Chapter 3

Chemical reactions

1. Chapter 3: Chapter 3: Chemical BondsChemical Bonds
• Octet Rule
• Naming anions and cations
• Ionic and Covalent bonds
• Electronegativity
• Drawing Lewis dot structures
• Octet rule exceptions
• Resonance
• Bond Angles
• Polarity of molecules

2. The Octet Rule
Main group elements react in ways that achieve an
electron configuration of eight valence electrons.
–An atom that loses one or more electrons
becomes a positively charged ion = cation.cation.
–An atom that gains one or more electrons
becomes a negatively charged ion = anion.anion.

3. By losing one electron, a sodium atom forms a sodium ion,
which has the same electron configuration as neon.

Na (11 electrons): 1s
2
2s
2
2p
6
3s
1
Na
+
(10 electrons): 1s
2
2s
2
2p
6
Neon (10 electrons): 1s
2
2s
2
2p
6

4. By gaining one electron, a chlorine atom forms a
chloride ion, which has the same electron
configuration as argon.
Chlorine atom (17 electrons): 1s
2
2s
2
2p
6
3s
2
3p
5
Chloride ion (18 electrons): 1s
2
2s
2
2p
6
3s
2
3p
6

5. When an atom gains electrons, it has a ________
charge and is called a(n) ________.
1.negative; anion
2.negative; cation
3.positive; anion
4.positive; cation

6. The Octet Rule
The octet rule gives us a good way to understand why Main Group elements
form the ions they do:
elements in group1 always loss 1 electron
group 2 always loss 2 electrons
group 7A (17) always gain 1 electron
group 6A (16) always gain 2 electrons
but it is not perfect:
–Ions of period 1 and 2 elements with charges greater than +2 (i.e.+3, +4)
or smaller than -2 (i.e. -3, -4) are unstable. For example, boron does not
lose its three valence electrons to become B
3+
, nor does carbon lose its
four valence electrons to become C
4+
or gain four valence electrons to
become C
4-
–The octet rule does not apply to transition elements, most of which form
ions with two or more different positive charges.

7. Forming Chemical Bonds
An atom may lose or gain enough electrons to
acquire a filled valence shell and become an
ion. An ionic bondionic bond is the result of the force of
attraction between a cation and an anion.
An atom may share electrons with one or more
other atoms to acquire a filled valence shell. A
covalent bond covalent bond results from two atoms that
share one or more pairs of electrons.

8. Ionic bondsIonic bonds form by transfer of one or more valence e
-
from
an atom that tends to give away electrons to another that
tends to accept electrons.
Cation Anion

9. Formulas of Ionic Compounds
In a formula of an ionic compound:
number of positive charges = number of negative charges
(+) LiBr(-) lithium bromide
(+2) BaI
2
(2)(-1)barium iodide
(+3)(2) Al
2
S
3
(3)(-2)aluminum sulfide
(+3) K
3
PO
4
(-3) potassium phosphate
Cation Anion

10. How to name Ionic Compounds
Cation
Anion
Ionic Compound
Name cation first then anion
Groups I and II
(Non transition metals)
i.e Na
+
sodium
Transition metals i.e. Fe
2+
Iron II
Fe
3+
Iron III
Ag
+
silver
Monoatomic
Stem part of the name plus sufix –ide
i.e. S
-2
sulfide Cl
-1
chloride
Oxoanions (contain oxygen)
i.e. non metals groups 15 to 17 plus oxygen
SO
4
-2
sulfate ClO
4
-
perchlorate
SO
3
-2
sulfite ClO
3
-
chlorate
ClO
2
-
chlorite
ClO
-
hypochlorite
Other polyatomic anions
(contain two or more different atoms)
i.e. CN
-
cyanide
Polyatomic cations
(two or more different atoms)
i.e. NH
4
+
ammonium
By M. Castillo
(metal and nonmetal)

11. CATIONS

12. For cations derived from transition and inner transition
elements, most of which form more than one type of cation,
– use either Roman numerals to show charge, or
– use the suffix-ous-ous to show the lower + charge,
-ic-ic to show the higher + charge.

13. Naming Anions
For monatomic (containing only one atom) anions, add
“ide” to the stem part of the name.

14. ANIONS

15. Examples
AlCl
3
= aluminum chloride
Ag
2
S = silver sulfide
FeO = iron(II) oxide; ferrous oxide
Fe
2
O
3
= iron(III) oxide; ferric oxide
NaH
2
PO
4
= sodium dihydrogen phosphate.
NH
4
OH= ammonium hydroxide.
FeCO
3
= iron(II) carbonate or ferrous carbonate
Fe
2
(CO
3
)
3
= iron(III) carbonate or ferric carbonate.

16. What is the systematic name for MnO?
1.manganese oxide
2.manganese trioxide
3.manganese (III) oxide
4.manganese (II) oxide

17. What is the formula for an ionic compound formed
between a calcium ion and a selenide ion?
1. 2.
3. 4.

18. Forming a Covalent Bond
A covalent bond is formed by sharing one or more pairs of
electrons. The pair of electrons is shared by both atoms
and, at the same time, fills the valence shell of each atom.

