Kotz: 6th edition

Module 4: I & II Chapter 9

Bonding and Molecular Structure:  Fundamental Concepts

By

 Dr Andrea Wallace

CGCC

 

Edited by

John Taylor

FCCJ-North Campus

Chapter 9:  Bonding and Molecular Structure:  Fundamental Concepts

 

9.1  Valence Electrons, p. 374  (CD-Rom, Screen 9.2)

 

 

________________________ - electrons in the outermost shell that participate in bonding.

 

________________________ - inner electrons the do not participate in bonding.

 

For main group elements, # of valence electrons = ________________.

 

Problem 9.1, p. 376 (modified)

Give the number of valence electrons in

Ba

 

As

 

Br

 

In Lewis dot or electron dot symbols, dots are used to represent valence electrons.

 

Lewis dot symbols emphasize full octets.  Full octets are accomplished by losing, gaining, or sharing electrons.

 

Problem 9.1, p. 376 (modified)

Provide Lewis dot symbols for the following

 

Ba                    As                    Br

 

 

See Table 9.2, p. 375 for more Lewis dot symbols.

 

9.2    Chemical Bond Formation, p. 376 (CD-Rom, Screen 9.6)

 

__________________________ - occurs when there is a transfer of electrons and occurs between a metal and a nonmetal.

 

 

 

 

 

 

Metal              Nonmetal                                      Ionic Compound

 

 

Ionic Lewis Structures

  1. Write the metal first with no dots and the correct positive charge and subscript.
  2. Write the nonmetal last with brackets and 8 dots and the correct negative charge and subscript.

 

Provide an ionic Lewis Structure for Al2S3.

 

 

 

 

 

 

 

______________________ is a sharing of electrons by two bonded nuclei.  Occurs between two nonmetals.  (Hydrogen is considered to be a nonmetal in this case.)

 

9.3  Bonding in Ionic Compounds, p. 377

(Skip)

 

9.4    Covalent Bonding and Lewis Structures, p. 382 (CD-Rom, Screen 9.7)

 

Provide a Lewis Structure for I2  (A full octet must be achieved.)

 

 

 

 

 

_______________ is  a nonbonding pair of electrons.

 

Provide a Lewis Structure H2. 

 

 

 

 

 

A single bond is a sharing of ____ pair of electrons.

A double bond is a sharing of ____ pairs of electrons.

A triple bond is a sharing of ____ pairs of electrons.

 

Give examples of molecules with multiple bonds.

 

 

 

 

 

 

Guidelines for Drawing Lewis Structures  (CD-Rom, Screen 9.8)

 

1.      Determine if your species is ionic (metal + nonmetal) or covalent (all nonmetals – consider H to be a nonmetal)

  1. If ionic, proceed to Step #2.
  2. If covalent, proceed to Step #3 for a two atom molecule and proceed to Step #4 for a three or more atom molecule.

 

2.      Ionic Lewis Structures

a.       Write the metal first with no dots and the correct positive charge and subscript.

b.      Write the nonmetal last with brackets and 8 dots and the correct negative charge and subscript.

 

3.      Covalent Lewis Structures (2 atoms)

a.       Add up the number of valence electrons.  For positively charged species, subtract electrons.  For negatively charged species, add electrons.

b.      Try a single bond, check to make sure each atom has 8 electrons around it.  If not, try a double bond.  If this doesn’t work either, try a triple bond.

 

4. Covalent Lewis structures (3 or more atoms)

  1. Place the atom with the lowest electron affinity at the center and surround it with the other atoms.  This is usually the atom you have the least of.  (Note:  H is always an outer atom.  Note:  With some larger organic molecules there may be more than one central atom.)
  2. Add up the number of valence electrons.  For positively charged species, subtract electrons.  For negatively charged species, add electrons.
  3. Singly bond all atoms together.
  4. Remaining electrons are used to provide a full octet for all outer atoms (except H which has a full octet with 2 electrons). 
  5. Further remaining electrons are placed on the central atom.  (Note:  Central atoms from period 3 or greater may have more than 8 electrons around them.  Note:  When Be and B are central atoms, they may have less than 8 electrons around them.)
  6. Check.  If the central atom has less than a full octet (except for Be and B), convert lone pairs into double or triple bonds. (Note:  In general, only C, N, O, Si, P, S, and Se can form multiple bonds.)

