Chemical Bonding

Properties depend on:

• Type of bonds between atoms– intramolecular

• Shape of the molecules

• Interactions between molecules– intermolecular

Bond Types

• Metallic (metal – metal)

• Ionic (metal – nonmetal)

• Covalent (nonmetal- nonmetal)

Key difference between the bonds is in the nature of the “positive/negative” attraction

Metallic Bonding

• Metals tend to form cations (lose electrons)

• Cations and “free electrons” are said to form a “cloud or sea of electrons” in which all atoms are able to share electrons

The electrons are not bound to an individual atom.

This atomic sharing of electrons is the “glue” that holds together the metal atoms.

Electronegativity

• Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons.

What happens if B is a lot more electronegative than A?

• In this case, the electron pair is dragged over to B's end of the bond.

• Ions have been formed.

Ionic Bonding

• Typically formed between metals and nonmetals (High difference in electronegativities)

• Electrons are transferred from one atom to another (from metal to nonmetal) resulting in the formation of positive and negative ions.

• The electrostatic attractions between the positive and negative ions hold the compound together.

Ionic BondingIonic Bonding

Consider the reaction between sodium and chlorine:Na(s) + ½Cl2(g) NaCl(s) Hºf = -410.9 kJ

Ionic BondingIonic BondingNa(s) + ½Cl2(g) NaCl(s) Hºf = -410.9 kJ

- -

-

--

+

++

+

Na(s) + 1/2 Cl2(g)

Na(g) + 1/2 Cl2(g)

Na(g) + Cl(g)

0

+100

+200

+300

+400

+500

+600

+700

+800

-400

-300

-200

-100

kJ

Na(s) + 1/2 Cl2(g)

Na(g) + 1/2 Cl2(g)

Na(g) + Cl(g)

0

+100

+200

+300

+400

+500

+600

+700

+800

-400

-300

-200

-100

kJ

Na(s) + 1/2 Cl2(g)

Na(g) + 1/2 Cl2(g)

Na(g) + Cl(g)

0

+100

+200

+300

+400

+500

+600

+700

+800

-400

-300

-200

-100

kJ

Na+(g) + Cl(g)

Na+(g) + Cl-(g)

NaCl(s)

Lattice enthalpy for sodium chloride

H = -786 kJmol-1

Calculation of enthalpy of formation

• The loss of an electron from an element– Ionization energy– Na(g)Na+(g) + 1e- ∆H = 496 kj/mol

• The gain of an electron by a nonmetal– Electron affinity– Cl(g) + 1e-Cl-(g) ∆H = -349 kj/mol

• Attraction of cation and anion– Lattice energy– Na+(g) + Cl-(g)NaCl(s) ∆H = -788 kj/mol

Ionic BondingIonic Bonding

Energetics of Ionic Bond Formation• Lattice energy: the energy required to completely

separate an ionic solid into its gaseous ions.

is a proportionality constant (depends on solid structure & e- configs of ions),

Q1 and Q2 are the charges on the ions

d is the distance between ions.

dQQ

El21

El

Charge

Distance

due to charge increase

due to size increase

What happens if two atoms of equal electronegativity bond together?

• If the atoms are equally electronegative, the bonding electrons are evenly shared

What happens if B is slightly more electronegative than A?

• B will attract the electron pair rather more than A does.

• Uneven sharing results in one side of the bond being more negative than the other (polarity)

Covalent Bonding

• Typically between two or more nonmetals

• No, or low, difference between electronegativities

• Positive nucleus is attracted to negative electron cloud of other atom

Nonpolar Covalent

• Both atoms have the same electronegativity.

• Usually two identical atoms, for example, H2 or Cl2 molecules

• This sort of bond could be thought of as being a "pure" covalent bond - where the electrons are shared evenly between the two atoms.

Polar Covalent

• A covalent bond in which there is a separation of charge between one end and the other - in other words in which one end is slightly positive and the other slightly negative.

Covalent BondingCovalent Bonding

Lewis Structures• Covalent bonds can be represented by the Lewis symbols

of the elements:

• In Lewis structures, each pair of electrons in a bond is represented by a single line:

Cl + Cl Cl Cl

Cl Cl H FH O

H

H N H

HCH

H

H

H

Covalent BondingCovalent Bonding

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 decrease as we move from single through double to triple bonds.

H H O O N N

Covalent BondingCovalent Bonding

Multiple Bonds

• Generally, bond distances decrease as we move from single through double to triple bonds.

• Generally, bond energies increase as we move from single through double to triple bonds.

H H O O N N

In order to determine# bonds, we need to learnhow to draw molecules!

Bond Energy and EnthalpyBond Energy and Enthalpy

The enthalpy of a reaction depends on the strength of the bonds of the molecules involved in the reaction.

