Chapter 11 Chemical Bonds: The Formation of Compounds from ...
Chapter 11 Chemical Bonds: The Formation of Compounds from ...
Chapter 11 Chemical Bonds: The Formation of Compounds from ...
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<strong>Chapter</strong> <strong>11</strong><br />
<strong>Chemical</strong> <strong>Bonds</strong>:<br />
<strong>The</strong> <strong>Formation</strong> <strong>of</strong><br />
<strong>Compounds</strong> <strong>from</strong> Atoms<br />
1
<strong>11</strong>.1 Periodic Trends in atomic properties<br />
<strong>11</strong>.1 Periodic Trends in atomic properties<br />
design <strong>of</strong> periodic table is based on observing<br />
properties <strong>of</strong> the elements<br />
periodic trends allow us to use the periodic table to<br />
predict properties and reactions <strong>of</strong> a wide variety <strong>of</strong><br />
substances<br />
1. metals and nonmetals<br />
• metals are usually lustrous, malleable and good<br />
conductors <strong>of</strong> heat and electricity<br />
• nonmetals are nonlustrous, brittle and poor<br />
conductors<br />
• metals tend to lose electrons and form cations<br />
• nonmetals tend to gain electrons and form<br />
anions<br />
a metal reacts with a nonmetal, electrons are<br />
<strong>of</strong>ten transferred <strong>from</strong> the metal to the nonmetal<br />
• hydrogen is a unique element<br />
2
2. atomic radius<br />
• the increase in radius down a group<br />
for each step down a group, an additional<br />
energy level is added to the atom<br />
• decrease in atomic radius across a period<br />
electrons are added the same energy level<br />
each time an electron is added, a proton is<br />
added to the nucleus and increase in positive<br />
charge that pulls the electrons closer to the<br />
nucleus<br />
3
3. ionization energy – the energy required to<br />
remove an electron <strong>from</strong> the (gaseous) atom<br />
Na + ionization energy Na + + e -<br />
first ionization energy for the elements in the<br />
first four periods<br />
• ionization energy in Group A elements decreases<br />
<strong>from</strong> top to bottom<br />
• ionization energy gradually increases <strong>from</strong> left 4<br />
to right across a period
<strong>11</strong>.2 Lewis structures <strong>of</strong> atoms<br />
metals tend to form cations and nonmetals tend to<br />
form anions in order to attain a stable valence<br />
electron structure which contains 8 electrons<br />
this rearrangements are accomplished by losing,<br />
gaining or sharing electrons with other atoms<br />
Lewis structure – using the symbol for element<br />
and dots for valence electrons<br />
paired dots represent paired electrons<br />
unpaired dots represent unpaired electrons<br />
ex. · unpaired electron<br />
:B<br />
paired electron symbol for element<br />
Lewis structures for the first 20 elements<br />
5
<strong>11</strong>.3 <strong>The</strong> ionic bond: transfer <strong>of</strong> electrons<br />
<strong>from</strong> one atom to another<br />
the chemistry <strong>of</strong> many elements is to attain an<br />
outer electron structure like that <strong>of</strong> noble gases<br />
ex. Na atom loses 3s electron to form Na +<br />
this process requires energy<br />
Cl atom gains an electron to form Cl -<br />
this process releases energy<br />
6
consider Na and Cl atoms react with each other<br />
Na + and Cl - strongly are attracted to each other<br />
by their opposite electrostatic charges<br />
Lewis representation:<br />
. . . .<br />
Na· + ·Cl: [Na] + [:Cl:] -<br />
·· ··<br />
NaCl is made up <strong>of</strong> cubic crystals<br />
each Na + is surrounded by 6 Cl - , each Cl - is<br />
surrounded by 6 Na +<br />
7
elative sizes <strong>of</strong> Na and Cl atoms with those <strong>of</strong><br />
their ions<br />
Na + is smaller than Na atom<br />
Cl - is larger than Cl atom<br />
ionic bond – the attraction between oppositely<br />
charged ions<br />
8
ex. <strong>11</strong>.2 how Mg and Cl combine to form MgCl 2<br />
ex. <strong>11</strong>.3 formation <strong>of</strong> NaF <strong>from</strong> its elements<br />
