chapter 8 – covalent bonding 8.1the covalent bond 8.2 naming molecules 8.3molecular structures...

Chapter 8 – Covalent Bonding

8.1 The Covalent Bond

8.2 Naming Molecules

8.3 Molecular Structures

8.4 Molecular Shape

8.5 Electronegativity and Polarity

Topic is Lewis Structures (combination of material found in 2 sections)

Sections 8.1/8.3 Covalent Bond/Molecular Structures

• Apply the octet rule to atoms that form covalent bonds.

• Describe the similarities and differences between ionic and covalent bonding.

• Describe the forces that act and energy changes that occur when atoms form a molecule.

Atoms gain stability when they share electrons and form covalent bonds. The sharing can be described by the Lewis structure of the compound.

• Categorize valance electrons as bonding or non-bonding.

• List the basic steps used to draw Lewis structures.

• Describe the formation of single, double, and triple covalent bonds using Lewis structures.

• Explain why resonance occurs, and identify resonance structures.

• Draw valid Lewis structures for molecules, including those involving multiple bonds, resonance, expanded octets, and electron deficient molecules.

Sections 8.1/8.3 Covalent Bond/Molecular Structures

Why Atoms Bond

Noble gas configuration especially stable

• ns2np6 (except for He)• Full outer energy level• Octet of electrons

Atoms bond to achieve a lower energy state (more stability)

Ionic vs Covalent Bonding

In ionic bonding, electrons transferred to achieve octet for each ion

• Number of ionic compounds small compared to total number of known compounds

In covalent bonding, electrons shared to achieve octet (mostly) for each atom

Covalent Bonding


• Sharing occurs when electronegativities of atoms same or similar

• Majority of covalent bonds formed between nonmetallic elements

• Electronegativity difference < 1.7 (see next slide) – bond will have more covalent character than ionic character

EN Difference & Bond Character

0 1.0 2.0 3.0

Electronegativity Difference

% Io

nic

Cha

ract

er

25

50

75

100

Ionic Bonds

Covalent Bonds

Covalent Bonding


Molecule formed when 2 or more atoms bond covalently

Covalent Bonding – Forces

Covalent Bonding – Forces

No interaction

Nucleus attracted to other atoms

electrons –

Not optimum distance

Net repulsion

from positive nuclei

Nucleus attracted to other atoms

electrons –

optimum distance

Covalent Bonding – Energy for H2

Internuclear Distance (pm)

100 200Pot

entia

l Ene

rgy

(kJ/

mol

)

-432 kJ/mol

Lewis Structures - Atoms (Electron Dot Diagrams)

Way of keeping track of valence electrons

To write for atom• Write symbol for element X• Put one dot for each

valence electron• Don’t pair up until you

have to (Hund’s rule)

Valence electrons of each element in molecule are divided into 2 categories:

• Bonding – pair of electrons shared by two atoms to form the covalent bond

Shared pair represented by a line connecting the element symbols H—H

• Nonbonding – called lone pairs A few molecules have odd # total

electrons – have unpaired nonbonding electron

Lewis Structure - Covalent Molecules

Lewis Structure - Covalent Molecules

Example – formation of H2 molecule

Bond = shared electron pair

H does not form octet

Space-Filling Model View Formation of H2

+

Bond = shared electron pair H_H

•HH••HH•

Ways of Representing Molecules: H2O

Structural Formula Space-Filling Model

Ball-and-Stick ModelOrbital Model

Ways of Representing Molecules: PH3

Covalent Bonding – F2

F 1s22s22p5

7 valence electrons

Forms F2 molecule

Each F shares 1 valence electron

Molecule is more stable than individual atoms

Lewis Structure - Covalent MoleculesExample – formation of F2 molecule

Bond = shared electron pairOctet

formed

Lewis Structure - Covalent MoleculesExample – formation of F2 molecule

Octet formed

Lewis Structure - Covalent MoleculesExample – formation of H2O molecule

Bonds = shared electron pairsOctet

formed

Two lone pairs

Shape of molecule

Lewis Structure - Covalent MoleculesExample – formation of ammonia, NH3

+ + +

Shape of molecule

Lone pair

Bonds = shared electron

pairs

Octet formed

Multiple Covalent Bonds

C, N, O, S often form multiple bonds

Double bond – O2 (6 valence e per O)

