Chemical Bonding
Chemical Bonds
are the forces that hold atoms together in compounds
Ionic bond
Covalent bond
Polar Covalent bond
Nonpolar Covalent bond
Metallic bond
Lewis theory,
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valence electrons play a fundamental role in chemical bonding
e- transfer leads to ionic bonds.
Sharing of e- leads to covalent bonds.
Atoms tend to have the electron configurations of the noble
gas, (octet rule.)
Lewis symbols : (electron dot symbols) is the simple way of showing the
valence electrons of atoms.
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• Si •
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P•
• As •
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• Sb •
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• Bi •
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Al
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Se
•
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I
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Ar
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•N •
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valence electrons are the electrons that occupy the outer shell (principal shell
number (n)) of an atom.
Core electrons are the electrons in inner shells.
Lewis Structures for Ionic Compounds
• O•
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2+ ••
Ba
O
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2-
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Ba•
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BaO
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Mg
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• Cl
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2 Cl
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-
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2+
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Mg •
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MgCl2
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• Cl
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İyonik Bileşikler
• Suda çözünürler
• Yüksek kaynama noktasına sahiptirler - iyonlar arası
çekim kuvveti yüksektir
• Eritilmiş halleri veya sulu çözeltileri elektriği iletir. İyonik
bileşikler katı halde elektriği iletmez.
Lewis Structures of covalent bonds
Lone pair (nonbonding or unshared) electrons are the electrons are not used for bonding.
Coordinate Covalent Bonds
If both electrons of the bond are contributed by the same atom, this type of bond
is called as a coordinate cocalent bond.
+
H
H
N
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H
H
H
Cl
H
H
H
Cu2+
N
+ 4 NH3
→
Cu(NH3)4
(type of covalent bond in which both electrons are donated by the same atom
Kovalent Bileşikler
• En temel yapıtaşları moleküllerdir
• Erime ve kaynama noktaları düşüktür
• Elektrik akımını iletmezler
• Çoğu suda çözünmez
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Bond Order
– Single bond, order = 1
Double bond, order = 2
Higher bond order
– Shorter bond - Stronger bond
• Multiple bonds In many molecules atoms attain complete octets by sharing
more than one pair of electrons between them.
The sharing of a pair of electrons represents a single covalent bond, usually
just referred to as a single bond,
if two electron pairs are shared it is a
double bond, if three electron pairs are shared it is called as a triple bond
• Bond length is the distance between the nuclei of two atoms joined by
covalent bond. As a general rule, the distance between bonded atoms
decreases as the number of shared electron pairs increases
• Bond energy (strength) is the quantity of energy required to break one mole
of covalent bond.
Bond energy (strength) is the quantity of energy required to break one mole of covalent
bond.
Bond Energies and Enthalpy of Reaction
ΔHrxn
= ΔH(product bonds) - ΔH(reactant bonds)
ΔHrxn
= ΔH(product bonds) - ΔH(reactant bonds)
= ΔH bonds formed - ΔH bonds broken
= -770 kJ/mol – (657 kJ/mol) = -114 kJ/mol
Bond polarity
In covalent bonds the atoms do not share the electrons equally. The atom
which has greater electronegativity, attract the electrons more strongly in
covalent bond.
nonpolar covalent bond
electrons are shared equally
between two atoms in a
chemical bond
polar covalent bond
one atom has a greater attraction for
the electrons than the other atom in a
chemical bond.
Electronegativity: is the relative ability of atoms to attract electrons in a
chemical bond.
if the electronegativity difference of atoms is 0, the bond is non-polar covalent
If the difference in electronegativities between the two atoms is greater than 0, but less than
2.0, the bond is polar covalent,
If the difference in electronegativities between the two atoms is 2.0, or greater, the bond is ionic
if the electronegativity difference of atoms is 0, the bond is non-polar covalent
If the difference in electronegativities between the two atoms is greater than 0, but less than
2.0, the bond is polar covalent,
If the difference in electronegativities between the two atoms is 2.0, or greater, the bond is ionic
Compound
Electronegativity
Difference
Type of Bond
F2
HF
LiF
4.0 - 4.0 = 0
4.0 - 2.1 = 1.9
4.0 - 1.0 = 3.0
Nonpolar covalent
Polar covalent
Ionic
(non-covalent)
Writing Lewis structure
• Skeletal structure shows how atoms are attached to one
another, the skeletal structure consist of one or more central
atoms and terminal atoms.
• A central atom bonds to two or more atoms in the structure.
The atom with the lowest electroneagtivity is generally central
atom.
• Terminal atom bonds to one another atom. Hydrogen atom
are terminal atoms.
In writing Lewis structure
• first determine the total number of valence electrons
• write a possible skeletal structure, connect the atoms by a single covalent
bond. (Molecules and polyatomic ions usually have compact, symmetrical
structures, but organic compounds are based on long chains of carbon
atoms. In oxyacids hydrogen atoms are usually bonded to oxygen atoms)
• place pairs of electrons as lone pairs around the terminal atoms, according
to octet rule.
• place the remaining electron as lone pairs around the central atom
• If there are not enough electrons to give the central atom an octet, try
multiple bonds (use one or more of the unshared pairs of electrons on the
atoms bonded to the central atom to form double or triple bonds, the double
bonds form generally among carbon, nitrogen oxygen and sulfur.)
• question: write a plausible Lewis
sturucture for phosgene COCl2 NO3-
Formal Charge is the differences between number of valence electrons in a free
(uncombined) atom and number of electrons assigned to that atom in the Lewis
structure.
FC= (number of valence electrons)- ½(number of bonding electrons)-(number of lone pair electrons)
-Total formal charges on the atoms in a lewis structure must be 0 to for a neutral atom
(and/or to the net charge of a polyatomic ion)
- negative formal charges should appear on the most electronegative atoms
- where formal charges are required, these should be as small as possible.
FC = #valence e- - #lone pair e- -
1
2
#bond pair e-
question : write a plausible Lewis structure and calculate
the formal charge each atom in that formula for each
compound given below,
NO3-, CO2, NH4+, HNO3, C2H4 SO42- PCl5, SF6
NO SOCl2 ICl4
Resonance:
sometimes a molecule or ion can be
represented by two or more plausible
Lewis structure that differ only by the
distribution of electrons. The different
plausible structures are called resonance
structures.
• O-O=O
 O-O=O
Exceptions to the octet rule
• Odd-electron species NO
• Incomplete octets BF3
• Expanded octets PCl5, SF6
H
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O—H
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H—C—H
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F
Cl
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Cl
Cl
P
F
Cl
F
Cl
F
Cl
F
S
P
Cl
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F
F
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Cl
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F
B
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+ F
F
+
-
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B
B
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F
F
F
F
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F
Dipole moments and molecular shape
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In many cases, covalent bonds are polar covalent because the bound
atoms have different electronegativities but whole molecule might be
nonploar. The "charge distribution" of a molecule is determined by
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A Polar Molecule:
The center of the overall negative charge on the molecule does not coincide
with the center of overall positive charge on the molecule
The molecule can be oriented such that one end has a net negative charge
and the other a net positive charge, i.e. the molecule is a dipole
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The shape of the molecule
The polarity of its bonds
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A Nonpolar molecule
Has no charges on the opposite ends of the molecule
Or, has charges of the same sign on the opposite ends of the molecule
Molecule is not a dipole
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question : determine the polarity of following molecules CH4, CH3Cl