4
MW = 180 g
II. Covalent Bonds – Organic Compounds – Aspirin Lab
1) Equation & Conversion Factors:
1 Salicylic Acid + 1 Acetic Anhydride -----) 1 Aspirin + 1 Acetic Acid
1 = molecules or moles; 1 mole = formula weight in grams = 6.0x1023 molecules
2) Lab Calculations (questions 2 & 3):
2.0 g SA x 1 mole SA = 0.014 mole SA
138 g SA
0.014 mole SA x 1 mole Aspirin = 0.014 mole Aspirin
1 mole SA
From the coefficients in the balanced chemical equation above.
III. Shapes
o
Molecular Shapes play a major role in:
1) Physical Properties
2) Chemical Properties
3) Biochemical Properties
o
To Obtain the shape of a molecule one draws the
Lewis Structure, counts the number of “things”
around the central atom, and uses simple geometry to
predict the shape.
III. Shapes
C. Simplified Examples
Bond angle = 180o
Bond angle = 120o
Bond angle = 109o
IV. Nomenclature A. Introductions

There are common & systematic names for chemicals. A
chemical may have scores of common names.

A systematic name must allow one to both obtain the
formula and derive the name from the formula.

There are two general rules for naming inorganic
compounds.

Ionic compounds use Rule #1.
compounds use Rule #2.
Molecular or Covalent
IV. Nomenclature B. Ionic Compounds
-
Rule #1 for ionic compounds: Name
the + element, then the – element and
change the ending to “ide.”
-
Examples:
NaCl = Sodium Chloride
Na2O = Sodium Oxide
IV. Nomenclature B. Ionic Compounds
Rule #1 – “ide” names
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Negative atoms have an “ide” ending.
Atom
Chlorine
Oxygen
Fluorine
Sulfur
Nitrogen
Iodine
Bromine
Phosphorus
Anion
Cl1O2F1S2N3I1Br1P3-
Name
Chloride
Oxide
Fluoride
Sulfide
Nitride
Iodide
Bromide
Phosphide
IV. Nomenclature B. Ionic Compounds
Examples
NaCl
Na2O
AlF3
Be3N2
Sodium Chloride
Sodium Oxide
Aluminum Fluoride
Beryllium Nitride
Calcium Sulfide
Barium Iodide
Barium Oxide
Magnesium Nitride
CaS
BaI2
BaO
Mg3N2
IV. Nomenclature C. Molecular Compounds
Rule #2
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When nonmetals & H combine with each other
through sharing electrons (covalent bonds),
they form molecules; there are no ions.
Rule #2 – When both elements are nonmetals
(molecular compounds), then Name the + &
the - & change ending to “ide” as before. Use
prefixes of di, tri, tetra, penta, etc to tell how
many of each element is present.
IV. Nomenclature C. Molecular Compounds
CO2
CCl4
N2O
P2S5
PBr3
BI3
=
=
=
=
=
=
Carbon Dioxide
Carbon Tetrachloride
Dinitrogen Oxide
Diphosphorus Pentasulfide
Phosporus Tribromide
Boron Triiodide
Notes: (1) Organic compounds like CH4 use their own
rules which we won’t cover.
(2) diatomic molecules named with the element
name. O2 = Oxygen
V. Polyatomic Ions
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Previous compounds formed from two elements.
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Frequently have compounds formed from three or
four elements. When this happens, then usually
have a polyatomic ion present.
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Polyatomic ions: stable ions formed from two or
more elements; held together by covalent bonds.
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Examples:
SO4-2 = Sulfate
NO2- = nitrite
PO4-3 = Phosphate
V. Polyatomic Ions
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Polyatomic ions are held together by covalent bonds,
and they form ionic bonds with metals.
Examples:
NaNO2
Na2SO4
Na3PO4
When have more than one polyatomic ion in a
compound then use parentheses around the ion.
Examples: Na2SO3 Ca(NO2)2
Ca3(PO4)2
Nomenclature: Simply use the polyatomic ion name.
Example: Calcium Nitrite & Calcium Phosphate above
Need to memorize the following polyatomic ions, their
names and their charges.
V. Polyatomic Ions - Memorize the Names,
Formulas and the Charges
Formula
NH4+
Name
Formula
Name
Ammonium (The Only Positive One in this list)
C2H3O2NO3OHHCO3-
Acetate
Nitrate
Hydroxide
Hydrogen Carbonate
CNNO2-
Cyanide
Nitrite
CO3-2
SO4-2
Carbonate
Sulfate
Cr2O7-2
SO3-2
Dichromate
Sulfite
PO4-3
Phosphate
V. Polyatomic Ions – Examples of Naming
and Obtaining Formulas
Aluminum Hydroxide
Calcium Cyanide
Barium Sulfate
Ammonium Nitrate
Al(OH)3
Ca(CN)2
BaSO4
NH4NO3
Ba(OH)2
LiNO2
KNO3
NaHCO3
Al2(SO4)3
Barium Hydroxide
Lithium Nitrite
Potassium Nitrate
Sodium Hydrogen Carbonate
Aluminum Sulfate
Naming: Mixed Examples
NaF
Sodium Fluoride
CS2
Carbon Disulfide
NI3
Nitrogen Triiodide
BaI2
Barium Iodide
K3PO4
Boron Trifluoride
Sodium Sulfite
Potassium Phosphate
BF3
Na2SO3