Chapter 11 Chemical Bonds: The Formation of Compounds from ...

Chapter 11
Chemical Bonds:
The Formation of
Compounds from Atoms
1

11.1 Periodic Trends in atomic properties
11.1 Periodic Trends in atomic properties
design of periodic table is based on observing
properties of the elements
periodic trends allow us to use the periodic table to
predict properties and reactions of a wide variety of
substances
1. metals and nonmetals
• metals are usually lustrous, malleable and good
conductors of heat and electricity
• nonmetals are nonlustrous, brittle and poor
conductors
• metals tend to lose electrons and form cations
• nonmetals tend to gain electrons and form
anions
a metal reacts with a nonmetal, electrons are
often transferred from the metal to the nonmetal
• hydrogen is a unique element
2

2. atomic radius
• the increase in radius down a group
for each step down a group, an additional
energy level is added to the atom
• decrease in atomic radius across a period
electrons are added the same energy level
each time an electron is added, a proton is
added to the nucleus and increase in positive
charge that pulls the electrons closer to the
nucleus
3

3. ionization energy – the energy required to
remove an electron from the (gaseous) atom
Na + ionization energy Na + + e -
first ionization energy for the elements in the
first four periods
• ionization energy in Group A elements decreases
from top to bottom
• ionization energy gradually increases from left 4
to right across a period

11.2 Lewis structures of atoms
metals tend to form cations and nonmetals tend to
form anions in order to attain a stable valence
electron structure which contains 8 electrons
this rearrangements are accomplished by losing,
gaining or sharing electrons with other atoms
Lewis structure – using the symbol for element
and dots for valence electrons
paired dots represent paired electrons
unpaired dots represent unpaired electrons
ex. · unpaired electron
:B
paired electron symbol for element
Lewis structures for the first 20 elements
5

11.3 The ionic bond: transfer of electrons
from one atom to another
the chemistry of many elements is to attain an
outer electron structure like that of noble gases
ex. Na atom loses 3s electron to form Na +
this process requires energy
Cl atom gains an electron to form Cl -
this process releases energy
6

consider Na and Cl atoms react with each other
Na + and Cl - strongly are attracted to each other
by their opposite electrostatic charges
Lewis representation:
. . . .
Na· + ·Cl: [Na] + [:Cl:] -
·· ··
NaCl is made up of cubic crystals
each Na + is surrounded by 6 Cl - , each Cl - is
surrounded by 6 Na +
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elative sizes of Na and Cl atoms with those of
their ions
Na + is smaller than Na atom
Cl - is larger than Cl atom
ionic bond – the attraction between oppositely
charged ions
8

ex. 11.2 how Mg and Cl combine to form MgCl 2
ex. 11.3 formation of NaF from its elements
ex. 11.5 formation of MgO from its elements
9

11.4 Predicting formulas of ionic
compounds
the concept forms the basis for our understanding
of chemical bonding:
in almost all stable chemical compounds of
representative elements, each atom attains a noble
gas electron configuration
predicting the formulas of ionic compounds
ex. Compound formed between Ba and S
Ba [Xe]6s 2
S [Ne]3s 2 3p 4
loses 2 electrons to achieve
Xe configuration Ba 2+
gains 2 electrons to achieve
Ar configuration S 2-
BaS
ex. 11.8 predict formulas for
(a) magnesium sulfide Mg 2+ S 2- MgS
(b) potassium phosphide K + P 3- K10
3 P
(c) magnesium selenide Mg 2+ Se 2- MgSe

11.5 The covalent bond: sharing
electrons
electron transfer between atoms does not occur in
molecular compounds in which a chemical bond
formed by sharing pairs of electrons between
atoms
1916 G. N. Lewis introduced
covalent bond consists of a pair of electrons
shared between two atoms
examples of molecular compounds:
H 2 , Cl 2 , H 2 O, HCl, CO 2 , sugar
ionic
compound
formation of hydrogen molecule (H 2 ) involves
overlapping and pairing of 1s electron orbitals
from two H atoms
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formation of chlorine molecule (Cl 2 ) involves
overlapping of unpaired 3p electron orbitals
from two Cl atoms
other examples
. . . . . . . . . . . .
:F : F: : I : I : :O :: O: :N ::: N:
·· ·· ·· ··
fluorine iodine oxygen nitrogen
using a dash — replace the pair of dots
. . . . . . . . . . . .
:F—F: : I—I : :O = O: :N≣N:
·· ·· ·· ··
ionic bond and covalent bond represent two
extremes
between two extremes
polar covalent bond – unequal sharing of
electrons between two atoms
. .
ex. H:Cl: :C:::O:
12
··