-A reaction with tightly bound reactants will require a higher input of energy to make the reaction proceed than one with loosely bound reactants.

-Likewise, the amount of energy required to form the bonds of the products affects the overall enthalpy.

Bond Energy and EnthalpyBond Energy and Enthalpy

The enthalpy of a reaction is given by:    ∆ H = Σbond energies of reactants -

Σbond energies of products

Note that bond energies are always positive quantities.

Bond Energy and EnthalpyBond Energy and Enthalpy

Example: Using bond energies, calculate the enthalpy of the following

reaction.

    C3H8(g ) + 5O2(g ) 3CO2(g ) + 4H2O(l )

H-C 414 J/mol C-C 347 J/mol H-O 464 J/mol   

C=O 730 J/mol C-O 351 J/mol   

O=O 502 J/mol O-O 142 J/mol

Bond Energy and EnthalpyBond Energy and Enthalpy

Σ bond energies of reactants = 8(414) + 2(347) + 5(502) = 6516 kJ/mol

Σ bond energies of products = 6(730) + 8(464) = 8092 kJ/mol

The enthalpy change is therefore -1576 kJ/mol.

Drawing Lewis StructuresDrawing Lewis Structures

1. Add valence electrons.

2. Write symbols for the atoms

3. Show initial bondings.

4. Try to complete the octets (8e-)

5. If there are not enough electrons try multiple bonds.

Drawing Lewis StructuresDrawing Lewis Structures

1. Add valence electrons.

2. Draw skeleton structure -put atom with lowest electronegativity

in middle (except hydrogen)

4. Show initial bondings.

5. Try to complete the octets (8e-)

6. If there are not enough electrons try multiple bonds.

NF3

5 + (3x7) = 26e-

.N .. ..

... .F. ..... .F. ..

... .F. ..

N..

... .F. ....F. ..

... F.. .

Polyatomic IonsPolyatomic Ions

Same procedure except:

•Take charge into account

-add electrons for negative charge

-subtract electrons for positive charge

•Brackets around structure with charge shown in upper right

[ structure ]charge

Resonance StructuresResonance Structures

The Lewis structure of ozone (O3)

Resonance StructuresResonance Structures

However... known facts about the structure of ozone:

The bond lengths between the central oxygen and

the other two oxygens are identical:

Resonance StructuresResonance Structures

We would expect that if one bond was a double bond that it should be shorter than the other (single) bond

Since all the atoms are identical (oxygens) which atom is chosen for the double bond?

Resonance StructuresResonance Structures

-These Lewis structures are equivalent except for the placement of the electrons

-Equivalent Lewis structures are called resonance structures, or resonance forms

-The correct way to describe ozone as a Lewis structure would be:

Resonance StructuresResonance Structures

The important points to remember about resonance forms are:

The molecule is not rapidly oscillating between different discrete forms

Resonance StructuresResonance Structures

There is only one form of the ozone molecule, and the bond lengths between the oxygens are intermediate between characteristic single and double bond lengths between a pair of oxygens

We draw two Lewis structures (in this case) because a single structure is insufficient to describe the real structure

Exceptions to Octet RuleExceptions to Octet Rule

Less than an octet:

•”Wimpy” atoms bonding with highly electronegative atoms

-typical of B, Be, Al

Exceptions to Octet RuleExceptions to Octet Rule

Greater than an octet:

(central atom must have a d sublevel)

-more than 4 atoms around central atom (PCl5)

-extra pairs of valence electrons (I3-)

Formal ChargeFormal Charge

Sometimes when writing a Lewis structure you come across two different ways to write the molecule, both which look fine.

In this case, you should use formal charge to decide which structure is correct for the molecule.

Formal ChargeFormal Charge

The formal charge is the difference in the number of valence electrons in the atom and the number of valence electrons in the Lewis structure.

Formal ChargeFormal Charge

The equation for the formal charge of any atom in a Lewis structure is

Cf = Ev - (Eu + 1/2Ep)

Formal ChargeFormal Charge

where Cf is the formal charge

Ev is the number of valence electrons in the bare atom Eu is the number of electrons in lone pairs on the atom

in the Lewis structure

Ep is the number of electrons in bonded pairs on the atom in the Lewis structure

Formal ChargeFormal Charge

To decide if a given structure is correct, check the formal charge on some atoms in all possible structures. In general the most likely Lewis structure has: -all formal charges as close to zero as possible

-negative formal charges on electronegative atoms like halogens or oxygen.

Formal ChargeFormal Charge

For example, consider the methanol molecule CH3OH. This can be written two different ways:

In both cases the octet rule is satisfied for all of the atoms in the structure. Which is correct?

