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Chapter 12
Chemical Bonding
CATIONS
These are species where a neutral atom loses electrons
to form a positive charge ion.
EXAMPLE:
Na
1s22s22p63s1
→
Na+
1s22s22p6
+ e-
This is the same electronic
configuration of Ne.
“ISOELECTRONIC”
ANIONS
These are species where a neutral atom gains electrons
to form a negative charge ion.
EXAMPLE:
Cl
+ e1s22s22p63s23p5
→
Cl1s22s22p63s23p6
This is the same electronic
configuration of Ar.
“ISOELECTRONIC”
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Other ANIONS
O
1s22s22p4
+ 2e-
→
O21s22s22p6
THIS IS THE SAME ELECTRONIC
CONFIGURATION as Neon.
As
+ 3e-
1s22s22p63s23p54s23d104p3
→
As3-
1s22s22p63s23p64s23d104p6
THIS IS THE SAME ELECTRONIC
CONFIGURATION as Krypton.
Monatomic Ions
Monatomic Ions
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Monatomic Ions
Figure 12.1-Moatomic ions with noble gas configurations. Each color group includes one
noble-gas atom and the monatomic ions that are isoelectronic with that atom.
Practice Problem
Identify which ions are isoelectronic to argon?
Rb+
Sr2+
S2Na+
Sc3+
Te2N3P3Ca2+
Cl-
Ionic Compounds
Ionic compounds are substances which are
comprised of an attraction of ions.
For example, table salt or sodium chloride is
comprised of sodium ions and chloride ions.
ILLUSTRATION:
Na
+
Cl
Na
[ Cl ]
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Ionic Bond
Sodium Chloride
In this compound, there
is an arrangement of
ions.
Gray is Na+ ions and
Green is Cl1- ions. In
this case, there is a 1:1
ratio between Na+ and
Cl1- ions in its cubical
structure.
Figure 12.2-Arrangement of ions in
sodium chloride.
Ionic Bond
Sodium Chloride
Chemical reaction
between solid sodium
with gaseous chlorine
from a macroscopic
level.
Illustration of the chemical reaction
between sodium and chlorine.
Ionic Compounds (cont’d)
Ionic bonds
It is a strong attractive force between ions in a crystalline
structure.
These strong attractive forces are precisely positioned in its
respective crystalline structure.
Crystal
A geometrical structure where the ions are arranged where
ions are attracted to one another.
The form of the crystal is dependent on the kinds of ions in
the compound, their sizes, and their corresponding ratio.
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Ionic Compounds (cont’d)
Typically, ionic compounds typically very strong and
exist naturally at room temperature as solids.
At extreme high temperatures, if the solid ionic crystal
is melted, the ionic bond holding the ions are locked in
because since the crystal is destroyed, the ions in the
bond are not able to be free from one another thus will
be poor conductors of electricity.
However, if the ionic compound is soluble in water, the
ions can interact with the water solvent. As a result,
liquid ionic compound solutions (aqueous matter) are
good conductors.
Illustration
Figure 12.4-Electricial conductivity. (a) Ionic solids do not conduct electrical current
but (b) are good conductors when melted or dissolved in water.
Covalent Bonds
Figure 12.5-Conductivity of solutions. Ions are held together by ionic bonds in salt,
but another type of bonding must be responsible for the properties of sugar.
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Covalent Bonds
Molecular compounds are defined as substances
comprised of nonmetal atoms.
A covalent bond is one where electrons are
shared between atoms to reach a noble gas
configuration.
RECALL: An ionic bond is one where electrons
are transferred from a metal atom to nonmetal
atom to form a cation and anion where each
obtains a noble gas configuration.
Illustration
Covalent Bond
H
+
H
H
H
H
H
The two dots or the straight line drawn
between the two atoms represent the
covalent bond that holds the atoms together.
Covalent Bond (cont’d)
The electron cloud or charge density formed by the
two electrons is concentrated in the region between
the two nuclei.
The atomic orbitals of the hydrogen atoms overlap,
coupling them together to form the hydrogen
molecule.
When bonding electrons are between two nuclei, both
nuclei are attracted to the electrons. The electrons are
the “glue” that bonds the atoms to each other.
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Illustration
Covalent Bond (OVERLAP)
Figure 12.6-The formation of a hydrogen molecule from two hydrogen atoms. Each
location in the electron cloud represents and instantaneous position of the electron
in the atom. To the left of the arrow, the circle minus represent the instantaneous
position of the 1s electron in each separate atom at the same instant, before the
bond is formed. To the right of the arrow, the circled minus signs represent the
location of those valence electrons at the same instant as they make up a bond
between the atoms. The electron pair is shared by both atoms.
Covalent Bond
Fluorine Molecule
F
+
F
F
F
F
F
A half-filled 2p orbital from one F atom
overlaps a half-filled 2p orbital from the other
F atom.
Covalent Bond (cont’d)
Lewis diagrams, Lewis formulas, or Lewis
structures are illustrations where one can
illustrate the bonding arrangement between
atoms in a molecule.
Lone pairs are unshared electron pairs for an
atom not involving in the bonding.
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Covalent Bond (cont’d)
Octet Rule
Covalent bonds tend to form between nonmetal
atoms by filling the overlapping valence electron
orbitals with the maximum number allowed, two
in the s orbital and two in each of the three p
orbitals, for a total of eight (octa-) valence
electrons.
Covalent bonds tend to form when half-filled
orbitals overlap.
Polar and Nonpolar Covalent Bonds
There are three types of bonds we deal with
respects to how electrons are behaving between
atoms.
1. Pure Covalent Bond
a)
b)
c)
A bond in which electrons between atoms
are shared equally.
