Chemical Bonds

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Chapter 4 Chemical Bonds
Outline 1. Introduction 2. Lewis bonding
theory 3. Molecular shapes 4 Valence bond theory
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1. Introduction to Chemical Bonds

• Atoms of elements are combined to form a variety
  of compounds with various shapes. e.g.

• Questions
• How and Why atoms attach together?
• What are forces responsible for holding atoms
  together in molecules, or ions in crystals?
• Why atoms are combined in a certain fixed
  proportion?
• Why molecules adopt different shapes?

• Purpose of this chapter
• Introduce some simple theory to answer the above
  questions

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Stable Electron Configurations

• Fact Noble gases, such as helium, neon, and
  argon are inert they undergo few, if any,
  chemical reactions.
• Theory The inertness of noble gases results
  from their electron structures each (except
  helium) has an octet of electrons in its
  outermost shell.
• Deduction Other elements may alter their
  electron structures to become like those of noble
  gases.

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Examples

• Sodium can lose a valence electron. In doing so,
  its core electrons are like the noble gas, neon.

• Chlorine can gain an electron. In doing so, its
  electron structure becomes like argon.

Isoelectronic species
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Indeed, sodium reacts with chlorine to give
sodium chloride

• Sodium dropped in flask filled with chlorine gas
• Produces sodium chloride (i.e., table salt)
• Chlorine extracts electron from sodium

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2. Lewis Bonding Theory

• Atoms come together for a single reason
• to produce a more stable electron configuration.
• More stable electron configuration by either
  transferring or sharing electrons.
• A lot of atoms like to have 8 electrons in their
  outer (valence) shell.
• Octet rule.
• There are some exceptions to this rulethe key to
  remember is to try to get an electron
  configuration like a noble gas.
• Li and Be try to achieve the He electron
  arrangement.

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Lewis Symbols of Atoms

• Lewis Bonding Theory is built on the idea of
  valence electrons. To illustrate bonding, it is
  more convenient to use Lewis Symbols (electron
  dot symbols)
• Lewis Symbols of atoms
• Use symbol of element to represent nucleus and
  inner electrons.
• Use dots around the symbol to represent valence
  electrons.
• Put one electron on each side first, then pair.

For example,
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Exercise Write the Lewis Symbol for arsenic
(As), and sulfur (S)

• As

S
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Lewis (Electron-Dot) Symbols
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Lewis Symbols of Ions
Lewis symbols can also be used for ions

• Cations have Lewis symbols without valence
  electrons.
• Since they are lost in the cation formation.
• Anions have Lewis symbols with 8 valence
  electrons.
• Electrons are gained in the formation of the
  anion.

Example
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Ionic Bonds

• Formed between metal and nonmetal.
• Metals (main group) give all its valence
  electrons (to nonmetal) to form cations.
• Nonmetals gain electrons from metal to form
  anions with 8 valence electrons.
• Ionic bond results from to - attraction.
• Larger charge stronger attraction.
• Smaller ion stronger attraction.

e.g. Sodium reacts with chlorine to form sodium
chloride
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Bonding in sodium chloride

• Na and Cl- ions have opposite charges and
  attract each other.
• The resulting attraction is an ionic bond.
• Ionic compounds are held together by ionic bonds
  and exist in a crystal (Ions organize themselves
  in orderly manner)

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Other examples
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Exercise 1. Predict the formula of the compound
that forms between calcium and chlorine.
Lewis symbols of the elements
Transfer all the valance electrons from the metal
to the nonmetal, adding more of each atom as
you go, until all electrons are lost from the
metal atoms and all nonmetal atoms have 8
electrons.
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Exercise 2. Use Lewis symbol to show the electron
transfer from sodium atoms bromine atoms to form
ions with noble gas configuration.
Lewis symbols of the elements Sodium, group
1A. Bromine, group 7A.
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Octet Rule and Cation formation

• Cation formation Metals of main group usually
  lose all their valence electrons to give the
  electronic structure of the previous noble gas
  (except helium, octet in the outmost shell). In
  doing so, they form positive ions (cations).
• The charge of a cation from the main group metal
  is the same as the family number. e.g.
• Na, group 1A? Na
• Al, group 3A ? ___
• Exceptions.
• Transition metal may form cations not having the
  electronic structure of the previous noble gas.
• Transition metal may form more than one kinds of
  cation.
• e.g. Fe3, Fe2

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Octet Rule and anion formation

• Nonmetals tend to gain electrons to give the
  electron structure of the next noble gas (except
  helium, octet in the outmost shell). In doing so
  they form negative ions (anions).
• The charge of an anion from the main group
  elements is equal to the family number 8. e.g.