 

 

 

Provide Lewis Structures for all of the following:

 

CO

 

 

 

 

 

CN-1

 

 

 

 

 

NH3

 

 

 

 

 

NO3-1

 

 

 

 

 

 

SO4-2

 

 

 

 

 

CH3OH

 

 

 

 

 

 

C and Si take ______ bonds.

N and P take ________ bonds.

O, S, and Se take _______ bonds.

F, Cl, Br, I, and H take _______ bond.

 

 

9.5    Resonance, p.  390, (CD-Rom, Screen 9.9)

 

___________________________ - contributing structures with different arrangements of electrons.

 

Example:  Ozone, O3

 

 

 

 

 

 

You would except to see different bond lengths, O=O should be shorter, O-O should be longer.  However, the bond distances are identical – 127.8 pm.

 

Why?

Both structures contribute, but it is their composite which is called the ___________ _____________________ which is the true structure.

 

 

 

 

 

 

 

Exercise 9.7, p. 392

Draw resonance structures for the nitrate ion, NO3-1.  Sketch a plausible Lewis dot structure for nitric acid.

 

 

 

 

 

 

9.6  Exceptions to the Octet Rule, p. 392

 

Fewer than 8 electrons (CD-Rom, Screen 9.10)

 

            Acceptable #                            Compound

 

 

B

 

Be

 

H

Example of a reaction using a Boron compound.  A _____________________ bond is formed – this is a bond in which both bonding electrons are contributed by the same atom.

 

Reaction:

 

 

 

 

 

 

 

 

More than 8 electrons

Occurs in element in periods 3 or greater only when they are acting as the central atom.

 

Examples:

 

 

PCl5

 

 

 

 

XeF2

 

 

 

 

ClF4-1

 

 

 

 

Molecules with Odd #’s of Electrons (CD-Rom, Screen 9.11)

 

Having unpaired electrons makes a species extremely reactive and they are referred to as _________________________.

 

Examples:

 

NO

 

 

 

NO2

9.7  Molecular Shapes, p. 397

 

Method for predicting molecular shapes is VSEPR.

 

VSEPR – Valence shell electron pair repulsion theory – based on the idea that electron pairs prefer to be as far apart from each other as possible due to electron repulsion.

 

See p. 398, 399, 400 and 402 for shapes.

 

Their shapes are based on their number of electron density groups.  An electron density group can be either a 1) lone pair or a 2) bonding group.

 

Electron Pair Geometry – is the geometry taken up by all the valence electron pairs around a central atom.

 

Molecular Geometry – describes the arrangement in space of the central atom and the atoms directly attached to it.

 

Multiple bonds to outer atoms count as only one bonding group.

 

Find the electron pair geometry, molecular geometry, and bond angles for the following:

 

BF3

 

 

 

BF4-1

 

 

 

H2O

 

 

 

NH3

 

 

 

ICl2-1

 

 

 

PO4-3

 

 

 

SO3-2

 

 

 

IF5

 

 

 

XeF2

 

 

 

NO2

 

 

 

Consider cysteine (p. 404).  Predict the Molecular Geometry about each central atom.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

9.8   Charge Distribution in Covalent Bonds and Molecules, p. 405

 

Oxidation # - charge an atom would have it all of its bonds were considered to be ionic.  It does not necessarily represent real charges.

 

Find the oxidation # of each atom in the following:

 

SF4

 

 

CO3-2

 

 

 

Formal Charge – assumes that each bond pair is shared equally by two atoms – covalent bonding.

 

Formal Charge = Group #  - # of nonbonded electrons - # bonds

 

Find the formal charges for all of the atoms in CO3-2

 

 

 

 

Formal Charges allow you to choose the best Lewis Structure.   The best Lewis structure will always be the one with the formal charges as close to zero as possible.