The charged density is centered in the region
between atoms in the molecule.
A bond between identical atoms are always
nonpolar.
Polar and Nonpolar Covalent Bonds
There are three types of bonds we deal with
respects to how electrons are behaving between
atoms.
2. Polar Covalent Bond
a)
b)
A bond in which electrons between atoms
are shared unequally.
The charged density is towards the region
which is more electronegative in the
molecule.
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Illustration
Polar Covalent Bond
Figure 12.7-A polar bond. Fluorine in a molecule of HF has a higher electronegativity
than hydrogen. The bonding electron pair is therefore shifted towards fluorine. The
uneven distribution of charge yields a polar bond.
Illustration
Bonding
Figiure 12.8-A comparison of nonpolar covalent, polar covalent, and ionic bonding.
Electronegativity
Bond polarity in covalent bonds may be described in
terms of the electronegativities of the bonded atoms.
Electronegativity
The ability of an atom of an element in a molecule to
attract bonding electron pairs to itself.
High electronegativity identifies an element with a
strong attraction for bonding electrons.
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Electronegativity
Figure 12.9-Periodic trends in electronegativity values. Electronegativity values
generally increase from left to right across any period of the table, and they decrease
from the top to the bottom of any group.
Electronegativity
Trend in electronegativity values of common
elements:
F > O > N = Cl > Br > C = I > H
(In this context, “=“ means the same to one
tenth of an electronegativity unit.)
Electronegativity
Classify each pairs of atoms based on their
electronegative difference.
Pure Covalent Bond
the electronegative difference between atoms is
between 0-0.3
Polar Covalent Bond
the electronegative difference between atoms is
between 0.4-1.8
Ionic Bond
the electronegative difference between atoms is greater
than 1.8
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Examples
Determine the electronegativity difference in a
typical C-H bond. Additionally, classify the type
of bond illustrated by its difference.
ANSWER: C has a 2.5 electronegativity value.
H has a 2.1 electronegativity value.
δ−
∆EN = 2.5-2.1 = 0.4
A Polar Covalent Bond
C
H
A dipole arrow points
to the atom with the
higher
electronegative
value.
C
δ+
H
δ- is labeled above the
atom whose
electronegative value is
large, δ- is label to the
lower electronegative
value.
Multiple Bonds
In illustrating the covalent bonds between any pair of
atoms in a covalent molecule, there are three types of
bonds:
1.
Single Bond
2.
Double Bond
3.
Triple Bond
Two electrons are share between two nearby atoms.
Four electrons are share between two nearby atoms.
Six electrons are share between two nearby atoms.
Multiple Bonds
Multiple bonds defines a molecule which have
more than one type of bond visually observed
in a molecule.
NOTE:
Triple Bd. > Double Bd. > Single Bd.
Bond Strength Strong
Weak
Energy
Large
Small
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MOLECULES
Central Atom (two or more atoms)
If you have two or more different atoms on a
central atoms, the objective is two obtain 8
electrons (shared or unshared.
EXAMPLE:
Formation of a water molecule
H
+
O
+
H
H
O H
or
H
O
H
Other Molecules
H
H
C
C
H
O
C
O
H
H
ethylene
C
C
H
acetylene
carbon dioxide
Octet Rule
Exceptions
There are exceptions to the octet rule.
1) Odd-Electron Molecules “RADICALS”
These are molecules which has an odd total
number of electrons.
N
N
O
nitrogen monoxide
5e- + 6e- = 11 valence e’s
O
O
nitrogen dioxide
5e- + 2(6e-) = 17 valence e’s
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Octet Rule
Exceptions
There are exceptions to the octet rule.
2) Molecules with More Than 4 Electron Pairs
Around the Central Atom
The central atom has more than 8 electrons in the
structure.
Cl
Cl
Cl
Cl
Cl
P
S
Cl
Cl
Cl
Cl
Cl
Cl
phosphorus pentachloride
5e- + 5(7e-)= 40 valence e’s
sulfur hexachloride
6e- + 6(7e-) = 48 valence e’s
Octet Rule
Exceptions
There are exceptions to the octet rule.
3) Molecules with Fewer Than 4 Electron Pairs
Around the Central Atom
The central atom has less than 8 electrons when shared between
atoms.
Cl
B
Cl
Cl
boron trichloride
3e- + 3(7e-) = 24valence e’s
Octet Rule
Exceptions
Special case: OXYGEN
3) In nature, oxygen is drawn on paper has a
double bond but liquid oxygen is paramagnetic,
meaning that it is attracted to a magnetic field.
O
O
diamagnetic
O
O
paramagnetic
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Illustration
OXYGEN
Figure 12.10-Physical properties of oxygen. When pure gaseous oxygen is cooled to
temperatures between -183oC and -218oC, it becomes a liquid (left). Liquid oxygen has
a pale blue color (center). If the liquid is poured into the field of a strong magnet
Metallic Bonds
A metallic bond is an attractive force which is
mainly between positive charged metal ions
and the negatively charge valence electrons
that move among them.
Metallic Bonds
Figure 12.11-The electron-sea model of a metallic crystal. The monatomic ions formed
by the metal remained fixed in a definite crystal pattern, but the highest-energy
valence electrons are relatively free to move, which explains the high electrical
conductivity of metals.
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Metallic Bonds
In a metallic bond, electrons are delocalized
because electrons do not stay near any single
atom or pair of atoms.
An alloy is combination of two or more metals
which is dependent on the composition of each
element in the mixture.
EXAMPLES:
Steel (Iron, Aluminum, Chromium)
Bronze (Copper, Tin)
Brass (Copper, Zinc)
Sterling Silver (Silver, Copper)
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