N, group 5A ? N3-
S, group 6A ? ___
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Octet Rule and ion formation

• Atoms of main group tend to gain or lose
  electrons to have electronic configuration of
  noble gases

Note Cations of transition metals (e.g. Fe2,
Zn2) may not have the electronic structure of
the previous noble gas (thus, octet rule is not
followed).
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Exercise. Electronic configuration of Fe2 and
Zn2?
Electronic configuration of transition metals in
Period 4
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Names of Binary Ionic Compounds

• Cation followed by anion
• The names of cations of main group are simply the
  name of the element. Examples
• Na sodium ion, Mg2 __________ ion
• The names of anions are the root name of the
  element plus the suffix ide. Examples
• Cl, chlorine, O, oxygen
• Cl- chloride ion, O2- _______ ion

Examples NaCl, sodium chloride Al2O3,
____________
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Names and charges of cations in Binary Ionic
Compounds of transition metals

• Many transition metals can exhibit more than one
  ionic charges. Roman numerals are used to denote
  the charge of such ions.
• Examples
• Fe2 iron(II) ion
• Fe3 iron(III) ion
• Cu2 ______ ion
• Cu _______ ion

FeCl3, ___________
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Formula of Binary Ionic Compounds

• Binary two components (cation anion)
• Cross-over method

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Covalent Bonds

• Many nonmetallic elements react by sharing
  electrons rather than by gaining or losing
  electrons.
• When two atoms share a pair of electrons, a
  covalent bond is formed.

• A shared electron pair form a a bond
• Shared electron pair bond bonding pair
• Electron pairs not involved in bonding
  nonbonding pairs

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Multiple Bonds

• Atoms can share one, two, or three pairs of
  electrons forming single, double, and triple
  bonds.
• Share two pairs double bond
• Share three pairs triple bond

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Naming Covalent Compounds
Covalent compounds have shared electron pair(s)
or covalent bond (s)

• Similar to ionic compound.
• Use prefixes to indicate the number of each kind
  of atom

N2O4 dinitrogen tetraoxide NO2 _______________
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Electronegativity of elements

• Measure of the attraction of an atom in a
  molecule for electrons
• Fluorine is the most electronegative element

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Polar Covalent Bonds

• When two atoms of different electronegativity
  form a bond, the bonding electrons are drawn
  closer to the atom with the higher
  electronegativity. Such a bond exhibits a
  separation of charge and is called a polar
  covalent bond.

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Types of Bonds

• The difference in electronegativity between two
  bonded atoms can be used to determine the type of
  bond. As a rule of thumb

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Covalent bond and Polyatomic Ions

• Covalent bonds can also be presented in
  polyatomic ions e.g.

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Main ideas of Lewis Bonding theory of covalent
compounds

• Nonmetal atoms tend to achieve an octet of
  valence electrons by sharing electrons.
• Though there are many exceptions to the Octet
  rule.
• This may involve sharing electrons with multiple
  atoms or sharing multiple pairs of electrons with
  the same atom.
• The shared electron pair provided the force
  linking atoms together.

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Exception to octet rule Molecules that Dont
Follow the Octet Rule
Molecules with an odd number of valence electrons
have at least one of them unpaired and are called
free radicals
An atom or molecule with an unpaired electron is
known as a free radical.
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Exception to octet rule Molecules that Dont
Followthe Octet Rule
-Some molecules have incomplete octets. These are
usually compounds of Be, B, and Al, generally
have some unusual bonding characteristics, and
are often quite reactive. e.g.
-Some compounds have expanded valence shells,
which means that the central atom has more than
eight electrons around it (may occur for elements
in period 3 or below).
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-An expanded valence shell may also need to
accommodate lone-pair electrons as well as
bonding pairs
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Drawing Lewis Structures
Suggested steps

• Determine the number of valence electrons.
• Each atom provide all its valence electrons
• (Only outshell ns and np electrons are
  considered for main group elements. v.e group
  for main group elements)
• a negative charge one more electron
• a positive charge one less electron
• b. Write the chemical symbols of the atoms in the
  arrangement that shows which bonds are formed.
  The less electronegative element is usually the
  central atom.
• c. Distribute the electron in pairs to achieve
  octet rule. Elements in period 3 and below could
  have more than 8 electrons in the valence shell.

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Suggested procedure to distribute electrons

• a. Put one pair between each pair of atom.
• Place electrons as lone pairs around outer atoms
  to fulfill the octet rule.
• c. Subtract the electrons used so far from the
  total number of valence electrons. Place any
  remaining electrons around the central atom.
• d. If the central atom lacks an octet, move one
  or more lone pairs from an outer atom to form a
  double or triple bond to complete an octet.
• Elements in period 3 or below may not follow
  the octet rule.

