Chemical Bonding

Chemical Bonding

is all about the electrons. in most of our discussions we will

concentrate on the outer ‘s’ and ‘p’ electrons. These are called the valence electrons.

since there are 4 of these orbitals in any quantum it requires 8 electrons (an octet) to fill them.

the tendency of atoms to try to fill out the outer ‘s’ and ‘p’ shells is the octet rule.

Chemical Bonding

Elements will combine in order to fill their valence shell with electrons, like the noble gases. This can happen in one of two ways:

To share electrons -

The outer orbitals of 2 atoms overlap so that each atom is in the vicinity of a full set of valence electrons.

This type of bonding is called covalent bonding.

Chemical bonding

Atoms gain or lose electrons to arrive at a full set of valence electrons.

When atoms gain or lose electrons they become ions.

Ions are attracted to ions of opposite charge and repelled by ions of the same charge.

This type of bonding is called ionic bonding

How Many Bonds ???

• the number of bonds made by atoms in the ‘s’ and ‘p’ blocks of the Periodic Table is determined by how many electrons they are away from an octet:

Group Valence Electrons

Covalent Bonds

IonicBonds

1 1 1 1 (1+)

2 2 21 2 (2+)

13 3 31 3 (3+)

14 4 41 -

15 5 3 (52) 3 (3-)

16 6 2 (62) 2 (2-)

17 7 1 (72) 1 (1-)

18 8 0 (82) 01electron promotion 2valence level expansion

How Many Bonds ???

• most of our discussion will be centred on covalent bonding.

Valence level expansion

Some compounds occur which cannot be easily explained:

PF5, SF6, ClF7, ArF8

in each case the number of chemical bonds is equal to the number of valence electrons.

this can only happen if electrons are promoted to a higher energy level. In this case it is the adjacent ‘d’ orbital.

PF5

the normal orbital diagram looks like this:

with valence level expansion it looks like this:

Valence level expansion

includes elements of groups 15 to 18, from period 3 down; periods 1 and 2 do not have a ‘d’ orbital to promote to.

Please note that valence level expansion is the exception, not the rule.

Polar-Covalent Bonding

Covalent bonding implies equal sharing of electrons.

If sharing is not equal, the electrons in a bond will spend more time with one atom than the other.

Polar-Covalent Bonding

The atom where the electrons spend more time will have a net negative charge, while the atom at the other end of the bond will be positive.

This type of bond is polar-covalent.

Electronegativity

is a measure of how strongly an atom is holding on to its valence electrons.

If an atom loses an electron fairly easily it has a low electronegativity (and tends to be a cation).

If an atom tends not to lose electrons, but tends to steal them from other atoms (and become an anion) it has a high electronegativity.

To determine what type of bond exists between two atoms you subtract their respective electronegativities:

if the electronegativity difference is 0.2 or less, the bond is covalent

if the electronegativity difference is 1.7 or greater the bond is ionic.

Electronegativity

If the electronegativity difference between two atoms is between 0.3 and 1.6 the bond is polar-covalent.

The greater the electronegativity difference the greater the ionic character of the bond:

Assignment

Determine the electronegativity difference for each chemical bond. If the bond is polar covalent draw an arrow in the direction of the dipole, from positive to negative:

C - HEND = | 2.5 - 2.1 | = 0.4polar covalent bond

N - H B - F S - OP - H Si - Cl Cu - Br N - IBr - Cl O - H C - Cl C - O

N - H• END = |3.0 – 2.1| = 0.9• polar covalent bond: N - H

B - F• END = |2.0 – 4.0| = 2.0• ionic bond: [B]3+[F]1-

S – O• END = |2.5 – 3.5| = 1.0• polar covalent bond: S - O

P - H• END = |2.1 – 2.1| = 0.0• covalent bond: P – H

Si - Cl• END = |1.8 – 3.0| = 1.2• polar covalent bond: Si - Cl

Cu - Br• END = |1.9 – 2.8| = 0.9• polar covalent bond: Cu - Br

N – I• END = |3.0 – 2.5| = 0.5• polar covalent bond: N - I

Br – Cl• END = |2.8 – 3.0| = 0.2• covalent bond: Br – Cl

O – H• END = |3.5 – 2.1| = 1.4• polar covalent bond: O – H

C – Cl• END = |2.5 – 3.0| = 0.5• polar covalent bond: C – Cl

C – O• END = |2.5 – 3.5| = 1.0• polar covalent bond: C – O

Covalently-Bonded Structures

we now have to consider molecules made of several atoms.

most of the following discussion will concern itself with molecules made with covalent or polar-covalent bonds.

ionic bonds (and others) will return later in the unit.