19. How to name Binary Covalent Compounds
Covalent Compound
(two non metals)
The name is made out of two words
First word
Name element that appears first
(usually the less electronegative)
Indicate the number of atoms by a Greek prefix
(di, tri, tetra, penta, hexa)
Second word
Name second element (the more electronegative)
Indicate the number of atoms by a Greek prefix
followed by the stem part of the name plus suffix
–ide
Example:
N
2
O
5
dinitrogen pentaoxide

20. NO is nitrogen oxide (nitric oxide)
SF
2
is sulfur difluoride
N
2
O is dinitrogen oxide (laughing gas)

21. 1.dinitrogen tetroxide
2.dinitrogen tetroxygen
3.nitrogen dioxide
4.nitrogen (IV) oxide
What is the systematic name
for N
2O
4?

22. Electronegativity
Electronegativity:Electronegativity: a measure of an atom’s attraction for
electrons.

23. Although all covalent bonds involve sharing of electron pairs, they differ in the
degree of sharing:
nonpolar covalent bond:nonpolar covalent bond: electrons are shared equally
polar covalent bond:polar covalent bond: electron sharing is not equal

24. H-Cl
Bond
Difference in
ElectronegativityType of Bond
3.5 - 2.1 = 1.4
3.0 - 2.1 = 0.9
4.0 - 0.9 = 3.1
2.5 - 1.2 = 1.3
polar covalent
polar covalent
ionic
polar covalent
2.5 - 2.5 = 0.0 nonpolar covalent
3.0 - 2.1 = 0.9 polar covalent
O-H
N-H
Na-F
C-Mg
C-S

25. In a polar covalent bond
–the more electronegative atom gains a greater fraction of the shared
electrons and acquires a partial negative charge ( dd-)-)
–the less electronegative atom acquires a partial positive charge (dd++)

26. An ionic bond has a ________ electronegativity difference between
atoms than a covalent bond and almost always involves ________.
1.greater; a metal and a nonmetal
2.greater; two nonmetals
3.lesser; a metal and a nonmetal
4.lesser; two nonmetals

27. Drawing Lewis Structures
Count the number of valence electrons
•For a molecule add up the valence electrons of the atoms present.
Draw a skeleton structure joining atoms by single bonds
•The central atom is usually written first in the formula
Determine the number of valence e
-
still available
•Subtract two e- for each single bond written
Determine the number of valence electrons required to fill an octet
for each atom (except H)
•If e
-
available = e
-
required distribute available e
-
as unshared pairs
•If the number of e- available is less than the number required by two e
-

change a single bond into a double bond.
•If you are four e
-
short convert a single bond into a triple bond.
•C, N, O and S can form multiple bonds. Hydrogen and halogens never form
double bonds.

30. –draw a Lewis structure for hydrogen peroxide, H
2
O
2
.
–draw a Lewis structure for methanol, CH
3
OH.
–draw a Lewis structure for acetic acid, CH
3
COOH.

31. Expanded octets:
exceptions to the Octet Rule
In these molecules the central atom is surrounded by more than 4
pairs of valence electrons.
In molecules of this type the terminal atoms are most often halogens
(F, Cl, Br, I) or oxygen. The central atom is a non metal in the third,
fourth, or fifth period of the periodic table. Most frequently one of the
following elements:
Group 15Group 16Group 17Group 18
3
rd
period P S Cl -
4
rd
period As Se Br Kr
5
rd
period Sb Te I Xe

32. Expanded Octets

33. There are a few species in which the central atom is surrounded by 2 or 3 e
-

pairs rather than 4. Although different structures can be written following
the octet rule, experimental evidence suggests the structures:
Exceptions to the Octet rule: Electron deficient molecules
For odd e
-
molecules (called free radicals) is not possible to come up with a
Lewis structure in which all atoms obey the octet rule. Example
NO number of valence e- = 5 + 6 = 11
NO
2
number of valence e- = 5 + 6(2) = 17

34. Resonance
Many molecules and ions are best described by writing two or more
Lewis structures. The true molecule is a hybridhybrid of the contributing
structures.

35. Which of these Lewis Structures has resonance?
1. 2.
3.
4.
1 2 3 4 5 6 7 8 9 1011121314151617181920
2122232425262728293031323334353637383940
4142434445464748495051525354555657585960
6162636465666768697071727374757677787980
81828384858687888990919293949596979899100
101102103104105106107108109110

36. Valence-Shell Electron-Pair Repulsion
(VSEPR) Model
–valence electrons of an atom may be involved in forming bonds
or may be unshared.
–each combination creates a negatively charged region of
electrons around the nucleus.
–because like charges repel each other, the various regions of
electron density around an atom spread so that each is as far
away from the others as possible.