11.6 Electronegativity
the bond formed between two different kinds of
atoms which exert unequal attraction for the pair
of electron
one atom assume a partial positive charge, the
other a partial negative charge
the attractive force that an element has for
shared electrons in a molecule or polyatomic ion
is known as its electronegativity
ex. hydrogen chloride HCl
Pauling scale assign electronegativity of F 4.0
increase
decrease
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• the highest electronegativity is 4.0 for F
the lowest electronegativity is 0.7 for Cs and Fr
• the higher the electronegativity, the stronger an
atom attracts electrons
the polarity of a bond is determined by the
difference in electronegativity values of the atoms
forming the bond
nonpolar covalent bond
polar covalent bond dipole + -
H Cl H Br I Cl O
H H
if the electronegativity difference between two
bonded atoms is greater than 1.7~1.9, the bond
will be more ionic than covalent
0 0.9 2.1
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carbon dioxide CO 2 is nonpolar
O = C = O
CCl 4 is nonpolar
H 2 O is a polar molecule
O
H H
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11.7 Lewis structures of compounds
the procedure helps to write the Lewis structure
ex. 11.9 how man valence electrons in each of
the following atoms:
Cl H C O N S P I
7 1 4 6 5 6 5 7
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ex. 11.10 write the Lewis structure for H 2 O
step 1 total valence electrons = 2 × 1 + 6 = 8
step 2 the skeleton structure
H O H : O· ·
H
H
step 3 subtract 4 electrons from 8
8 – 4 = 4
step 4 distribute 4 electrons around O atom
·· ··
H : O : H—O :
·· |
H
H
ex.11.11 write the Lewis structure for CH 4
step 1 total valence electrons = 4 × 1 + 4 = 8
step 2 H H· ·
H C H H : C : H
··
H
H
step 3 8 – 8 = 0
H
|
H—C—H
|
H
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ex.11.11 write the Lewis structure for CO 2
step 1 total valence electrons = 4 + 2 × 6 = 16
step 2
O C O O : C : O
step 3 16 – 4 = 12
step 4 . . . . . . . . . . . . . . . .
: O : C : O : : O : C : O : : O : C : O :
·· ·· ·· ··
I II III
step 5 : O :: C :: O : : O = C = O:
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11.8 Complex Lewis structures
for some molecules or polyatomic ions, no single
Lewis structure consistent with all characteristics
and bonding information can be written
ex. nitrate ion NO
-
3
step 1 total valence electrons
= 5 + 3 × 6 + 1 = 24
step 2 O O· ·
O N O O : N : O
step 3 24 - 6 = 18
step 4 · ·
: O electron deficient
·· ·· ··
: O : N : O :
·· ·· ··
step 5 · · - · · - · · -
: O : O : : O :
·· || ·· ·· ·· ·· ·· ·· ··
: O : N : O : : O =N : O : : O : N= O :
·· ·· ·· ··
there are three possible Lewis structures
a molecule or ion that has multiple correct
Lewis structures shows resonance
each of these Lewis structures is called a
resonance structure
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ex. 11.14 write the Lewis structure for
CO 3
2-
step 1 total valence electrons
= 4 + 3 × 6 + 2 = 24
step 2 O O· ·
O C O O : C : O
step 3 24 - 6 = 18
step 4 · ·
: O :
·· ·· ··
: O : C : O :
·· ··
C atom is electron deficient
step 5
·· 2- ·· 2- ·· 2-
: O : O : : O :
·· || ·· ·· ·· ·· ·· ·· ··
: O : C : O : : O =C : O : : O : C= O :
·· ·· ·· ··
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11.9 Compounds containing polyatomic
ions
sodium carbonate Na 2 CO 3 has both ionic and
covalent bonds
ionic bonds exist between Na + and CO 3
2-
covalent bonds present between C and O atoms
an important difference between ionic and
covalent bonds can be demonstrated by
dissolving Na 2 CO 3 in water:
Na 2 CO 3 (s)
water
2 Na + (aq) + CO 3
2-
(aq)
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11.10 Molecular shape
Lewis structures do not indicate anything regarding
the geometric shape of a molecule
geometric shapes for several common examples
how do we predict the geometric shape of a
molecule
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11.11 The valence shell electron pair
repulsion (VSEPR) model
VSEPR model is based on the idea that electron
pairs will repel each other electrically and will seek
to minimize this repulsion
to accomplish this minimization, the electron pairs
will arrange around the central atom as far apart as
possible
1. linear structure 180 o
ex. BeCl 2 Cl—Be—Cl
2. trigonal planar 120 o
ex. BF 3 F F
B
|F
3. tetrahedral 109.5 o
ex. CH 4 H H
C
H H
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one or more electron pairs may be nonbonding
or lone pairs
ex. NH 3 . .
H : N : H
. .
H
ex. H 2 O . .
H : O :
. .
H
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ex. 11.15 predict the molecular shape for
(1) H 2 S ..
H : S : H tetrahedral bent
..
(2) CCl 4 .. tetrahedral
: Cl :
.. .. ..
: Cl : C : Cl :
.. .. ..
: Cl :
..
(3) AlF 3 .. trigonal planar
: F :
.. .. ..
: F : Al : F :
.. ..
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