Formal ChargeFormal Charge

For the leftmost structure

The carbon has four bonds, each worth 2 electrons, for a total of eight. It has no lone pairs. Thus, Cf = 4 - (0 + 1/2*8) =0

The oxygen has two bonds, each worth 2 electrons, for a total of four. It has two lone pairs. Thus, Cf = 6 - (4 + 1/2*4) =0

Formal ChargeFormal Charge

For the rightmost structure

The carbon has three bonds, each worth 2 electrons, for a total of six. It has one lone pair. Thus, Cf = 4 - (2 + 1/2*6) = -1

The oxygen has three bonds, each worth 2 electrons, for a total of six. It has one lone pair. Thus, Cf = 6 - (2 + 1/2*6) = +1

Formal ChargeFormal Charge

The leftmost structure has the formal charges closer to zero, and thus is probably the correct structure.

Molecular Geometry…

• …is simply the shape of a molecule.

• Molecular geometry is found using the Lewis structure, but the Lewis structure itself does NOT necessarily represent the molecule’s shape.

A water molecule is angular or

bent.

VSEPR

• Valence-Shell Electron-Pair Repulsion (VSEPR) is a simple method for determining geometry

• Basis: pairs of valence electrons in bonded atoms repel one another.

• These mutual repulsions push electron pairs as far from one another as possible.

B A B

B

B B

B

A

Electron Geometries

• An electron group is any collection of valence electrons, localized in a region around a central atom, that repels other groups of valence electrons.

• The mutual repulsions among electron groups lead to an orientation of the groups that are called electron geometry.

• These geometries are based on the number of electron groups:

2 linear

3 trigonal planar

4 tetrahedral

5 trigonal bipyramidal

6 octahedral

A Balloon Analogy

Each electron group may be:

-an unshared pair of electrons, or

-a bond (single, double, triple bonds are each counted as one electron group).

VSEPR Notation

• In the VSEPR notation used to describe molecular geometries, the central atom in a structure is denoted as A, terminal atoms as X, and the lone pairs of electrons as E.

• Example: ClF3 is designated AX3E2. It has three groups (atoms in this case) around the Cl atom, and two lone pairs of electrons on the Cl (draw the Lewis structure to see).

Polar MoleculesAnd Dipole Moments

• A polar bond is a bond with separate centers of positive and negative charge.

• A molecule with separate centers of positive and negative charge is a polar molecule.

• The dipole moment () of a molecule is the product of the magnitude of the charge () and the distance (d) that separates the centers of positive and negative charge.

= d• A unit of dipole moment is the debye (D).• One debye (D) is equal to 3.34 x 10-30 C m.

Polar Molecules In An Electric Field

An electric field causes polar molecules to “line up” but has no

effect on nonpolar molecules.

Bond Dipoles AndMolecular Dipoles

• A polar covalent bond has a bond dipole; a separation of positive and negative charge centers in an individual bond.

• Bond dipoles have both a magnitude and a direction (they are vector quantities).

• A molecule can have polar bonds, but may be a nonpolar molecule – IF the bond dipoles cancel.

Bond Dipoles AndMolecular Dipoles

• CO2 has polar bonds, but is a linear molecule; the bond dipoles cancel and it has no net dipole moment ( = 0 D)

• The water molecule has polar bonds also, but is an angular molecule.

• The bond dipoles do not cancel, so water is a polar molecule.

Net dipole

No net dipole

Molecular ShapesAnd Dipole Moments

• Molecular polarity can be predicted based on the following three-step approach:1. Use electronegativity values to predict bond dipoles.2. Use the VSEPR method to predict the molecular

shape.3. From the molecular shape, determine whether bond

dipoles cancel to give a nonpolar molecule or combine to produce a resultant dipole moment for the molecule.

DETERMINING MOLECULAR POLARITY

• 1. Molecules are non-polar or polar• 2. Non-polar molecules have an even (symmetrical) distribution

of charge (+ or – )– If all atoms are the same in a 2-atom molecule (non-polar bonds)

(H2 , N2 , Br2 )– If there are no lone pairs on the central atom and the attached

atoms are all the same

(CO2 , BCl3 , CH4) • 3. Polar molecules have an uneven (assymetrical) distribution

of charge. The molecule has a dipole (+ side and a – side, like a bar magnet)– If the outer atoms are different from each other

(HCl , H2CO , CH3F)OR

– If there are lone pairs on the central atom

(H2O , NH3 , SO2 , O3)

Localized Electron Model

• A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms– Molecular Structure (Lewis Structures/VSEPR)– Bonds within the molecule

• Valence Bond Theory• Hybridization of orbitals

Atomic Orbital Overlap

• Valence Bond (VB) theory states that a covalent bond is formed when atomic orbitals (AOs) overlap.