ex. <strong>11</strong>.5 formation <strong>of</strong> MgO <strong>from</strong> its elements<br />
9
<strong>11</strong>.4 Predicting formulas <strong>of</strong> ionic<br />
compounds<br />
the concept forms the basis for our understanding<br />
<strong>of</strong> chemical bonding:<br />
in almost all stable chemical compounds <strong>of</strong><br />
representative elements, each atom attains a noble<br />
gas electron configuration<br />
predicting the formulas <strong>of</strong> ionic compounds<br />
ex. Compound formed between Ba and S<br />
Ba [Xe]6s 2<br />
S [Ne]3s 2 3p 4<br />
loses 2 electrons to achieve<br />
Xe configuration Ba 2+<br />
gains 2 electrons to achieve<br />
Ar configuration S 2-<br />
BaS<br />
ex. <strong>11</strong>.8 predict formulas for<br />
(a) magnesium sulfide Mg 2+ S 2- MgS<br />
(b) potassium phosphide K + P 3- K10<br />
3 P<br />
(c) magnesium selenide Mg 2+ Se 2- MgSe
<strong>11</strong>.5 <strong>The</strong> covalent bond: sharing<br />
electrons<br />
electron transfer between atoms does not occur in<br />
molecular compounds in which a chemical bond<br />
formed by sharing pairs <strong>of</strong> electrons between<br />
atoms<br />
1916 G. N. Lewis introduced<br />
covalent bond consists <strong>of</strong> a pair <strong>of</strong> electrons<br />
shared between two atoms<br />
examples <strong>of</strong> molecular compounds:<br />
H 2 , Cl 2 , H 2 O, HCl, CO 2 , sugar<br />
ionic<br />
compound<br />
formation <strong>of</strong> hydrogen molecule (H 2 ) involves<br />
overlapping and pairing <strong>of</strong> 1s electron orbitals<br />
<strong>from</strong> two H atoms<br />
<strong>11</strong>
formation <strong>of</strong> chlorine molecule (Cl 2 ) involves<br />
overlapping <strong>of</strong> unpaired 3p electron orbitals<br />
<strong>from</strong> two Cl atoms<br />
other examples<br />
. . . . . . . . . . . .<br />
:F : F: : I : I : :O :: O: :N ::: N:<br />
·· ·· ·· ··<br />
fluorine iodine oxygen nitrogen<br />
using a dash — replace the pair <strong>of</strong> dots<br />
. . . . . . . . . . . .<br />
:F—F: : I—I : :O = O: :N≣N:<br />
·· ·· ·· ··<br />
ionic bond and covalent bond represent two<br />
extremes<br />
between two extremes<br />
polar covalent bond – unequal sharing <strong>of</strong><br />
electrons between two atoms<br />
. .<br />
ex. H:Cl: :C:::O:<br />
12<br />
··
<strong>11</strong>.6 Electronegativity<br />
the bond formed between two different kinds <strong>of</strong><br />
atoms which exert unequal attraction for the pair<br />
<strong>of</strong> electron<br />
one atom assume a partial positive charge, the<br />
other a partial negative charge<br />
the attractive force that an element has for<br />
shared electrons in a molecule or polyatomic ion<br />
is known as its electronegativity<br />
ex. hydrogen chloride HCl<br />
Pauling scale assign electronegativity <strong>of</strong> F 4.0<br />
increase<br />
decrease<br />
13
• the highest electronegativity is 4.0 for F<br />
the lowest electronegativity is 0.7 for Cs and Fr<br />
• the higher the electronegativity, the stronger an<br />
atom attracts electrons<br />
the polarity <strong>of</strong> a bond is determined by the<br />
difference in electronegativity values <strong>of</strong> the atoms<br />
forming the bond<br />
nonpolar covalent bond<br />
polar covalent bond dipole + -<br />
H Cl H Br I Cl O<br />
H H<br />
if the electronegativity difference between two<br />
bonded atoms is greater than 1.7~1.9, the bond<br />
will be more ionic than covalent<br />
0 0.9 2.1<br />
14
carbon dioxide CO 2 is nonpolar<br />
O = C = O<br />
CCl 4 is nonpolar<br />
H 2 O is a polar molecule<br />
O<br />
H H<br />
15
<strong>11</strong>.7 Lewis structures <strong>of</strong> compounds<br />
the procedure helps to write the Lewis structure<br />
ex. <strong>11</strong>.9 how man valence electrons in each <strong>of</strong><br />
the following atoms:<br />
Cl H C O N S P I<br />
7 1 4 6 5 6 5 7<br />
16
ex. <strong>11</strong>.10 write the Lewis structure for H 2 O<br />
step 1 total valence electrons = 2 × 1 + 6 = 8<br />
