Chapter 8: Covalent Compounds

Polarity

A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.

Polarity of MoleculesPolarity of MoleculesDipole: def: two charges, equal in magnitude and opposite in sign, are separated by a distance

Polarity of Polyatomic MoleculesPolarity of Polyatomic Molecules•Each bond can be polar.•The orientation of these polar bonds determines whether the molecule is polar overall.•It is possible for a molecule with polar bonds to be either polar or non-polar.

Polarity of MoleculesPolarity of MoleculesDipole Moments of Polyatomic MoleculesDipole Moments of Polyatomic MoleculesExample: in CO2, each C-O dipole is canceled because the molecule is linear. In H2O, the H-O dipoles do not cancel because the molecule is bent.

Polarity of MoleculesPolarity of MoleculesDipole Moments of Polyatomic MoleculesDipole Moments of Polyatomic Molecules

Models Models are attempts to explain how nature operates

on the microscopic level based on experiences in the macroscopic world.

Properties of Models: A model does not equal reality.

Models are oversimplifications, and are therefore often wrong.

Models become more complicated as they age.

We must understand the underlying assumptions in a model so that we don’t misuse it.

Lewis Structure Shows how valence electrons are

arranged among atoms in a molecule.

Valence electrons are found in the outermost energy level of an atom. They are involved in bonding.

Reflects the central idea that the stability of a compound relates to noble gas electron configuration.

Lewis Dot Structure Rules:

Treat ions separately (e.g. NH4Cl) Count only valence electrons Assemble bonding framework Fill up non-bonding electrons on outer

atoms Fill up non-bonding electrons on inner

atoms Calculate Formal Charge Minimize Formal Charge

To Complete a Lewis Structure: Must be able to recognize polyatomic ions Must be able to identify valence electrons

Must be able to construct Bond framework

Hints on Lewis Dot Structures1. Octet “rule” is the most useful guideline.2. Carbon always forms 4 bonds.3. Other 2nd row elements (N, O, F) observe

the octet rule.4. Hydrogen forms one bond with other atoms

to complete its octet with 2 electrons.5. 2nd row elements B and Be often have

fewer than 8 electrons around themselves - they are very reactive.

6. 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals.

Notes About the Octet Rule

7. When multiple bonds are forming, they are usually between C, N, O or S.

8. Nonmetals can form single, double, and triple bonds, but not quadruple bonds.

9. Always account for single bonds and lone pairs before forming multiple bonds.

10. Look for resonance structures. 11. When writing Lewis structures, satisfy

octets first, then place electrons around central atoms having available d orbitals.

PCl3

Bonding Pairs

Lone Pairs(a.k.a. nonbonding electrons)

5+(3*7)=26 e-

Try Some Examples:

C2H5OH

CH3CH2NH2

Cl2CO

Ozone (O3)

NO2

Formal ChargeDifference between the # of valence electrons in the free atom and the # of electrons assigned to that atom in the Lewis structure.

FC = formal charge; G.N. = Group Number#BE = bonding electrons; #LPE = lone pair electrons

If Step 4 leads to a positive formal charge on an inner atom beyond the second row, shift electrons to make double or triple bonds to

minimize formal charge, even if this gives an inner atom with more than an octet of electrons.

LPEBE ##2

1- G.N. FC

Formal Charge

Not as good Better

CO O(-1) (0) (+1)

CO O(0) (0) (0)

Multiple Bonds• It is possible for more than one pair of electrons to

be shared between two atoms (multiple bonds):• One shared pair of electrons = single bond (e.g. H2);

• Two shared pairs of electrons = double bond (e.g. O2);

• Three shared pairs of electrons = triple bond (e.g. N2).

• Generally, bond distances shorten with multiple bonding.

H H O O N N

Covalent BondingCovalent Bonding

Octet in each case

Resonance

Occurs when more than one valid Lewis structure can be written for a particular molecule.

Odd Number of Electrons…

NO Number of valence electrons = 11

N O N O

NO2Number of valence electrons = 17

Resonance occurs when more than one valid Lewis structure can be written for a particular molecule (i.e. rearrange electrons)

Molecules and atoms which are neutral (contain no formal charge) and with an

unpaired electron are called Radicals

N OO N OO N OO

Beyond the Octet

Elements in the 3rd period or higher can have more than an octet if needed.