+

Triple bond – N2 (5 valence e per N)

+

Guide for Writing Lewis Structures

Step 1 – Write skeletal structure• least electronegative atom usually

occupies central position

Step 2 – Count total number of valence electrons

• polyatomic anions, add # of - charges e.g. CO3

2- add 2 electrons to total• polyatomic cations, subtract # of +

charges

Similar to procedure on p. 254, but without # of bonding pairs

Step 3 – Place single bond between central atom and surrounding atoms

Step 4 – Complete octet for terminal atoms (not for H)

Step 5 – Add remaining to central atom

Step 6 – If octet rule not satisfied for central atom, add multiple (double, triple) bonds between terminal and central atom, using the lone pairs from the terminal atoms

Lewis Structures – Common Bonding Patterns

C 4 bonds & 0 lone pairs

4 single (CH4), or 2 double (CO2), or single + triple (HCCH), or 2 single + double (CH2CH2)

N 3 bonds & 1 lone pair (NH3)

O 2 bonds & 2 lone pairs (H2O)

H & halogen 1 bond (CH4, CF4)

Be 2 bonds & 0 lone pairs (BeH2, electron def.)

B 3 bonds & 0 lone pairs (BH3, electron def.)

B C N O F

Lewis Structure Examples

Draw Single Bonds

Total Valence

Electrons

Calculate Number of Electrons

Remaining

Use Remaining

Electrons to Achieve

Noble Gas Configuration

Check Number of Electrons

a, HF 1 + 7 = 8 H-F 6

b, N2 5 + 5 = 10 N-N 8 c, NH3 5 + 3(1) = 8 2

d, CH4 4 + 4(1) = 8 0

e, CF4 4 + 4(7) = 32 24

f, NO+ 5 + 6 - 1 = 10 N-O 8

NH H

H

CH H

H

H

CF F

F

F

H, 2

F, 8

N,8

N, 8

O, 8

H, 2

N, 8

H, 2

C, 8

F, 8

C, 8

NH H

H

N N

FH

CH H

H

H

CF F

F

F

N O+

Practice

Problems 1-5 page 244

Problems 37-38, page 255

Problems 39-40, page 256 (mult bonds)

Problems 41-42 page 257 (ions)

Problems 104(a-d), page 275

Problems 1(a-d), page 979

Problems 4(a-e) page 980

Lewis Structure Example: NO3─

1. Write skeletal structure

N central because it is least electronegative

2. Count valence electrons

ONO

O

N = 5

3O = 3 x 6 = 18

(-) = 1

Total = 24 e-

Example NO3─ , Continued

3. Attach atoms with single bonds (pairs of electrons) & subtract from total ONO

O

——

Electrons

Start 24

Used 6

Left 18

4. Complete octets, outside-in

Keep going until all atoms have an octet or you run out of electrons

::

::

—— ONO

O

Electrons

Start18

Used18

Left 0


5. If central atom does not have octet, bring in electron pairs from outside atoms to shareIf structure is an ion, use brackets and indicate the charge

6. For this ion an extra step is needed – draw resonance structures

::

::

— ONO|

O

::

::

—— ONO

O


-1

Can have more than one correct Lewis structure for molecules or ions with double and single bonds

-


Resonance StructuresResonance structures differ only in position of electron pairs, never the atom positions

Molecule behaves as if it had only one structure (the average one)

• NO3- has all bond lengths

identical

-

Practice (Resonance Structures)

