Chapter Nine

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Chapter Nine

• Chemical Bonds

• IntraMolecular Force
• Forces hold atoms together in molecules and/or
  keep ions in place in solid ionic compounds.
• Caused by electrostatic forces that reflect a
  balance of attraction and repulsion between
  charged particles

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Contents
1. Electrostatic Interaction 2. Lewis Theory (1)
Lewis Dot Symbol (2) Lewis Theory of Chemical
Bonding Overview (3) Electronegativity (4)
Electronegativity Difference and Bond
Type 3. Ionic Bonding (1) Ionic Bonds and Ionic
Crystals (2) Lewis Structure for Ionic
Compounds (3) Lattice Energy
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4. Covalent Bonding (1) Lewis Structure for
Simple Molecules (2) Coordinate Covalent
Bond (3) Multiple Covalent Bond (4) Writing
Lewis Structure for Covalent Compounds (5)
Formal Charge (6) Resonance (7) Exception of
Octet Rule (8) Bond Order, Bond Length, Bond
Energy 5. Miscellaneous (1) Alkenes and
Alkynes (2) Polymers
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1. Electrostatic Interaction Covalent
interaction potential of two hydrogen atoms for
example
H nuclei far apart little attraction.
(Not in Textbook)
lt 74 pm repulsion increases
At 74 pm attraction max. energy min.
H nuclei closer attraction increases
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2. Lewis Theory

• Lewis Dot Symbols
• Combination of a chemical symbol for the element
  and valence electrons (the dots around the
  element symbol) of the atom
• The first four dots are placed singly on each of
  the four sides of the chemical symbol, and
  pairing the dots as the next four are added.

• For representative elements

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(2) Lewis Theory of Chemical Bonding Overview

• Valence electrons act a fundamental role in
  chemical bonding.
• Metals and Nonmetals combine valence electrons
  transferred from the metal to the non-metal atoms
  giving rise to ionic bonds.
• Example NaCl, NH3 and MgO.
• Nonmetals and Nonmetals combine one or more
  pairs of valence electrons are shared between the
  bonded atoms producing covalent bonds.
• Example H2O, NH3 and CH4.

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• In losing, gaining, or sharing electrons to form
  chemical bonds atoms tend to acquire the
  electron configurations of noble gases.
• - H, Li, Be duet rule
• - Representative elements octet rule (ns2np2)
• - Transition elements irregular

Example 9.1 Give Lewis symbols for magnesium,
silicon, and phosphorus.
Solution
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(3) Electronegativity (EN)
1) The character derived from ionization energy
and electron affinity, it is a measure of the
ability of an atom to attract bonding
electrons. 2) The greater the electronegativity
of an atom, the more strongly it attracts the
electrons in a chemical bond.

• General trend increases from bottom to top,
• left to right

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4) Paulings Electronegativities

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Example 9.4 Referring only to the periodic table
inside the front cover, arrange the following
sets of atoms in the expected order of increasing
electronegativity. (a) Cl, Mg, Si (b)
As, N, Sb (c) As, Se, Sb

• Solution
• EN increases from left to right
• Ans Mg lt Si lt Cl
• EN increases from bottom to top
• Ans Sb lt As lt N
• (c) As and Se, same period, As lt Se
• As and Sb, same group, Sb lt As
• Ans Sb lt As lt Se

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• (4) Electronegativity Difference (?EN)
• and Bond Type
• ?EN The difference in EN of bonded atom
• ?EN versus bond type

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• In general
• nonmetal-nonmetal 0 lt ?EN lt 1.5 covalent bond
• metal-nonmetal ?EN gt 2.0 ionic bond
• metal-metal all atoms low EN metallic bond
• ?EN versus percent ionic character

• For nonmetal-nonmetal bond
• ?EN 0 nonpolar (pure) covalent bond
• ?EN ? 0 polar covalent bond

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• Ways for expressing electron distributions of
  covalent bond

d
d
HH
HCl
HCl

• Polyatomic ion such as CO32
• Covalent bond between C and O
• Ionic compounds such as Na2CO3
• Ionic bond between Na and CO32

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Example 9.5 Use electronegativity values to
arrange the following bonds in order of
increasing polarity BrCl, ClCl, ClF, HCl,
ICl
Solution
EN and ?EN are EN 2.8 3.0 3.0 3.0 3.0
4.0 2.1 3.0 2.5 3.0 Br Cl Cl Cl
Cl F H Cl I Cl ?EN 0.2
0.0 1.0 0.9 0.5 The
order of increasing polarity is Cl Cl lt Br Cl
lt I Cl lt H Cl lt Cl F.
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• Ionic Bonding
• Ionic Bonds and Ionic Crystals
• The atoms may lose or gain electrons to acquire a
  noble gas configuration, for example

• Two ions formed between a metal and a nonmetal
  have opposite charges, they are strongly
  attracted to one another and form an ion pair,
• For example NaCl

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• The net attractive electrostatic forces that hold
  the cations and anions together called ionic
  bonds.