• In the overlap region, electrons with opposing spins produce a high electron charge density.

Overlap region

between nuclei has

high electron density

• In general, the more extensive the overlap between two orbitals, the stronger is the bond between two atoms.

Bonding In H2S

Hydrogen atoms’ s-orbitals can overlap with the two half-filled

p- orbitals on sulfur.

The measured bond angle in H2S is 92°;

good agreement

Points of VB Theory

• Most of the electrons in a molecule remain in the same orbital locations that they occupied in the separated atoms.

• Bonding electrons are localized in the region of AO overlap.

• For AOs with directional lobes (such as p-orbitals), maximum overlap occurs when the AOs overlap end to end.

• VB theory is not without its problems…

Hybridization Of Atomic Orbitals• VB theory: carbon should have only two

bonds, and they should be about 90° apart.

• We can hybridize the four orbitals; mathematically combine the wave functions for the 2s orbital and the three 2p orbitals on carbon.

• The four AOs combine to form four new hybrid AOs.• The four hybrid AOs are equivalent, and each has a single

electron (Hund’s rule).

• Reality: carbon has four bonds, which (singly bonded) are about 109° apart.

Four equivalent hybrid orbitalscan now form four bonds

sp3 Hybridization

• Hybridizing an s-orbital with three p-orbitals gives rise to four hybrid orbitals called sp3 orbitals.

• The number of hybrid orbitals is equal to the number of atomic orbitals combined.

• The four hybrid orbitals, being equivalent, are about 109° apart.

The sp3 Hybridization Scheme

Four AOs…

…form four new hybrid AOs.

Methane and Ammonia

In methane, each hybrid orbital is a

bonding orbital

In ammonia, one of the hybrid orbitals (top) contains

the lone pair that is on the nitrogen atom

Four sp3 hybrid orbitals: tetrahedralFour electron groups: tetrahedral

Coincidence? Hardly…

sp2 Hybridization

• Three sp2 hybrid orbitals are formed from an s-orbital and two p-orbitals.

• The empty p-orbital remains unhybridized. It may be used in a multiple bond.

• The sp2 hybrid orbitals are in a plane, 120o apart.• This distribution gives a trigonal planar molecular

geometry, as predicted by VSEPR.

The sp2 Hybridization Scheme in Boron

A 2p orbital remains unhybridized.

Three AOs combine to form…

…three hybrid AOs

sp Hybridization

• Two sp hybrid orbitals are formed from an s-orbital and a p-orbital.

• Two empty p-orbitals remains unhybridized; the p-orbitals may be used in a multiple bond.

• The sp hybrid orbitals are 180o apart.• The geometry around the hybridized atom is linear,

as predicted by VSEPR.

sp Hybridization in Be

Two unused p-orbitals

Hybrid Orbitals Involvingd Subshells

• This hybridization allows for expanded valence shell compounds.

• By hybridizing one s, three p, and one d-orbital, we get five sp3d hybrid orbitals.

• This hybridization scheme gives trigonal bipyramidal electron-group geometry.

• By hybridizing one s, three p, and two d-orbitals, we get five sp3d2 hybrid orbitals.

• This hybridization scheme gives octahedral geometry.

The sp3d and sp3d 2 Hybrid Orbitals

sp3d sp3d 2

Predicting Hybridization Schemes

In the absence of experimental evidence, probable hybridization schemes can be predicted:

1. Write a plausible Lewis structure for the molecule or ion.

2. Use the VSEPR method to predict the electron-group geometry of the central atom.

3. Select the hybridization scheme that corresponds to the VSEPR prediction.

4. Describe the orbital overlap and molecular geometry.

Hybrid Orbitals and TheirGeometric Orientations

Hybrid Orbitals AndMultiple Covalent Bonds

• Covalent bonds formed by the end-to-end overlap of orbitals, regardless of orbital type, are called sigma () bonds.

• All single bonds are sigma bonds.• A bond formed by parallel, or side-by-side, orbital

overlap is called a pi () bond.• A double bond is made up of one sigma bond and one

pi bond.• A triple bond is made up of one sigma bond and two pi

bonds.

VB Theory for Ethylene, C2H4

σ-bond: overlap of s-orbital of hydrogen

and sp2 hybrid orbital.

σ-bond: sp2 - sp2 overlap

π-bond has two lobes (above and below plane), but is one

bond. Side overlap of 2p–2p.

Summary of VB theory of Ethylene

VB Theory: Acetylene

σ-bond: sp - sp overlap

σ-bond: s - sp overlap

Two π-bonds (above and below, and front and back) from 2p–

2p overlap…

…form a cylinder of π-electron density around the two carbon atoms.