Lewis Structures

Lewis structures allow us to predict how atoms will come together to make molecules.

Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.

Writing Lewis Structures

1. Find the sum of valence electrons of all atoms in the polyatomic ion or molecule. If it is an anion, add

one electron for each negative charge.

If it is a cation, subtract one electron for each positive charge.

PCl35 + 3(7) = 26

Writing Lewis Structures

2. The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds.

Keep track of the electrons:

26 6 = 20

Writing Lewis Structures

3. Fill the octets of the outer atoms.

Keep track of the electrons:

26 6 = 20 18 = 2

Writing Lewis Structures

4. Fill the octet of the central atom.

Keep track of the electrons:

26 6 = 20 18 = 2 2 = 0

Writing Lewis Structures

5. If you run out of electrons before the central atom has an octet…

…form multiple bonds until it does.

Exceptions to Octet Rule

Electron Promotion group 2 central atom will have 4 electrons group 13 central atom will have 6 electrons

Valence Level Expansion group 15 central atom will have 10 electrons group 16 central atom will have 12 electrons group 17 central atom will have 14 electrons group 18 central atom will have 16 electrons

Writing Lewis Structures

Write Lewis Structures for the following molecules:

F2 SiF4

PCl3 C3H8

SeH2 C3H6

MgH2 C3H4

Coordinate Covalent bonds

is defined as a covalent bond where one atom provides both of the electrons:

Polyatomic Ions

we can use coordinate covalent bond theory to explain most ions:

Resonance

This is the Lewis structure we would draw for ozone, O3. -

+

Resonance

But this is at odds with the true, observed structure of ozone, in which… …both O—O bonds

are the same length.

…both outer oxygens have a charge of 1/2.

Resonance

One Lewis structure cannot accurately depict a molecule such as ozone.

We use multiple structures, resonance structures, to describe the molecule.

Resonance

Just as green is a synthesis of blue and yellow…

…ozone is a synthesis of these two resonance structures.

Resonance

In truth, the electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon.

They are not localized, but rather are delocalized.

Resonance

The organic compound benzene, C6H6, has two resonance structures.

It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.

Molecular Shapes

The shape of a molecule plays an important role in its reactivity.

By noting the number of bonding and nonbonding electron pairs we can easily predict the shape of the molecule.

What Determines the Shape of a Molecule?

Simply put, electron pairs, whether they be bonding or nonbonding, repel each other.

By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

Valence Shell Electron Pair Repulsion Theory (VSEPR)

“The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”

Electron-Domain Geometries

All one must do is count the number of electron domains in the Lewis structure.

The geometry will be that which corresponds to that number of electron domains.

Molecular Geometries

The electron-domain geometry is often not the shape of the molecule, however.

The molecular geometry is that defined by the positions of only the atoms in the molecules, not the nonbonding pairs.

Molecular Geometries

Within each electron domain, then, there might be more than one molecular geometry.

Group 2 Geometries

In this domain, there is only one molecular geometry: linear.

NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is.

Central Atom

Bonding Electron

s

Lone Pair Electrons

Bond Type

Shape

Example

Magnesium

2 0 all single linear MgI2

one double

linear MgO

Group 13 Geometries

because there are no lone pair electrons the molecular is planar (flat).

Central Atom

Bonding Electron

s

Lone Pair Electrons

Bond Type

Shape Example

Boron 3 0 all single trigonal planar

BI3

one double

linear BIO

one triple

linear BN

Group 14 Geometries

Central

Atom

Bonding Electron

s

Lone Pair Electrons

Bond Type

Shape Example

Carbon 4 0 all single tetrahedral

CH4

one double

trigonal planar

COH2

one triple, ortwo double

linear HCN

CO2

Summary for Groups 2, 13 & 14

because there are no lone pair electrons: central atom bonded to 1 atom: linear central atom bonded to 2 atoms: linear central atom bonded to 3 atoms:

trigonal planar

central atom bonded to 4 atoms: tetrahedral

Group 15 Geometries

Central Atom

Bonding Electron

s

Lone Pair Electrons

Bond Type

Shape Example

Nitrogen 3 1 all single

trigonal pyramidal

NH3

one double

angular NOH

one triple

linear N2

5 0 all single

trigonal bipyramidal

NCl5

Group 16 Geometries

Central Atom

Bonding Electron

s

Lone Pair Electrons

Bond Type

Shape Example

Oxygen 2 2 all single

angular H2O

one double

linear O2

6 0 all single

octahedral

OF6

Group 17 & 18 Geometries

group 17 normally makes 1 chemical bond

(linear) with valence level expansion can

make 7 bonds (ClF7).

group 18 normally makes no bonds with valence level expansion can

make 8 bonds (ArF8).

