CHAPTER 9: BONDING AND MOLECULAR STRUCTURE

CHAPTER 9: BONDING AND MOLECULAR

STRUCTURE

Understand the basic process of ionic bonding: identify ionic compounds and describe their internal structure and properties, understand how size of ions affects ionic properties, calculate lattice energy, and draw Lewis diagrams of ionic structures.

Identify covalent compounds and characterize their properties.

Draw Lewis structures of covalent substances including exceptions such as reduced and expanded octets, radicals, and resonance structures.

Define and predict trends in bond order, bond length, and bond dissociation energy. Use bond energy to predict enthalpy of a reaction.

Understand the concept of electronegativity and how it is used to predict polarity of individual bonds and entire molecules.

Use VSEPR Theory to predict the shapes of simple covalent molecules.

9.0 Objectives:

HW#1 – 29, 33, 38, 39, 40, 41, 43, 45 Valence e-, LEWIS STRUCTURES!

HW#2 – 47, 49, 51, 53, 95 Formal Charge, Polarity, Electonegativity

HW#3 – 69, 93, 109 Bond Energy

HW#4 – 73, 75, 77, 79, 81, 89, 99, 103 Molecular Geometry

HW#5 – 83, 85, 97 Polarity

Homework:

1. Bonding – definition “forces that hold atoms together” What role do the e- play?

2. Valence electrons vs. Core electrons MAIN GROUP Elements

Outermost “s” and “p” e-

Transition Metals Outermost “s” and “p” e- as well as (n-1) “d” e-

When in doubt, write the noble gas e- configuration

9.1 VALENCE ELECTRONS

3. Lewis dot diagrams of elements Diagrams that showcase valence e- Lewis says, “Place the first four dots

separately!”

Ex. Li, Be, B, C, N, O, F, NeLi, Be, B, C, N, O, F, Ne

9.1 VALENCE ELECTRONS

1. Ionic bonding – definition and Lewis representation of NaCl Bond between metal and nonmetal due to

“electrostatic interactions” Metal donates an e- Nonmetal accepts an e-

Ex. NaCl

9.2 CHEMICAL BOND FORMATION

2. Covalent bonding and Lewis representation of Cl2 Bond in which e- are shared Overlap of e- density between 2 orbitals Ex. Cl2


3. Continuum Complete ionic or complete sharing of e- is a

bit extreme; most bonding has uneven sharing of e- (sometimes ionic, sometimes covalent)

4. Other bond types Metallic bonding

Ex. Alloys


1. Steps in formation of NaCl 1. Na(g) Na+

(g) + e- E = +496 kJ/mol

2. Cl(g) + e- Cl-(g) E = -349 kJ/mol

3. Na+(g) + Cl-(g) [Na+, Cl-] E = -498 kJ/mol

Eoverall = -351 kJ/mol

9.3 BONDING IN IONIC COMPOUNDS


2. Lattice energy “energy for the formation of 1 mol of solid

crystalline ionic compound when ions in the gas phase combine”


3. Formula units Smallest whole number

Ratio -repeating unit of an ionic compound

1. Diagram of H2

H H Both want 1s2

H:H Share the electron pair

H—H Bonded stable H2

9.4 COVALENT BONDING AND LEWIS STRUCTURES

2. Orbital overlap diagrams of H2, HCl, Cl2


3. Terminology – single, double, and triple bonds, bonding pairs and nonbonding or lone pairs of electrons Single Bond: 2 e- shared between 2 atoms

Ex. H2

Double Bond: 4 e- shared between 2 atoms Ex. O2

Triple Bond: 6 e- shared between 2 atoms Ex. N2


Bonding Pairs: e- involved in bonding (See preceding examples)

Nonbonding (lone) pairs: e- that are not involved with bonding but help provide the octet for an atom Ex. Cl2


4. Octet Rule Noble-gas configuration “tendency for molecules/polyatomic ions to

have structures in which 8 e- surround each atom”

H, He have a “duet” Be has 4 electron max and B has 6 electron

max


5. Rules for drawing Lewis structures a. Choose a central atom

Usually the atom with the lowest e- affinity Usually makes a lot of bonds Halogens are generally terminal atoms

b. Count the total number of valence electrons Neutral Molecule: sum of valence e- for each atom Anions: sum of valence e- and negative charge Cations: valence e- minus the total positive charge


c. Draw a skeleton structure Use one pair of electrons to form a bond between each

pair of bound atoms

d. Place the remaining electrons to fulfill the octet rule Do this for each atom Hydrogen gets a duet


e. Lack of electrons: Requires multiple bonds (double, triple) Could be more than one multiple bond

f. Too many electrons: Verify that your structure is correct (octets for all?) Watch anions!