step 2 the skeleton structure<br />
H O H : O· ·<br />
H<br />
H<br />
step 3 subtract 4 electrons <strong>from</strong> 8<br />
8 – 4 = 4<br />
step 4 distribute 4 electrons around O atom<br />
·· ··<br />
H : O : H—O :<br />
·· |<br />
H<br />
H<br />
ex.<strong>11</strong>.<strong>11</strong> write the Lewis structure for CH 4<br />
step 1 total valence electrons = 4 × 1 + 4 = 8<br />
step 2 H H· ·<br />
H C H H : C : H<br />
··<br />
H<br />
H<br />
step 3 8 – 8 = 0<br />
H<br />
|<br />
H—C—H<br />
|<br />
H<br />
17
ex.<strong>11</strong>.<strong>11</strong> write the Lewis structure for CO 2<br />
step 1 total valence electrons = 4 + 2 × 6 = 16<br />
step 2<br />
O C O O : C : O<br />
step 3 16 – 4 = 12<br />
step 4 . . . . . . . . . . . . . . . .<br />
: O : C : O : : O : C : O : : O : C : O :<br />
·· ·· ·· ··<br />
I II III<br />
step 5 : O :: C :: O : : O = C = O:<br />
18
<strong>11</strong>.8 Complex Lewis structures<br />
for some molecules or polyatomic ions, no single<br />
Lewis structure consistent with all characteristics<br />
and bonding information can be written<br />
ex. nitrate ion NO<br />
-<br />
3<br />
step 1 total valence electrons<br />
= 5 + 3 × 6 + 1 = 24<br />
step 2 O O· ·<br />
O N O O : N : O<br />
step 3 24 - 6 = 18<br />
step 4 · ·<br />
: O electron deficient<br />
·· ·· ··<br />
: O : N : O :<br />
·· ·· ··<br />
step 5 · · - · · - · · -<br />
: O : O : : O :<br />
·· || ·· ·· ·· ·· ·· ·· ··<br />
: O : N : O : : O =N : O : : O : N= O :<br />
·· ·· ·· ··<br />
there are three possible Lewis structures<br />
a molecule or ion that has multiple correct<br />
Lewis structures shows resonance<br />
each <strong>of</strong> these Lewis structures is called a<br />
resonance structure<br />
19
ex. <strong>11</strong>.14 write the Lewis structure for<br />
CO 3<br />
2-<br />
step 1 total valence electrons<br />
= 4 + 3 × 6 + 2 = 24<br />
step 2 O O· ·<br />
O C O O : C : O<br />
step 3 24 - 6 = 18<br />
step 4 · ·<br />
: O :<br />
·· ·· ··<br />
: O : C : O :<br />
·· ··<br />
C atom is electron deficient<br />
step 5<br />
·· 2- ·· 2- ·· 2-<br />
: O : O : : O :<br />
·· || ·· ·· ·· ·· ·· ·· ··<br />
: O : C : O : : O =C : O : : O : C= O :<br />
·· ·· ·· ··<br />
20
<strong>11</strong>.9 <strong>Compounds</strong> containing polyatomic<br />
ions<br />
sodium carbonate Na 2 CO 3 has both ionic and<br />
covalent bonds<br />
ionic bonds exist between Na + and CO 3<br />
2-<br />
covalent bonds present between C and O atoms<br />
an important difference between ionic and<br />
covalent bonds can be demonstrated by<br />
dissolving Na 2 CO 3 in water:<br />
Na 2 CO 3 (s)<br />
water<br />
2 Na + (aq) + CO 3<br />
2-<br />
(aq)<br />
21
<strong>11</strong>.10 Molecular shape<br />
Lewis structures do not indicate anything regarding<br />
the geometric shape <strong>of</strong> a molecule<br />
geometric shapes for several common examples<br />
how do we predict the geometric shape <strong>of</strong> a<br />
molecule <br />
22
<strong>11</strong>.<strong>11</strong> <strong>The</strong> valence shell electron pair<br />
repulsion (VSEPR) model<br />
VSEPR model is based on the idea that electron<br />
pairs will repel each other electrically and will seek<br />
to minimize this repulsion<br />
to accomplish this minimization, the electron pairs<br />
will arrange around the central atom as far apart as<br />
possible<br />
1. linear structure 180 o<br />
ex. BeCl 2 Cl—Be—Cl<br />
2. trigonal planar 120 o<br />
ex. BF 3 F F<br />
B<br />
|F<br />
3. tetrahedral 109.5 o<br />
ex. CH 4 H H<br />
C<br />
H H<br />
23
one or more electron pairs may be nonbonding<br />
or lone pairs<br />
ex. NH 3 . .<br />
H : N : H<br />
. .<br />
H<br />
ex. H 2 O . .<br />
H : O :<br />
. .<br />
H<br />
24
ex. <strong>11</strong>.15 predict the molecular shape for<br />
(1) H 2 S ..<br />
H : S : H tetrahedral bent<br />
..<br />
(2) CCl 4 .. tetrahedral<br />
: Cl :<br />
.. .. ..<br />
: Cl : C : Cl :<br />
.. .. ..<br />
: Cl :<br />
..<br />
(3) AlF 3 .. trigonal planar<br />
: F :<br />
.. .. ..<br />
: F : Al : F :<br />
.. ..<br />
25