Problems 43-46 page 258

Problems 101,103 pages 274-5

Problems 5, 6 page 980

Practice—Lewis Structures

NClO

H3BO3

NO2-1

H3PO4

SO3-2

P2H4

NClO

H3BO3

NO2-1

H3PO4

SO3-2

P2H4

O P

O

O

O

HH

H

••

••

••

••

••

••

••

••

••

O S

O

O

••

••

•• •

•••

••

••

••

••

••

O N O ••

••

••

••

••••

18 e-

26 e-

32 e-

14 e- H P P H

HH

•• ••

O B

O

OH H

H••

••

••

••

••

••

24 e-

O N Cl ••

••

••

••

••••18 e-

Practice—Lewis Structures

*

*

* Has resonance structures

-1

-2

Exceptions to Octet RuleMolecules with odd number of total valence electrons

NO2 – 17 valence electrons

Also ClO2, NO

Exceptions to Octet RuleElectron deficient – form with fewer than 8 electrons around atom

• Be, B • Rare

+

Tend to form coordinate covalent bonds – both electrons in shared pair donated by single atom

Exceptions to Octet Rule

BeH2 – 4 electrons

BF3 – 6 electrons

Exceptions to Octet RuleMore than 8 valence electrons = expanded octet

PCl5 SF6

d orbitals involved• Only can occur for period 3 and higher,

not periods 1 or 2

Practice (Octet Exceptions)

Problems 47 - 49 page 260

Problems 102 (a-d), 104(a-d) page 273

Problem 7, page 980





8.4 Molecular Shape


Section 8.2 Naming Molecules

• Translate molecular formulas into binary molecular compound names and also the reverse process.

• Name acidic solutions

Specific rules are used when naming binary molecular compounds, binary acids, and oxyacids.

Naming Binary Covalent CompoundsFirst element named first, using entire element name

Second element named using same procedure as for ionic compounds – root of element name + ide ending

Use prefixes except if first element = 1• Drop final letter in prefix if precedes

vowel• Carbon monoxide , not monooxide

Prefixes in Covalent CompoundsTable 9-1, page 248

# Atoms Prefix # Atoms Prefix

1 mono- 6 hexa-

2 di- 7 hepta-

3 tri- 8 octa-

4 tetra- 9 nona-

5 penta- 10 deca-

Naming Binary Covalent Compounds

Name of AlCl3 ?Aluminum chloride

Name of PCl3 ?Phosphorus trichloride

Name of Al2O3?Aluminum oxide

Name of P2O5 ?Diphosphorus pentoxide

The naming systems for ionic and covalent compounds are different!!!

Common NamesTable 9-2, page 249

Formula

Common Name

Molecular Compound Name

H2O Water Dihydrogen monoxide

NH3 Ammonia Nitrogen trihydride

N2H4 Hydrazine Dinitrogen tetrahydride

N2O Nitrous oxide (laughing gas)

Dinitrogen monoxide

NO Nitric oxide Nitrogen monoxide

Naming Acids

For our purposes, acids are what result when molecules dissolved in water produce H+ (hydrogen ions)

• HCl(g) in water H+(aq) + Cl-(aq)• Product is hydrochloric acid

Two common types• Binary – H and one other element• Oxyacid – H and an oxyanion

Naming True Binary Acids

Use prefix hydro- to name hydrogen part of compound

For remainder, use a “form of the root” of 2d element plus suffix –ic followed by word acid

HCl – hydrochloric acid

H2S – hydrosulfuric acid• Root of S for acid name not “sulf” as in

Na2S (sodium sulfide)

Naming Acids Similar to Binary Acids (Rare)

If second part of compound is a polyatomic anion that does not contain oxygen (rare), use same system as for a true binary acid employing the root name for the anion