NOT suv….SOV!

S = O-VS = O-VS: SShared e- in bondsO: total # e- required for an OOctet

V: VValence e- for all elements

9.6 Lewis Structures of Some Simple Molecules

6. Diagrams of H2 F2 CH4 NH3 H2O HF OH- NH4

+


H2 F2 CH4 NH3 H2O HF OH- NH4+


7. Isoelectronic species: NO+ N2 CO CN-


9.5 RESONANCE

1. Definition Alternative and equivalent Lewis structure

“created” by shifting the e- in a structure Spinning Rim Analogy

2. Examples: NO3- and NO2

-

9.5 RESONANCE

9.5 RESONANCE

3. Experimental evidence says: “It’s a combination of both” There are however, MORE PREVALENT

resonance structures for some molecules

Benzene is the most classic of all resonance structures

9.5 RESONANCE

1. Reduced octets for H, B and Be Ex. BeCl2, BCl3 (Be = 4 e-, B = 6e-)

9.6 EXCEPTIONS TO THE OCTET RULE

2. Expanded octets: PF5 SF6 ClF4- XeF2

Watch these elements (and some others) for expanded octets: P, S, Cl, As, Se, Br, Kr, Xe


3. Radicals (paramagnetic): NO and NO2

Structure that has unpaired e- Extremely Reactive O2 ?!?!


3. Problems with Lewis structures Only show 2-D view life (chemistry) is 3-D Works for most molecules, but not all Doesn’t show how evenly/unevenly e- are

being shared


1. Definitions: polar and nonpolar bonds Nonpolar bonds: 2 e- in a bond are “evenly”

shared between the 2 atoms

Polar bonds: 2 e- in a bond are unevenly shared; one atom is taking more of the e- density; atoms have a partial charge

9.7 CHARGE DISTRIBUTION IN COVALENT BONDS AND

MOLECULES

POLARITY RULES

2. Electronegativity a. definition

(e/n): “ability of an atom to attract bonding e- to itself when the atom is in a molecule”

b. Table and Periodic trends See Pg.10 in Reference Booklet Increases going left to right and bottom to top

(Points toward Fluorine) ACS Periodic Table


MOLECULES


MOLECULES

3. EN – parameters Prediction of “Ionic Character”

Pure Covalent Pure Ionic

0 .5 1 1.5 2 2.5 3

In General: 0.0 < 0.5 Nonpolar0.5 ≤ 1.8 Polar Covalent> 1.8 Ionic


MOLECULES

4. Ex9.1 Arrange the following bonds in order of increasing polarity: F-Cl, F-F, F-Na


MOLECULES

5. Central atom in Lewis structure: Many times has a formal charge Making more/less bonds than it “normally”

does


MOLECULES

6. Formal Charge a. Definition and Use

Charge for an atom in a molecule based on premise that bonding e- are evenly shared

b. Calculating – equation Formal Charge = Group # - [Lone Pair e- + ½ Bonding

e-]


MOLECULES

c. Examples: OH- and NO3-


MOLECULES

Ex9.2 Calculate the formal charge on each atom in CO3

2- and NH4+


MOLECULES

7. Electroneutrality – Definition The e- in a molecule are distributed so that the

formal charge is minimal

Most Probable Lewis Structure = one with minimal formal charge

Negative charge should reside on the most electronegative element

Formal charge > +/- 2 is not likely


MOLECULES

a. Example: CO2


MOLECULES

b. Ex9.3 Use formal charge and the electroneutrality concept to determine the most likely structures for N2O and OCN1-


MOLECULES

1. Bond order a. definition and examples

“number of bonding e- pairs shared between 2 atoms” Usually an integer (1, 2, or 3) BOND ORDER = (# shared pairs linking X-Y)

(number of X-Y links in the molecule)

Ex. CH4, CO2

b. resonance structures Bond order are fractions e- residing over both locations evenly Ex. O3