Atoms of these elements have valence d orbitals, which allow them to accommodate more than eight electrons.

More than an Octet…

PCl5

Elements from the 3rd period and beyond, have ns, np and unfilled nd orbitals which can be used in bonding

P : (Ne) 3s2 3p3 3d0

Number of valence electrons = 5 + (5 x 7) = 40 P

Cl

ClCl

ClCl

SF4S : (Ne) 3s2 3p4 3d0

Number of valence electrons = 6 + (4 x 7) = 34

SF

F

F

FThe Larger the central atom, the more atoms you can bond to it – usually small atoms such as F, Cl and O allow central atoms such as P and S to expand their valency.

Less than an Octet…

BCl3Group 13 atom only has six electrons around itB

Cl

Cl Cl

Molecular ShapesMolecular Shapes• Lewis structures give atomic connectivity:

they tell us which atoms are physically connected together. They do not tell us the shape.

• The shape of a molecule is determined by its bond angles.

• Consider CCl4: experimentally we find all Cl-C-Cl bond angles are 109.5.Therefore, the molecule cannot be planar.All Cl atoms are located at the vertices of a

tetrahedron with the C at its center.

Molecular Shape of CClMolecular Shape of CCl44

VSEPR TheoryVSEPR Theory

In order to predict molecular shape, we assume the valence electrons repel each other. Therefore, the molecule adopts whichever 3D geometry minimized this repulsion.

We call this process Valence Shell Electron Pair Repulsion (VSEPR) theory.

Why is VSEPR Theory Important? Gives a specific shape due to the

number of bonded and non-bonded electron pairs in a molecule

Tells us the actual 3-D structure of a molecule

In bonding, electron pairs want to be as far away from each other as possible.

VSEPR and Resulting GeometriesVSEPR and Resulting Geometries

How does VSEPR THEORY work?

We can use VSEPR theory using 4 steps.

1. Draw the Lewis Structure for the molecule.Example: SiF4 F

F-Si-F

F

How does VSEPR THEORY work?

We can use VSEPR theory using 4 steps

1. Draw the Lewis Structure for the molecule.

2. Tally the number of bonding pairs and lone (non-bonding) pairs on the center atom.

F

F-Si-F

F

Bonding pairs: 4

Lone pairs on central atom: 0

How does VSEPR THEORY work?

We can use VSEPR theory using 4 steps

1. Draw the Lewis Structure for the molecule

2. Tally the number of bonding pairs and lone pairs on the center atom.

3. Arrange the rest of the atoms so that they are as far away from each other as possible. Si

F

FF F

How does VSEPR THEORY work?We can use VSEPR theory using 4 steps1. Draw the Lewis Structure for the

molecule2. Tally the number of bonding pairs

and lone pairs on the center atom.3. Arrange the rest of the atoms so that

they are as far away from each other as possible

4. Give the type of geometry the molecule has:

Tetrahedral

Another Example:

To determine the electron pair geometry:1) draw the Lewis structure;2) count the total number of electron pairs around the central atom.3) arrange the electron pairs in one of the geometries to minimize e-e repulsion.4) multiple bonds count as one bonding pair for VSEPR

The VSEPR ModelThe VSEPR Model

Predicting Molecular GeometriesPredicting Molecular Geometries

The VSEPR ModelThe VSEPR Model

Predicting Molecular GeometriesPredicting Molecular Geometries

The VSEPR ModelThe VSEPR Model

Difference between geometry and shapeDifference between geometry and shapeGeometry:We determine the geometry only looking at electrons.All the atoms that obey the octet rule have the same tetrahedral-like geometry.

Shape:We name the shape by the positions of atoms.We ignore lone pairs in the shape.

The VSEPR ModelThe VSEPR Model

Predicting ShapePredicting Shape

Shape

The VSEPR ModelThe VSEPR Model

Predicting ShapePredicting Shape

Shape

The VSEPR ModelThe VSEPR ModelThe Effect of Nonbonding Electrons and The Effect of Nonbonding Electrons and Multiple Bonds on Bond AnglesMultiple Bonds on Bond AnglesBy experiment, the H-X-H bond angle decreases on moving from C to N to O:

Since electrons in a bond are attracted by two nuclei, they do not repel as much as lone pairs.Therefore, the bond angle decreases as the number of lone pairs increase.