CN- – cyanide anion

HCN – hydrocyanic acid

Naming OxyacidsName is based solely on the anion

“A form of the root name of the anion” + suffix + acid

Anion suffix Acid Suffix

-ate -ic

-ite -ous

HNO3 Nitric acid NO3- = nitrate

HNO2 Nitrous acid NO2- = nitrite

Naming Molecular CompoundsFlow Chart, Fig 9-9, page 251

Naming Molecular CompoundsFlow Chart, Fig 9-9, page 251

Acidic Not Acidic

Practice

Problems 13-17 page 249 (binary covalent)Problems 18-22 page 250 (acids)Problems 27-29 page 251 (mixed)Problems 94-96(all a-d) page 273Problems 97-98(all a-d) page 273Problems 2 (a-f) page 874 (binary cov)Problem 3 page 875 (acids)


8.1 The Covalent Bond – Bond Strength



8.4 Molecular Shape


Section 8.1 The Covalent Bond

• Relate the strength of a covalent bond to its bond length, bond order, and bond dissociation energy.

• Describe how the overall energy of a reaction (i.e., whether it is an endo- or exothermic reaction) is related to the bond energies of the reactant and product molecules.

Covalent Bonding – Energy for H2

Internuclear Distance (pm)

100 200Pot

entia

l Ene

rgy

(kJ/

mol

)

-432 kJ/mol

Bond Strength & Bond Length/Order

Distance between bonding nuclei at position of max attraction = bond length

Scale of bond length: ~10-10 m =100 pmBond order: Single 1 Double 2 Triple 3


Strength of bond related to bond lengthBond dissociation energy = energy needed to break bond Triple bond > double bond > single bond

Molecule Bond Length (pm)

Dissoc. Energy kJ/mol

F2 143 159

O2 121 498

N2 110 945

Reaction Energies & Bond Energies

Chemical reaction

Bonds in reactant molecules broken

New bonds formed in product molecules

CH4 + 2O2 2H2O + CO2

Breaking C-H bonds and O=O bonds

Making O-H bonds and C=O bonds

Reaction Energies & Bond Energies

CH4(g) + 2O2(g) 2H2O(g) + CO2(g)

Total energy change determined by difference of energy of bonds broken (reactant side) and formed (product side)• Endothermic – need more energy to

break than get back in formation• Exothermic – bond formation energy

larger than energy needed to break bonds

Reaction Energies & Bond EnergiesE

ntha

lpy

-SBE (products)

-SBE (products)

SBE (reactants)

SBE (reactants)

BE = Bond energy

Chapter 8 – Covalent Bonding8.1 The Covalent Bond



8.4 Molecular Shape


Section 8.4 Molecular Shapes

• Summarize the VSEPR bonding theory, including the role of bonding and nonbonding pairs of electrons.

• Predict the shape of, and the bond angles in, a molecule using VSEPR theory.

The VSEPR model is used to determine molecular shape.

2 Simple Theories Related to Covalent Bonding

Valence Shell Electron Pair Repulsion Theory (VSEPR)• Use Lewis structures to predict shape

Valence Bond Theory• Extends Lewis bonding model to focus

on orbitals, particularly hybridized orbitals

VSEPR

Valence Shell Electron Pair Repulsion Theory - allows us to predict geometry

Lewis structures tell us how the atoms are connected to each other

Lewis structures don’t tell us anything about shape

Shape of a molecule can greatly affect its properties

Molecular Shape & Biological Sensors

For some biological systems, a response is generated or a chemical change is initiated when a molecular key fits into correspondingly shaped molecular lock• Key is typically small molecule• Lock is typically large molecule with a

shaped receptor siteOnly interacts with key of a specific shape

Lewis Structure (a) & Tetrahedral Geometry (b) for Methane (CH4)

VSEPR

Molecules take a shape that puts electron pairs as far away from each other as possible (electron pair repulsion)