 

 

 

 

Draw the best Lewis structures for CO2 and SO4-2.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Bond Polarity and Electronegativity

 

Electronegativity – of an atom in a molecule is the measure of the ability of the atom to attract electrons to itself.   (Follows the same trend as ionization energy and electron affinity.)  The higher the electronegativity value, the stronger the attraction.  See p. 409, Figure 9.14, for a Table of Electronegativity values.

 

 

 

Type of Bonding

Definition

Charges on Atoms

Electronegativity Difference

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

What type of bonding occurs in the following?  (Calculate electronegativity difference and show charges where applicable.)

 

 

H-O

 

 

 

C-H

 

 

 

Cs-Cl

 

 

 

 

H-Cl

 

 

 

C-C

 

 

 

 

Exercise 9.15, p. 413

Consider all possible resonance structures for SO2.  What are the formal charges on each atom in each resonance structure?  What are the bond polarities?  Do they agree with formal charges?

 

 

 

 

 

 

 

9.9  Molecular Polarity, p. 413

 

Polarity results from an unequal sharing of electrons.

 

Dipole moment is a measure of polarity.  The larger the value, the more polar the molecule.

 

See Table 9.8, p. 414.

 

A molecule is polar if

1)      There are lone pairs on the central atom.

2)      It contains polar bonds which do not cancel out.   (Canceling out occurs if all of the bonding groups are identical.)

 

Determine if the following molecules are polar or nonpolar:

 

CO2

 

 

 

 

BF3

 

 

 

 

NH3

 

 

 

 

H2O

 

 

 

 

CFCl3

 

 

 

 

CF4

 

 

 

 

9.10          Bond Properties:  Order, Length, and Energy, p. 419

 

Bond Order - # of bonding electron pairs shared by two bonded atoms in a molecule.

 

Type of Bond

Bond Order

 

 

 

 

 

 

 

Fractional Bond Order – occurs when resonance structures exist.

 

Bond Order =      # shared pairs linking X and Y_____

                        # of X-Y links in the molecule or ion

 

 

Calculate the bond order of a C-O bond in CO3-2.

 

 

 

 

 

 

 

Bond Length – distance between the nuclei of two bonded atoms.  It is based on two factors:

1)      Bond Order

2)      Atoms Size

 

1) Bond Order

Bond

C-O

C=O

C triple O

Bond Order

1

2

3

Bond Length in pm

143

122

113

 

                                                ί____________________________ -----------------

 

2)      Atom Size

 

C-N

C-C

C-P

147 pm

154 pm

187 pm

 

---------____________________________ ΰ

 

 

 

 

 

Average Bond Lengths are given in Table 9.9 on p. 420.

Exercise 9.17, p. 421

a)      Give the bond order of each of the following bonds and arrange them in order of decreasing bond distance:  C=N, C triple N, C-N

 

 

 

 

 

 

b)      Draw resonance structures of NO2-1.  What is the NO bond order in this ion?  Consult Table 9.9 for for N-O and N=O bond lengths.  Compare these with the NO bond length in NO2-1 (124 pm).  Account for any differences you observe.

 

 

 

 

 

 

 

 

 

 

 

 

 

Bond Dissociation Energy (D) – energy required to break a bond.

 

See Table 9.10, p. 422.

 

The shorter the bond length, the greater the energy needed to break the bond.

 

Bond

Bond Dissociation Energy

C-C

346 kJ/mol

C=C

610 kJ/mol

C triple C

835 kJ/mol

 

The process of breaking a bond is _____________________.

The process of forming a bond is ______________________.

 

DHrxn  = Summation of D (bonds broken) – Summation of D (bonds formed)

 

 

 

 

 

 

 

Exercise 9.18, p. 424

Using the bond energies, in Table 9.10, estimate the heat of combustion of gaseous methane, CH4.  That is, estimate DHrxn for the reaction of methane with O2 to give water vapor and carbon dioxide gas.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

9.11         The DNA Story – Revisited, p. 424

 

Tetrahedral Structure of Carbon in DNA double helix – see p. 425, Figure 9.20