Larger Molecules In larger molecules, it makes more sense to talk

about the geometry about a particular atom rather than the geometry of the molecule as a whole.

This molecule is tetrahedral about the first carbon, planar trigonal about the second and angular about the single-bonded oxygen.

Larger Molecules

This approach makes sense, especially because larger molecules tend to react at a particular site in the molecule.

Using VSEPR Theory

Determine the shape of each molecule, or atom within the molecule:

F2 SiF4

PCl3 C3H8

SeH2 C3H6

MgH2 C3H4

Polarity

previously we discussed bond dipoles.

But just because a molecule possesses polar bonds does not mean the molecule as a whole will be polar.

Polarity

By adding the individual bond dipoles, one can determine the overall dipole moment for the molecule.

In other words we can see if the dipoles reinforce each other, or cancel out.

Polarity

Assignment Perform the following activities for the

molecules at the bottom of the slide:1. Draw a lewis structure2. Draw a structural diagram and indicate

shape.3. Determine the electronegativity difference

for each bond type and determine whether the bond is covalent, polar covalent or ionic

4. Determine if they will be polar, or non-polar.

F2 PCl3 SeH2 MgH2

SiF4 C3H8 C3H6 C3H4

Complete “Fun With Balls & Sticks”

Intermolecular Forces

Properties of Molecules

Properties of Molecules

are determined by how the molecules interact with each other.

How they interact is determined by the forces of attraction between molecules.

These are called intermolecular forces; the forces which act between molecules, to draw them together, forming the various phases of matter.

How do we measure force ?

the best way is by temperature. temperature is a measure of kinetic energy;

the lowest temperature represents zero kinetic energy.

for a substance to melt or boil the kinetic energy must overcome the intermolecular force.

the higher the melting temperature or boiling temperature, the greater the intermolecular force.

Van der Waals Forces

occur between covalently bonded molecules.

three kinds: London disersion dipole-dipole attraction hydrogen bonding

London Dispersion Forces

are the dominant forces between covalently bonded, non-polar molecules

based on the formation of instantaneous dipoles.

the more electrons in a molecule, the stronger the force.

Relationship of Boiling Point to Number of Electrons and Molar Mass

Melting Point Boiling Point

# e- °C °C

CH4 10 - 182 -161

C2H6 18 - 183 - 88

C3H8 26 - 190 - 44

C4H10 34 - 138 - 0.5

C8H18 66 - 57 +125

London dispersion forces

are influenced by shape: Normal Pentane

(C5H12)

m.p. -130C, b.p. 36°C

Neopentane (C5H12)

m.p. -20°C, b.p. 9°C

Dipole-Dipole attraction

is a force which acts between polar molecules (ex. H2S).

results from the attraction of the opposite poles of the permanent molecular dipoles.

These substances generally have higher melting and boiling points than non-polar molecules with similar molecular weights (or numbers of electrons).

Ion-Dipole Interactions A fourth type of force, ion-dipole interactions

are an important force in solutions of ions. The strength of these forces are what make

it possible for ionic substances to dissolve in polar solvents.

Dipole-Dipole Interactions

The more polar the molecule, the higher is its boiling point.

How Do We Explain This?

The nonpolar series (SnH4 to CH4) follow the expected trend.

The polar series follows the trend from H2Te through H2S, but water is quite an anomaly.

Hydrogen Bonding a specialized form of dipole-dipole attraction. It occurs as when O, N, and F are bonded to H, owing to the

large electronegativity difference: O - H 3.5 - 2.1 = 1.4 N - H 3.1 - 2.1 = 1.0 F - H 4.1 - 2.1 = 2.0

This is a stronger force than standard dipole-dipole attraction.

Molecules with hydrogen bonding will have boiling points and melting points quite a bit higher than molecules that have only dipole-dipole or London dispersion forces.