Have to draw the Lewis structure to determine categories of electron pairs

• bonding• nonbonding lone pair

Lone pair take more space

Multiple bonds count as one pair

Balloon Analogy for the MutualRepulsion of Electron Groups

Two Three Four Five Six

Number of Electron Groups

VSEPR

The number of pairs determines• bond angles• underlying structure

The number and position of atoms determines

• actual molecular shape

VSEPR – Underlying Shapes

# Elec. pairs Bond Angles Shape2 180° Linear

3 120° Trigonal Planar

4 109.5° Tetrahedral

590° &120°

Trigonal Bipyramidal

6 90° Octahedral

Actual Molecular Shapes

3 3 0 trigonal planar

4 4 0 tetrahedral4 3 1 trigonal pyramidal

2 2 0 linear

3 2 1 bent

4 2 2 bent

ElectronPairs

BondingPairs

Non-Bonding

Pairs Shape


5 5 0 trigonal bipyrimidal

5 4 1 See-saw

5 3 2 T-shaped5 2 3 linear

ElectronPairs

BondingPairs

Non-Bonding

Pairs Shape


6 6 0 Octahedral

6 5 1 Square Pyramidal

6 4 2 Square Planar6 3 3 T-shaped6 2 4 linear

ElectronPairs

BondingPairs

Non-Bonding

Pairs Shape

Relative Sizes: Bonding Pairs vs Lone Pairs

CH4 NH3H2O

Molecular GeometryCan predict geometry around each atom center and build overall molecular geometry piece by piece

GlycineN

C1

C2

O1

O2


8.1 The Covalent Bond8.2 Naming Molecules8.3 Molecular Structures8.4 Molecular Shape (extension of book) Valence Bond Theory - Orbital Overlap Hybrid Orbitals

Quantum mechanical calculations8.5 Electronegativity and Polarity


• Describe the valence bond model of bonding

• Explain the similarities and differences between the Lewis and valence bond models of chemical bonds.

• Describe sigma and pi bonds and identify these bonds within molecules.

• Define hybridization.

The Valence Bond model is used to determine molecular shape via the concept of overlap of orbitals, particularly hybrid orbitals.


• Relate the type of hybridization (sp3, sp2, etc.) to the VSEPR geometry of a molecule

• Identify the specific type of hybridization that occurs within a given molecule and identify the specific orbitals (hybrid or non-hybrid) that are involved in each sigma and pi bond.

• Explain how quantum mechanics and the wave function concept can be applied to a molecule.

Valence Bond Theory

Lewis structures indicate status of electrons• Shared in bond• Lone pair

No information about orbitals involved

Valence bond theory• Bonds are formed by overlap of half-filled

atomic orbitals• Orbital geometry can give direct information

about molecular shape

Sigma Bonds

Single covalent bonds = sigma bond• Symbol Greek letter

Occurs when electron pair shared in area centered between two atoms

Atomic orbitals overlap end to end, forming a bonding orbital

• Localized region where bonding electrons will most likely be found

Sigma Bond Formation by Orbital Overlap

Two s orbitals overlap

Two s orbitals overlap

Two p orbitals overlap

H2

HF

F2

Sigma Bond Formation

Sigma Bonding – F2

px1py

2pz2 px

1py2pz

2

F—F

Area of overlap for atomic

orbitals

Pi Bond ()Formed when parallel orbitals overlap to share electrons

Shared pair occupies space above and below a line connecting atoms

Multiple bonds always have one sigma and at least one pi bond

• Double: 1 , 1 bond• Triple: 1 , 2 bonds

Sigma & Pi Bonding

Sigma () and Pi () Bonds

Hybrid Orbitals

For correct geometry of polyatomic molecules using the valence bond model, have to use concept of hybrid orbitals

• CH4 has 109.5 angles, but atomic p orbitals are at right angles to each other

Hybrid Orbitals

Hybrid orbitals – orbitals obtained when 2 or more nonequivalent orbitals combine to form an equal number of identical, degenerate orbitals

Hybridization – mixing of atomic orbitals in an atom (usually a central atom) to generate a set of hybrid orbitals