38. Predict the shape of methane, CH
4
–The Lewis structure shows carbon surrounded by four regions of electron
density.
–According to the VSEPR model, the four regions radiate from carbon at angles of
109.5°, and the shape of the molecule is tetrahedral.
–The measured H-C-H bond angles are 109.5°.

39. Predict the shape of ammonia, NH
3
–nitrogen is surrounded by four regions of electron density (3 with single
pairs of electrons, and 1 with an unshared pair of electrons).
–According to the VSEPR model, the four regions radiate from nitrogen
at angles of 109.5°, and the shape of the molecule is pyramidal.
–The measured H-N-H bond angles are 107.3°

40. Predict the shape of water, H
2
O
–The Lewis structure shows oxygen with four regions of electron density
(2 regions with single pairs of e
-
, and 2 with unshared pairs of e
-
.
–According to the VSEPR model, the four regions radiate from oxygen at
angles of 109.5°, and the shape of the molecule is bent.
–The measured H-O-H bond angle is 104.5°.

41. Predict the shape of formaldehyde, CH
2
O
–The Lewis structure shows carbon surrounded by 3 regions of electron
density; 2 with single pairs of e
-
and one with 2 pairs of e
-
forming the
double bond to oxygen.
–According to the VSEPR model, the three regions radiate from carbon
at angles of 120°, and the shape of the molecule is planar (trigonal
planar).
–The measured H-C-H bond angle is 116.5°.

42. Predict the shape of ethylene, C
2
H
4
–The Lewis structure shows carbon surrounded by 3 regions of e
-
density;
2 with single pairs of e
-
and 1 with two pairs of electrons forming the
double bond to the other carbon.
–According to the VSEPR model, the three regions radiate from carbon
at angles of 120°, and the shape of the molecule is planar (trigonal
planar).
–The measured H-C-H bond angle is 117.2°.

43. Predict the shape of acetylene, C
2
H
2
–The Lewis structure shows carbon surrounded by 2 regions of electron
density; one region with a single pair of e
-
, and the other one with three
pairs of e
-
forming the triple bond to carbon.
–According to the VSEPR model, the two regions radiate from carbon at
an angle of 180°, and the shape of the molecule is linear.
–The measured H-C-C bond angle is 180°.

44. Polarity of Molecules
A molecule will be polar if:
–it has polar bonds, and
–its centers of partial positive and partial
negative charges lie at different places
within the molecule.

45. Carbon dioxide, CO
2
, has two polar bonds but, because
of its geometry, is a nonpolar molecule

46. Water, H
2
O, has two polar bonds and, because of its
geometry, is a polar molecule.

47. Ammonia, NH
3
, has three polar bonds and, because of its
geometry, is a polar molecule.

48. Both dichloromethane, CH
2
Cl
2
, and formaldehyde, CH
2
O,
have polar bonds and are polar molecules.

49. 1.linear.
2.bent.
3.trigonal planar.
4.trigonal pyramidal.
The VSEPR shape of SO
3
is predicted to be:
1 2 3 4 5 6 7 8 9 1011121314151617181920
2122232425262728293031323334353637383940
4142434445464748495051525354555657585960
6162636465666768697071727374757677787980
81828384858687888990919293949596979899100
101102103104105106107108109110

50. 1.polar; polar
2.polar; nonpolar
3.nonpolar; polar
4.nonpolar; nonpolar
SO
3
is a _______ molecule that contains _______ bonds.

51. Chapter 4 Review Questions
1.Write the ion magnesium forms. What is the name?
2.Write the ion does chlorine forms. What is the name?
3.Write the symbols for the following ions.
a) Silver(I) ion
b) Iron (III) ion
c) Cuprous ion
4.Name the following polyatomic ions.
a) NO
3
-
b) CO
3
-2
c) OH
-
d) PO
4
-3

52. 5.Name the following compounds, using Roman numerals
to indicate the charges on the cations where necessary.
a) KF
b) MgCl
2
c) (NH
4
)
2
CO
3
d) MgSO
4
e) Fe
2
O
3
6.Write the formula for the following compounds.
a) sodium hydroxide
b) Magnesium chloride
c) copper(II) carbonate
d) calcium phosphate
•Draw a Lewis structure for vinyl chloride, C
2
H
3
Cl, a substance
used in making PVC plastic.

53. 9. What shape do you expect for the hydronium ion, H
3
O
+
?
10.Predict the geometry of an acetaldehyde molecule, CH
3
CHO?

54. 12.Use electronegativity differences to classify bonds between
the following pairs of atoms as ionic, nonpolar covalent, or
polar covalent.
a) I (2.7) and Cl (3.2)
b) Li (1.0) and O (3.4)
c) Br (3.0) and Br (3.0)
d) P (2.2) and Br (3.0)
13.Use the symbols d+ and d- to identify the location of the partial
charges on the polar covalent bonds formed between the
following:
a) Fluorine and sulfur
b) phosphorus and oxygen
14. Name the following compounds.
a) N
2
O
3
b) PCl
5
c) SeF
4