• The highly ordered solid collection of ions is
  called an ionic crystal.
• NaCl crystal formation for example

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• Lewis Structure for Ionic Compounds
• (Stoichiometry should be considered)

Example 1 For NaCl Solution
Ans
Example 2 For MgO
Ans
Example 3 For Li2O
Ans
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Example 9.2 Use Lewis symbols to show the
formation of ionic bonds between magnesium and
nitrogen. What are the name and formula of the
compound that results?
Ans Lewis structure
Name magnesium nitride Formula Mg3N2
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• Lattice Energy
• Born-Haber cycle
• An application of Hess's law to analyzing ionic
  compounds involved reaction energies, used
  primarily as a means of calculating lattice
  enthalpies.

Born-Haber cycle for elements Na and Cl for
example
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• Lattice Energy
• The enthalpy change (?H) that accompanies the
  formation of one mole of an ionic compound from
  its gaseous atomic ions.
• Lattice energy of NaCl for example
• Na(g) Cl(g) ? NaCl(s) ?H Lattice energy
  787 kJ
• large negative value of the lattice energy
  (exothermic, energetically favorable) makes ionic
  compound formation

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3) Example of calculating lattice energy from
Born-Haber cycle (NaCl for example)
1. Enthalpy of sublimation Na(s) ? Na(g) ?H1
107 kJ 2. Bond-dissociation energy ½Cl2(g) ?
Cl(g) ?H2 122 kJ 3. First ionization
energy Na(g) ? Na(g) e ?H3 496
kJ 4. Electron affinity Cl(g) e ? Cl(g)
?H4 349 kJ 5. Lattice energy Na(g)
Cl(g) ? NaCl(s) Lattice energy ?H5
? Overall Enthalpy of standard formation Na(s)
½ Cl2(g) ? NaCl(s) ?Hfo 411 kJ
?Hfo ?H1 ?H2 ?H3 ?H4 ?H5 411
kJ Lattice energy ?H5 ?Hfo (?H1 ?H2 ?H3
?H4) 787 kJ
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• Covalent Bonding
• Lewis Structure for Simple Molecules
• Lewis structure of molecule shows the bonded
  atoms with the electron configuration of a noble
  gas, that is, atoms of representative elements
  obey the octet rule (H obeys the duet rule).
• The shared electrons is counted for each atom
  that shares them.
• The shared pairs of electrons (two electrons) in
  a molecule are called bonding pairs, commonly
  represented by a dash ().
• The electron pairs which are not shared, called
  nonbonding pairs, or lone pairs.

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• Example of Lewis structure (Cl2)

• Non-metals of the second period (except boron)
  form a number of covalent bonds equal to eight
  minus the group number

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(2) Coordinate Covalent Bonds

• One atom provides both electrons of the shared
  pair to form a covalent bond called a coordinate
  covalent bond (also called dative bond).
• Examples

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(3) Multiple Covalent Bonds

• 1) Single bond one shared pair of electrons,
• represented by ().
• e.g., H Cl or HCl
• 2) Double bond two shared pairs of electrons,
  represented by ().
• e.g.,
• 3) Triple bond three shared pairs of electrons,
  represented by (?).
• e.g.,

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• Writing Lewis Structure for Covalent Compounds
• 1) Steps
• 1. Determine the total number of valence
  electrons.
• 2. Write a skeletal structure and connect the
  atoms by single dashes (covalent bonds).
• 3. Place lone pairs electrons around the terminal
  atoms to give each terminal atom (except H) an
  octet.
• 4. Place the remaining electrons around the
  central atom to give an octet.
• 5. If necessary, move one or more lone pairs of
  electrons from terminal atom to form a multiple
  bond to the central atom

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2) Skeletal structure (the arrangement of
atoms) 1. Hydrogen atoms are always terminal
atoms 2. The central atom of a structure usually
has the lowest electronegativity 3. Oxoacids such
as HClO4, HNO3, etc., hydrogen atoms are usually
bonded to oxygen atoms 4. Either molecules or
polyatomic ions usually have compact and
symmetrical structures
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Example 9.6 Write the Lewis structure of nitrogen
trifluoride, NF3.
Solution 1. Total of valence electron 1 N
atom, 3 F atoms
5 (37) 26
2. Skeletal structure EN of N is 3.0 EN of F is
4.0
3. Octet of terminal atoms remaining e 26
24 2
4. Assign remaining e to central atom
Ans
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Example 9.7 Write a plausible Lewis structure for
phosgene, COCl2.
Solution 1. Total of valence electron 4 6
(27) 24
2. Skeletal structure
3. Octet of terminal atoms (none of e remained)
4. Central atom is not octet
5. Form multiple bonds to complete octet of
central atom
Ans
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Example 9.8 Write a plausible Lewis structure for
the chlorate ion, ClO3.
Solution 1. Total of valence electron 7
(36) 1 26
2. Skeletal structure
3. Octet of terminal atoms remaining e 26
24 2
4. Assign remaining e to central atom
Ans
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(5) Formal Charge

• Definition
• Formal charge is the difference between the
  number of valence electrons in a free
  (uncombined) atom and the number of electrons
  assigned to that atom when bonded to other atoms
  in a Lewis structure
• Estimation