Use VSEPR logic to determine geometry of hybrid orbitals formed

Valence Orbitals on a Free Carbon Atom: 2s, 2px, 2py, and 2pz

s py

px pz

Formation of sp3 Hybrid Orbitals From Original Valence Orbitals

Hybridization

Cross Section of sp3 Orbital

Energy-Level Diagram Showing Formation of Four sp3 Orbitals

C Orbitals in CH4 molecule

Orbitals in free C atom

Hybridization

2s 2p

C

1s 1s 1s 1s

4 H

C*

Valence Bond Theory Treatment of CH4

Overlap of sp3 hybrid orbitals on C with 1s orbitals on H atoms gives 4 C-H (sp3)-1s bonds oriented 109.47° from each otherHas tetrahedral geometry predicted by VSEPR

HC

H

HH

sp3

C* (sp3)

Tetrahedral Set of Four sp3 Orbitals Forming Sigma Bonds with s Orbitals of Four Hydrogen Atoms(CH4)

Formation of sp2 Hybrid Orbitals from s, px, and py Atomic Orbitals

Hybridization

Energy-Level Diagram Showing Formation of Three sp2 Orbitals

Orbitals in sp2 hybridized C


Hybridization

Note: Inconsistent with actual bonding – 4 valence electrons populate only sp2 orbitals (Aufbau) leaving only 1 unpaired electron in sp2

An sp2 Hybridized C Atom

Formation of sp Hybrid Orbitals from s and px, Atomic Orbitals

Hybridization

Energy-Level Diagram Showing Formation of Two sp Hybrid Orbitals

Orbitals in sp hybridized C


Hybridization

Note: Inconsistent with actual bonding – 4 valence electrons should populate only sp orbitals (Aufbau) leaving no unpaired electrons

Orbitals of sp Hybridized Carbon Atom

sp3d (dsp3) Hybrid Orbitals

3s 3p

P

P*

3d

P* (sp3d) 3d

Can only occur for periods 3 & higher (need d orbitals) – example shown is for PLinked to geometry with 5 pairs (trigonal bipyramid)

3dz23pz 3py 3px3s sp3dz2

Set of dsp3 Hybrid Orbitals on a Phosphorus Atom

sp3d2 (d2sp3) Hybrid Orbitals

Can only occur for periods 3 & higher (need d orbitals)Linked to geometry with 6 pairs (octahedral)Example on next slide for S

S - Octahedral Set of d2sp3 Orbitals

Relationship among the number of effective pairs, geometry, and the

hybrid orbital set required to obtain this geometry shown on the following two

slides

Linear sp2

3

4

Trigonal sp2

planar

Tetra sp3

hedral

# Geometry Hybridization

Trigonal sp3dbipyramidal

5

6

# Geometry Hybridization

Octa sp3d2

hedral

Geometry & Hybridization - Steps1. Draw Lewis structure

2. Determine # of effective electron pairs (count double & triple bonds as one pair)

3. Determine basic geometry from number of pairs (e.g., 5 pairs = trigonal bipyramid)

4. Determine hybridization type from number of pairs (e.g., 5 pairs = sp3d)

5. Form single (sigma) bonds from hybrid orbitals; lone pairs also go in hybrid orbitals

6. Form pi bonds using unhybridized orbitals

Geometry & Hybridization - StepsFollowing slides give examples of using the steps listed on previous slide for these molecules:

1. Ammonia

2. Ethylene

3. Diatomic nitrogen

4. Acetylene

5. Carbon dioxide

6. Phosphorus pentachloride

N in Ammonia

sp3 Hybridized (4 pairs)

N in Ammonia

Trigonal pyramidal molecule with lone pair occupying hybrid orbital

Sigma & Pi Bonds Using Hybrid Orbitals - Ethylene

Three electron pairs for C sp2 hybridization & trigonal planar geometry

C 1s22s22p2 1s22(sp2)32p1

Hydrogens have 1s1 orbitals (spherical)

s Bonds in Ethylene – Top View

Sigma () bonds

Sigma and Pi Bonds in Ethylene

Pi () bond

Sigma Bonds in Ethylene

Because each C has trigonal planar geometry, entire molecule is planar

N2 Bonding

Two sp hybrid orbitals and two normal p orbitals

sp

py

pz

sp

Two pi bonds

One sigma bond

lone pair sigma lone pair

sp hybridized (2 pairs)