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Exercise

• Give the Lewis structure for the following
  compounds a) NH3, b) CO2, c) XeF4 and d) BF4-

• a) NH3 N 2s2 2p3 H 1s1
• No. of Valence electrons V.e 5(N) 3(H)
  8e-
• 2) Arrange atoms
• N
• H H
• H
• 3) Distribute electron pairs

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b). CO2

• C O
• 1) V. e
• 2) Arrange atoms
• 3) Distribute electrons

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c) XeF4

• Xe F
• 1). . V. e
• 2) Arrange atoms
• 3) Distribute electrons

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d) BF4-

• B F
• 1). . V. e
• 2) Arrange atoms
• 3) Distribute electrons

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Concept of Resonance More than one possible
Lewis structures. a single Lewis structure is
insufficient to electron delocalization. The
actual structure of a molecule is taken as a
blend of all the feasible Lewis structures. The
various structures are called Resonance Structure
Does NO3- has one short and two long bonds?
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3. Molecular Shape

• Molecules and ions can adopt various shapes.
  e.g.,

• Shapes can be important in determining
  reactivity, properties, especially in biological
  systems.

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Valence Shell Electron Pair Repulsion (VSEPR)
Theory

• The most stable arrangement of groups attached
  to a central atom is the one that has the maximum
  separation of electron pairs (bonded or
  non-bonded)

Assumptions
a). Each valence electron pair is important. A
multiple bond is treated as a single electron
pair. b). Repulsion of the electron pairs
determines the shape of molecules.
6 pairs
2 pairs
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No. of e- pairs and geometry
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No. of e- pairs and geometry
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PCl5
SF6
CX4
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Modifications and Examples

• Spatial requirement for lone pair (lp) is
  larger than those of bonding electron pair (bp).
  Thus the order of repulsion of electron pair is
  following
• lp/lp gt lp/bp gtbp/bp

The spatial requirement for a triple bond
electron pair is larger than that of a double
bond which is larger than that of single bond.
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General procedure to predict the shape of
molecules

• a). Draw a Lewis structure.
• b) Determine the number of electron groups around
  the central atom.
• Remember a multiple bonds count as one group.
• c). Determine the geometry of the central atom
  based on the number of electron pairs.
• d). Distribute the lone pairs and surrounding
  atoms such that the repulsion is minimized.
• e). Only atoms are then considered for the shape.

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Exercise. Predict the geometry of (a) H2O, (b)
NO2, (c) SO2 (d) NH3, (e) SF6, (f) SF4 and (g)
XeF4
(a) H2O
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(b) NO2
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(c) SO2
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(d) NH3
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(e) SF6
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(f) SF4
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(g) XeF4
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One more exercise. Predict the geometry of XeF2.
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Shapes and Properties Polar and Nonpolar
Molecules
A molecule is polar, if one part of the molecule
has a slight positive charge and the other a
slight negative charge. e.g.
e.g. A diatomic species of two different elements
Polar molecule has a dipole. The size of the
polarity is measured by Dipole moment.
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Shapes and Properties Polar and Nonpolar
Molecules
A polyatomic molecule can also polar, if

• 1. It has polar bonds.
• 2.The bonds are arranged such that a separation
  of charge exists. Have an unsymmetrical shape

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Molecule Polarity
The OC bond is polar. The bonding electrons
are pulled equally toward both O ends of the
molecule. The net result is a nonpolar molecule.
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Molecule Polarity
The HO bond is polar. Both sets of bonding
electrons are pulled toward the O end of the
molecule. The net result is a polar molecule.
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Adding Dipole Moments
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• Exercise. Decide Whether Each of the Following
  Molecules Is Polar

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4. Valence Bond Theory describes covalent bond
in terms of the overlap of atomic orbitals
A. Nature of the bond. The chemical bond is
formed by overlap of atomic orbitals of the two
atoms so that the two orbitals can be occupied by
a electron pair. e. g.
The electron density of shared electron pair is a
maximum between the bonded atoms. The electron
pair between the atoms provided the attractive
force for the bonding atoms.
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Comments on bond formation and types of bonds.

• (a) Only atomic orbitals that match with each
  other can overlap to form a bond. The better the
  overlap, the stronger the bonds.

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• (b) Type of bonds. Overlap of atomic orbitals can
  give chemical bonds having different symmetry
  properties.
• Based on symmetry properties, chemical bonds can
  be classified into ?, and p bonds.

s-bonds symmetric around the internuclear axis.
e.g.
p-bonds has one nodal plane through the
internuclear axis. e.g.
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B. Why electron pair? Energy consideration.
The energy of the system is lowered when the two
electrons of the bonding pair are of opposite
spin. e.g. in the case of H2.
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Another view
If the spin of the two electrons is parallel.
if the spin of the two electrons is antiparallel
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C. Number of bonds around an atoms

• Since pairing up the electrons lowers the
  energy, atoms with unpaired electrons tends to
  form chemical bonds to pair up the electrons.
• An atom can form covalent bonds with each of its
  stable valence orbitals.
• bonds unpaired e- (or orbitals with
  unpaired e-

Valence electrons may need to be re-distributed
to generate enough unpaired electrons.
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D. Molecular geometry
The shape of the molecules is determined
primarily by the directional characteristics of
the orbitals involved. e.g.
AOs can hybridize to generate orbitals of
required directional properties