1 Chapter 3 Chemical Bonding Ref: Ralph H. Petrucci, F. Geoffrey Herring, Jeffry D. Madura, Carey Bissonnette. General Chemistry: Principles and Modern Applications. Prentice Hall 2010, English 10 th ed. Theodore E. Brown, H. Eugene H LeMay, Bruce E. Bursten, Catherine Murphy, Patrick Woodward. Chemistry: The Central Science. Prentice Hall 2011, English 12 th ed. Raymond Chang and Jason Overby. General Chemistry: The Essential Concepts. McGraw-Hill Science/Engineering/Math 2010, 6 th ed.
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Chapter 3 Chemical Bonding - Kittisak Choojun€¦ · Electronegativity is a measure of an atom’s ability to attract electrons from a neighboring atom to which it is bonded; it
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1
Chapter 3
Chemical Bonding
Ref: Ralph H. Petrucci, F. Geoffrey Herring, Jeffry D. Madura, Carey Bissonnette. General Chemistry:
Principles and Modern Applications. Prentice Hall 2010, English 10th ed.
Theodore E. Brown, H. Eugene H LeMay, Bruce E. Bursten, Catherine Murphy, Patrick Woodward.
Chemistry: The Central Science. Prentice Hall 2011, English 12th ed.
Raymond Chang and Jason Overby. General Chemistry: The Essential Concepts. McGraw-Hill
Science/Engineering/Math 2010, 6th ed.
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Lewis Theory1. Electrons, especially those of the outermost (valence) electronic shell, play a
fundamental role in chemical bonding.
2. In some cases, electrons are transferred from one atom to another. Positive
and negative ions are formed and attract each other through electrostatic forces
called ionic bonds.
3. In other cases, one or more pairs of electrons are shared between atoms. A
bond formed by the sharing of electrons between atoms is called a covalent
bond.
4. Electrons are transferred or shared in such a way that each atom acquires an
especially stable electron configuration. Usually this is a noble gas configuration,
one with eight outer-shell electrons, or an octet.
Lewis Theory
A Lewis symbol consists of a chemical symbol to represent the nucleus and core
(inner-shell) electrons of an atom, together with dots placed around the symbol to
represent the valence (outer-shell) electrons.
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Lewis Structure
Write the Lewis Structure for nitrogen trifluoride (NF3), nitric acid (HNO3),
carbonate (CO32-)
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Some Exceptions to the Octet Rule
1. The Incomplete Octet
Fewer than eight :
Group 2A (stable with 4 e-), BeH2,
Group 3A (stable with 6 e-), BF3
2. Odd-Electron Molecules: NO, NO2
3. The Expanded Octate
Atoms of the second-period elements cannot have more than eight
valence electrons around the central atom, but atoms of elements in
and beyond the third period of the periodic table form some
compounds in which more than eight electrons surround the central
atom. SF6, PCl5, XeF4
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Formal Charge and Lewis Structure
Formal Charge is the electrical charge difference between the valence
electrons in an isolated atom and the number of electrons assigned to the atom
in a Lewis Structure.
FC = [(number valence e- in free atom) – (number lone-pair e-) - ½
(number bond-pair e-)]
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Formal Charge and Lewis Structure
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Calculate the formal charge of the following molecules:
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Resonance
The two resonance structure are equivalent.14
Resonance
Which resonance form of azide anion, N3- are the most stable.
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Covalent Bonding
Bonding between a hydrogen atom and a chlorine atom involves the sharing of
electrons, which leads to a covalent bond.
The sharing of a single pair of electrons between bonded atoms produces a
single covalent bond.
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Coordinate Covalent Bonds
The Lewis theory of bonding describes a covalent bond as the sharing of a pair of
electrons, but this does not necessarily mean that each atom contributes an
electron to the bond. A covalent bond in which a single atom contributes both of
the electrons to a shared pair is called a coordinate covalent bond.
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Multiple Covalent Bonds
The Covalent Bond
A bond in which two electrons are shared by two atoms
Low ∆EN, H2, F2,
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Types of Covalent Bonds
Single Bond – two atoms are held together by one electron
pair.
Multiple bonds – two atoms share two or more pairs of
electrons.
Double bond – two atoms share two pairs of electrons
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Types of covalent bonds
Multiple bonds – two atoms share two or more pairs of electrons.
Triple bond – two atoms share three pairs of electrons
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Electronegativity
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Polar Covalent Bond
Even though, covalent bond is formed from the two atoms sharing
their electrons, if the atoms have different EN the electrons
will spend more time in the vicinity of one atom that has higher
EN than the other.
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Bond Enthalpy
A measurement of the stability of a molecule is its “bond enthalpy”
which is the enthalpy change required to break a particular bond
in 1 mole of gaseous molecules.
Bond enthalpies in solids and liquids are affected by neighboring
molecules.
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Bond Enthalpy
Measuring the strength of covalent bonds in polyatomic molecules
is more complicated.
In each case, and O-H bond is broken, but the first step is more
endothermic than the second. The difference between the two ∆Ho values
suggests that the second O-H bond itself has undergone change, because of
the changes in the chemical environment. Thus, for polyatomic molecules
the bond enthalpy is the average of a particular bond.
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Shapes of Molecules
Think about the molecules in 3D
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Valence shell electron-pair repulsion
(VSEPR)
The basis of the VSEPR approach is that electrons repel each
other because they are negatively charged.
Molecules adopt geometries such that valence electron pairs
position themselves as far from each other as possible to
minimize electron–electron repulsions.
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Valence shell electron-pair repulsion
(VSEPR)
A is the central atom, X stands for any atom or group of atoms
surrounding the central atom, and E represents a lone pair of electrons.
The steric number (SN) = m + n
CO2 --SN = 2
SO3 -- SN = 3
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Valence-Shell Electron-Pair Repulsion
(VSEPR) Theory
Central Atoms with Two Electron Groups
Central Atoms with Three Electron Groups
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Valence-Shell Electron-Pair Repulsion
(VSEPR) Theory
Central Atoms with Three Electron Groups
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Valence-Shell Electron-Pair Repulsion
(VSEPR) TheoryCentral Atoms with Four Electron Groups
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Lone-pair repulsion
The VSEPR model predicts that electron-pair repulsions involving
lone pairs ( lp ) are stronger than those involving bonding pairs ( bp ) in