Sigma and Pi Bonds in Acetylene

Two electron pairs for C sp hybridization, linear geometry (triple bond = single pair)

C 1s22s22p2 1s22(sp)22p2

Hydrogens have 1s1 orbitals (spherical)


sp hybrid orbitals on C form single (sigma) bond with H and other C

Remaining two unhybridized p orbitals overlap to form two pi bonds

Sigma & Pi Bonds Using Hybrid Orbitals in CO2

Two electron pairs for C sp hybridization, linear geometry (double bond = single pair)

C 1s22s22p2 1s22(sp)22p2

Three electron pairs for O sp2 hybridization & trigonal planar geometry

O 1s22s22p4 1s22(sp2)5p1

Orbital Arrangement for an sp2 Hybridized Oxygen Atom

Sigma Bonds using Hybrid Orbitals in CO2 Molecule

Sigma () bonds

Sigma and Pi Bonds Using Hybrid Orbitals in CO2

Sigma Bonds Using Hybrid Orbitals in PCl5

Five electron pairs for P sp3d hybridization & trigonal bipyramidal geometry

P [Ne]3s23p3 [Ne](sp3d)5

Four electron pairs for Cl sp3 hybridization & tetrahedral geometry

Cl [Ne]3s23p5 [Ne]3(sp3)7

Set of dsp3 Hybrid Orbitals on a Phosphorus Atom

Structure of PCI5 and Orbitals Used to Form Sigma Bonds

Sigma () bond

Lone pairs on Cl in sp3 orbitals

Geometry & HybridizationSupply for each indicated atom in structure

# of sigma & pi bonds in molecule?

HC

HC

C O

O

C

H

H

H

C

N

3 pairs

Trigonal planar

sp2

3 pairs

Trigonal planar

sp2

2 pairs

Linear

sp

4 pairs (2 lone)

Bent

sp3

4 pairs

Tetrahedral

sp3

12 , 4

Practice (Shape, Angles, Hybridization)

Problems 56 – 60, 65 - 67 page 264

Problems 108,110 - 112 page 275

Problem 8 page 980

Quantum Mechanics & Molecules

Quantum Mechanics & Molecules

Y (wave function) exists for entire molecule and can be obtained from solution to Schrodinger wave equation written for the molecule

Y2 - Square of Y gives probability of finding electron at particular position around molecule – defines what is called a molecular orbital (MO)

Quantum Mechanics & MoleculesUsing certain types of approximations and today’s computers, wave functions for molecules (not individual atoms) can be obtained and molecular properties calculated from this information

Energy, absorption spectrum, dipole moment, etc

Molecular orbital theory is most advanced way of describing covalent bonding





8.4 Molecular Shape


Section 8.5 Electronegativity and Polarity

• Describe how electronegativity is used to determine bond type and characterize bonds between given pairs of atoms as being polar or nonpolar.

• Compare and contrast polar and nonpolar covalent bonds and polar and nonpolar molecules.

• Describe the term “dipole moment” and relate it to the terms polar and nonpolar.

A chemical bond’s character is related to each atom’s attraction for the electrons in the bond.

Section 8.5 Electronegativity and Polarity

• Identify molecules as being polar or nonpolar.

• Describe how polarity affects the solubility of one substance in another substance.

• Describe how polarity can give rise to intermolecular forces.