Formal charge
Number of valence electrons in the uncombined atom
Number of lone-pair electrons on the bound atom
Number of electrons in bonds to the atom


½
Simply expressed by FC V LP ½ BP
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• Select plausible Lewis structure by formal charge
• The most plausible Lewis structure is one with no
  formal charges in all atoms
• If the formal charge is required, it should be as
  small as possible
• The negative formal charges should appear on the
  most electronegative atoms.
• The adjacent atoms in a structure should not
  carry formal charges of the same sign

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4) Estimation example for formal charge on
individual atom
Lewis structure (a) is more plausible than Lewis
structure (b)
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Example 9.9 For the molecule COCl2, shown here as
structure (a). Show that structure (a) is more
plausible than (b) or (c).
Solution Formal charge estimation
0
0
0
0
1
1
2
2
0
0
0
0
Most plausible structure
Least plausible structure
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• (6) Resonance (Delocalized Bonding)
• The situation that a molecule or a polyatomic ion
  is presented by several plausible Lewis
  structures, called resonance structures,
• O3 for example

• The actual molecule or ion that is a hybrid of
  the resonance structures, called a resonance
  hybrid,
• O3 for example

• Electrons in resonance hybrid are spread out over
  several atoms, called delocalized electrons

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Example 9.10 Write three equivalent Lewis
structures for the SO3 molecule that conform to
the octet rule, and describe how the resonance
hybrid is related to the three structures.
Solution Total of valence electron
24 Skeletal structure
Complete Octet of terminal atoms
Resonance structures
The SO3 molecule is the resonance hybrid of the
three resonance structures.
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• Exception of Octet Rule

• Odd number of valence electrons
• These kind molecules called free radicals, must
  have unpaired electron, NO2 for example

• Incomplete octets
• Compounds of Be, B, and Al may present
  incomplete octets, BF3 for example

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3) Expanded valence shells For elements in the
third and higher periods (such as S, Cl, P), the
unfilled d orbital provide additional shared and
lone pairs position for valence electron.
Therefore, the central atom has more than eight
electrons around it, PCl5 and SF6 for examples
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Example 9.11 Write the Lewis structure for
bromine pentafluoride, BrF5.
Solution 1. Total of valence electron
7 (57) 42
2. Skeletal structure
3. Complete octet of terminal atoms
4. Assign remaining e to central atom
Ans
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(8) Bond Order, Bond Length, Bond Energy

• Definition

• Bond order (b.o.) The number of shared electron
  pairs in a bond single bond (b.o. 1), double
  bond (b.o. 2), triple bond (b.o. 3).
• Bond length is the distance between the nuclei of
  two atoms joined by a covalent bond.
• Bond (dissociation) energy (BE or D) The energy
  required to break one mole of a particular type
  of covalent bond in a gas-phase compound

• Trend in general
• Bond length Single gt Double gt Triple
• Bond energy Single lt Double lt Triple

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3) Some representative bond lengths and bond
energy

• Red marked Diatomic molecule, pure covalent,
  definite BE.
• Others Affected by the linked atoms or groups,
  it is average BE!!

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• Estimation of enthalpy of formation by tabulated
  BE
• (a) Visualization example for N2 2H2(g) ?
  N2H4(g)

• Estimation equation
• ?Ho SBE(reactants) SBE(products)
• (c) More accurate estimation is calculated by
  thermodynamic method (Chapter 6)
• SHo S?Hfo(products) - S?Hfo(reactants)

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Example 9.13 Estimate the length of (a) the
nitrogen-to-nitrogen bond in N2H4 and (b) the
bond in BrCl.
Solution (a) Lewis structure
From Table 9.1, NN single bond is 145 pm Ans
145 pm
(b) Lewis structure
From Table 9.1, ClCl is 199 pm, BrBr is 228 pm
Ans 214 pm
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Example 9.14 Use bond energies from Table 9.1 to
estimate the enthalpy of formation of gaseous
hydrazine. Compare the result with the value of
?Hf N2H4(g) from Appendix C.
Solution
SBE(reactants)
SBE(products)
?Hf of N2H4(g) value listed in Appendix C. is
95.40 kJ/mol.
Ans
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• 5. Miscellaneous
• Unsaturated hydrocarbon
• Definition
• The hydrocarbons with double or triple bonds
  between carbon atoms
• Subcategory
• Alkene hydrocarbon with one or more CC double
  bonds. For example ethene (C2H4), common name
  ethylene.
• Alkyne hydrocarbon with one or more CC triple
  bonds. For example ethyne (C2H2), common name
  acetylene.

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3) Examples of hydrocarbons Lewis structures
ethyne
ethene
ethane
4) Hydrogenation reaction

• Example of hydrogenation reaction application
• unsaturated fatty acid to saturated fatty acids

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• Polymer
• Polymer The long molecule consists of structural
  units and repeating units jointed together
  through chemical bonds.
• Monomer The units of the polymer
• Polymerization The process of converting these
  units to a polymer is called.
• Biological polymer macromolecules formed by
  repeating units in living organisms, such as
  proteins, nucleic acids, and polysaccharides.

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End of Chapter 9