9.8 BOND PROPERTIES

2. Bond Length – definition and examples Bond length: distance between nuclei of a

covalent bond (little variation) More Polar bonds = shorter length

Ex. C-C C=C C≡C

1.54Å 1.34Å 1.20Å

9.8 BOND PROPERTIES

1. Bond dissociation energy – definition and examples Bond Dissociation Energy (D):

“energy needed to break a covalent bond in the gas phase”

Higher Bond Order Higher Dissoc. Energy

9.8 BOND PROPERTIES

a. Estimating Enthalpy of reaction from bond energies – equation

Hrxn = D (bonds broken) - D (bonds formed)

Energy is required to break bondsEnergy is required to break bonds Energy is released when bonds are formedEnergy is released when bonds are formed

9.8 BOND PROPERTIES

9.8 BOND PROPERTIES

b. Example: Estimate the HR for the synthesis reaction between gaseous hydrogen and chlorine.

9.8 BOND PROPERTIES

c. Ex9.4 Estimate the enthalpy of reaction for the combustion of methane, CH4, to produce gaseous carbon dioxide and water vapor.

9.8 BOND PROPERTIES

9.9 MOLECULAR SHAPES

1. Gumdrops, toothpicks, and protractors A tasty way to practice chemistry at home!

I’ll use balloon models


2. VSEPR Theory and importance of shapes VSEPR: VValence SShell EElectron PPair RRepulsion VSEPR gives our 2-D Lewis structures LIFE (3-D) Geometry has HUGE impact on properties

Based on idea pairs of e- in bonded atoms repel one another Want to be as far apart as possible Gives shape

Electron Group Any collection of valence e- around an atom that repel other e-

Single unpaired e- Lone pair e- Bonding pairs of e- (1, 2, 3)


NOTATION:AXnEm

A = central atomX = terminal atomsE = lone pair e- on central atom


3. Single bonds, no unshared pairs of electrons Lewis structure, geometry and bond angles

of:

a. BeH2


BH3

CH4



4. Unshared pairs of electrons on the central atom a. NH3

b. H2O


c. analogs (H2S, PCl3, etc.)


Ex9.5 Predict the molecular geometry and bond angles of HOCl and SiO4

4-


6. Multiple Bonds a. CO2

b. H2CO


c. HCN


7. Ex9.6 Predict the molecular geometry and bond angles in the following species: C2H2 C2H4 ClO3

1- NO31-

N2O ONCl


7. Expanded octets a. definition and recognizing

More than 4 e- groups around a central atom Use formal charge to help guide Lewis

Structure/Geometry


b. Examples: PCl5

SF6


ClF5

XeF4


8. Ex9.7 Determine the molecular geometry and bond angles of ICl2-1 IF3 XeOF4 How does the electronic geometry differ from the molecular geometry for these species?


Song Time: “The Devoted Chemist”

Review and Visualization

9.10 MOLECULAR POLARITY

1. Nonpolar molecules – definition and examples, CH4 CO2 Molecule consisting entirely of nonpolar

bonds OR

Molecule with polar bonds that cancel one another out


2. Dipoles – definition and examples, NH3 H2O Molecule with separate centers of (+) and (-)

charge Polar bonds present with no canceling out

DEMO: So how do we KNOW that water has a bent geometry?


4. Dipole moments and vectors () = d

= magnitude of charged = distance

Measured in debye (D) Nonpolar: = 0 Polar: ≠ 0


5. Rules for determining polarity of a molecule. A molecule is a dipole if: (uneven balance of e- density)

RULES: 0. Draw Lewis Structure 1. Use VSEPR to predict molecular shape 2. Electronegativity to predict bond dipoles 3. Determine whether bond dipoles cancel to produce

nonpolar or combine to give polar molecule

KEY CLUES: 1. Polar Bonds 2. Lone Pair e- 3. 2 different atoms bonded to central atom


6. Ex9.8 Determine which of the following are dipoles: SO2 BF3 CO2 N2O ClO3

- ONCl NCl3 BFCl2 SCl2


6. Exceptions: XeF4


7. Properties of dipoles Special properties observed due to

interactions between molecules (intermolecular forces)

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CHAPTER 9: BONDING AND MOLECULAR STRUCTURE