Polar Covalent Bonds

Polarity of bond determined by electronegativity difference

Difference = 0 Nonpolar

Difference > 0 Polar

Very large differences • No longer covalent compound

EN Difference & Bond Character

0 1.0 2.0 3.0

Electronegativity Difference

Ionic Bonds

Covalent Bonds

% Io

nic

Cha

ract

er

25

50

75

100

Relationship Between EN Difference and Bond Type

Relationship Between EN Difference and Bond Type

EN=0

Nonpolar Covalent

EN=medium

Polar Covalent

EN=large

Ionic

Scalars & VectorsScalar

• Completely specified by magnitude and units

Vector• Has magnitude, direction, and units

v = 3.5 m/s (scalar)

v = 3.5 m/s to northeast (vector)

Trigonometric Functions

Pythagorean Theorem

Dipole MomentTwo equal and opposite charges +Q and -Q separated by a distance l have a dipole moment p:

(vector points from –Q to +Q) Q lp = p =

r

+Q -Q

l

pr

Polarity and Dipole Moment

Dipole moment is a vector pointing from center of - charge to center of + chargeMagnitude proportional to size of charges and to separation distanceAll polar covalent bonds have a dipole moment

+ _Dipole

Polarity and Dipole Moment

Units of p are Debye units (D)% ionic character of bond determined by size of measured dipole moment relative to value calculated from using full (ionic) charges as Q

Q lp = p = +Q -Q

l

pr

Dipole Moments of Gas Phase Molecules

Dipole Moment

moment) dipole(net 21 ppp += r rr

Dipole moments from bonds add as vectors to give dipole moment of molecule

Molecular Polarity – Linear Molecule

O=C bond polar; bonding electrons pulled equally toward both O ends of molecule

Net result is nonpolar molecule (dipole moments of bonds cancel each other) [note: red arrows are opposite dipole direction]

HO bond polar

Both sets of bonding electrons pulled toward O end; net result is polar molecule (y components of bond dipole moment add, x components cancel)

[note: red arrows are opposite dipole direction]

Molecular Polarity – Bent Molecule

Polar Molecules

Molecule can have polar bonds but be a nonpolar molecule

-+

+-

-

-

+

-

Polar Nonpolar

Polar Bonds in Nonpolar Molecules

In symmetric molecules, vector addition of bond dipole moments results in zero dipole moment for the molecule

All molecules having basic VSEPR shapes & equal bonds are nonpolar

Linear+

Trigonal Planar +

Practice (Polar Bonds & Polar Molecules)

Problems 74 – 77 page 270

Problems 117 – 123 page 275

Problem 9 page 980

Polarity Effects

Polarity of molecule determines solubility characteristics – “like dissolves like”

Oil (nonpolar) and water

(polar) don’t mix

Dipole in an Electric FieldThe + and – charges in an electric dipole are pulled in opposite directions in an electric field, producing a net torque on the dipole, and orienting it.

+

-δ + δ

F HFieldOff

FieldOn

Dipole in Electric Field – HF Molecule

Polar Molecule & Electric FieldPolar molecules affected by electric field in an EM wave

Oscillating field twists water molecule and energy transferred (heats up)

Basis for microwave oven operation

Properties of Covalent Compounds

Bonding types affect properties

Many properties controlled by intermolecular forces

• Forces between molecules• Also known as van der Waals forces

Intermolecular forces are weaker than chemical bonds

[Note: intermolecular forces treated in more depth in section 12.2 – Forces of Attraction]

Intermolecular Forces

Forces between nonpolar molecules relatively weak

• Tend to be gases or volatile liquids• O2, N2, small hydrocarbons

Forces between polar molecules are stronger due to dipole-dipole forces

• Hydrogen bonding a particular strong version - H and F, O, or N

Hydrogen Bonding – Water Pentamer

Hydrogen Bonds

Hydrogen Bonding in Nylon

Hydrogen bonding helps make nylon strong

End of Chapter 8

chapter 8 – covalent bonding 8.1the covalent bond 8.2 naming molecules 8.3molecular structures...

Documents

covalent bonding energy

covalent character

covalent bonding forces

slide bond

ionic vs covalent bonding

ionic bonding

ionic bonds covalent

bonding pair of electrons