OHH

104.5O107O

NHH

HC

H

HHH109.5O

The VSEPR ModelThe VSEPR Model

The Effect of Nonbonding Electrons and The Effect of Nonbonding Electrons and Multiple Bonds on Bond AnglesMultiple Bonds on Bond AnglesSimilarly, electrons in multiple bonds repel more than electrons in single bonds.

C OCl

Cl111.4o

124.3o

The VSEPR ModelThe VSEPR Model

Molecules with Expanded Valence ShellsMolecules with Expanded Valence ShellsAtoms that have expanded octets have AB5 (trigonal bipyramidal) or AB6 (octahedral) electron pair geometries.

Examples:PF5 trigonal bipyramidal SCl6 octahedral

The VSEPR ModelThe VSEPR Model

Molecules with Expanded Valence ShellsMolecules with Expanded Valence Shells

The VSEPR ModelThe VSEPR Model

Molecules with Expanded Valence ShellsMolecules with Expanded Valence Shells

The VSEPR ModelThe VSEPR ModelMolecules with More than One Central AtomMolecules with More than One Central AtomIn acetic acid, CH3COOH, there are three central atoms.We assign the geometry about each central atom separately.

Hybrid Orbitals In bonding, s and p orbitals are

used in bonding. It is easy to tell which ones are used by looking at our molecule.

For example, CH4. Looking again at the Lewis structure, we see that there are 4 bonds. We call this sp3 hybridized.

Hybrid Orbitals

Regions of electron density-EACH BOND AND LONE PAIR OF ELECTRONS ON THE CENTRAL ATOM IS KNOWN AS A REGION OF ELECTRON DENSITY.

2 regions of electron density-sp hybridized

3 regions of electron density-sp2 hybridized

4 regions of electron density-sp3 hybridized

Hybridization

spsp Hybrid Orbitals Hybrid OrbitalsThe two lobes of an sp hybrid orbital are 180 apart.

Hybrid Orbitals

spsp22 Hybrid Orbitals Hybrid OrbitalsImportant: when we mix n atomic orbitals we must get n hybrid orbitals.sp2 hybrid orbitals are formed with one s and two p orbitals. (Therefore, there is one unhybridized p orbital remaining.)The large lobes of sp2 hybrids lie in a trigonal plane.All molecules with trigonal planar electron pair geometries have sp2 orbitals on the central atom.

Hybridization

Hybridization

spsp33 Hybrid Orbitals Hybrid Orbitalssp3 Hybrid orbitals are formed from one s and three p orbitals. Therefore, there are four large lobes.Each lobe points towards the vertex of a tetrahedron.The angle between the large lobes is 109.5All molecules with tetrahedral electron pair geometries are sp3 hybridized.

Hybridization

Hybrid OrbitalsHybrid Orbitals

Hybrid OrbitalsHybrid Orbitals

Hybrid OrbitalsHybrid OrbitalsSummarySummaryTo assign hybridization:

1. Draw a Lewis structure.2. Assign the geometry using VSEPR theory.3. Use the geometry to determine the hybridization.4. Name the shape by the positions of the atoms.

Hybridization and Multiple Bonds

•Multiple bonds overlap differently and are called -bonds and -bonds

•All single bonds are •Double bonds contain 1 and 1 bond

•Triple bonds contain 1 and 2 bonds

Bond Energy

Covalent Bonding & Orbital Overlap

•As two nuclei approach each other their atomic orbitals overlap.•As the amount of overlap increases, the energy of the interaction decreases.•At some distance the minimum energy is reached.•The minimum energy corresponds to the bonding distance (or bond length).

Covalent Bonding & Orbital Overlap

•As the two atoms get closer, their nuclei begin to repel and the energy increases.

•At the bonding distance, the attractive forces between nuclei and electrons just balance the repulsive forces (nucleus-nucleus, electron-electron).

Bond Energies

Bond breaking requires energy (endothermic).

Bond formation releases energy (exothermic).

H = D(bonds broken) D(bonds formed)

energy required energy released

Bond Energies