Hydrogen bonding is responsible for many of the unusual properties of water.

Hydrogen bonding

is responsible for folding, final structure and function of proteins

holds the DNA strands together, but allows them to be unzipped for copying and gene expression.

Relationship Between Polarity, Force and Boiling Point

# e- B.P. (°C) Polarity Force

C2H6 18 - 161 Non-polar London

Dispersion

H2S 18 - 60 Polar Dipole-

Dipole

H2O 10 + 100 Very polar Hydrogen Bonds

Ionic Bonding generally occurs in compounds of metals and non-metals

(salts). It is the result of the attraction of oppositely charged ions. The structures formed are very orderly and are given the

name crystal lattice. Ionic solids are called crystals. No sharing of electrons occurs between the ions in the

crystal lattice. As a result, ionic solids are brittle. Ionic solids conduct electricity only in the molten state, and

not very well. Ionic solids are characterized by very high melting and

boiling points.

Metallic Bonding is the bonding which occurs between metals in the Periodic

Table. It is characterized by close packing of the atoms, with the

electrons delocalized; that is, they are free to jump from atom to atom, filling unoccupied orbitals.

This free sharing of electrons allows metals to conduct electricity freely (copper conducts electricity 100 000 times better than molten NaCl).

The free electrons also act as a lubricant, allowing metal atoms to slide over one another without affecting the integrity of the material. Thus metals are malleable and ductile.

This bond is strong, giving most metals high melting and boiling points.

This bond is also variable.

Network Covalent Bonding This is the traditional covalent bond, expanded to 2 or 3

dimensions in a network which is theoretically infinite (much like an ionic crystal lattice).

Network solids include diamonds, graphite, quartz, and most rocks.

Because the covalent bond is stronger than any other bond, the network solid is very hard (diamonds are the hardest substance known).

Because electrons are held tightly in their bonds, network solids are brittle, and they do not conduct electricity.

Because the orientation of the atoms is very specific to the bonding orientation (tetrahedral, planar trigonal), network solids form distinct crystals.

Because of the strength of the bonds, network solids have very high melting and boiling points

Summarizing Intermolecular Forces

Intermolecular Forces Affect Many Physical Properties

The strength of the attractions between particles can greatly affect the properties of a substance or solution.

Viscosity• Resistance of a liquid

to flow is called viscosity.

• It is related to the ease with which molecules can move past each other.

• Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

Surface Tension

Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

Vapor Pressure

• At any temperature, some molecules in a liquid have enough energy to escape.

• As the temperature rises, the fraction of molecules that have enough energy to escape increases.

Vapor Pressure

As more molecules escape the liquid, the pressure they exert increases.

Vapor Pressure

The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.

Vapor Pressure

• The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure.

• The normal boiling point is the temperature at which its vapor pressure is 760 torr.

Intermolecular Forces in Summary

Strength of Attraction

Molecular (London dispersion)

1 to 50

Molecular (dipole-dipole)

5

Molecular (hydrogen bonding)

15

Ionic 50

Metallic 20 to 80

Network Covalent 100

Melting and Boiling PointMolecular (London dispersion)

very low (nitrogen boils at - 196 C) VARIABLE

Molecular (dipole-dipole)

low (H2S boils at -61 C)

Molecular (hydrogen bonding)

medium (H2O boils at +100 C)

Ionic high (NaCl boils at +1413 C)

Metallic variable (Hg @ +357 C, W @ +5660 C)

Network Covalent very high (SiO2 boils at 2600 C)

Properties of SolidsMolecular (London dispersion)

soft, waxy

Molecular (dipole-dipole)

more rigid

Molecular (hydrogen bonding)

crystalline, brittle

Ionic long-range crystalline,hard, brittle

Metallic short-range crystalline,ductile, malleable

Network Covalent long-range crystalline,hard, brittle

Conductance of Heat and Electricity

Molecular (London dispersion)

non-conductor

Molecular (dipole-dipole)

non-conductor

Molecular (hydrogen bonding)

non-conductor

Ionic conductor in liquid or dissolved phase

Metallic conductor in solid or liquid phase

Network Covalent non-conductor in 3d form, some conductance in 2d

Solubility in H2OMolecular (London dispersion)

Not soluble

Molecular (dipole-dipole)

Soluble

Molecular (hydrogen bonding)

Soluble

Ionic Soluble

Metallic Not soluble

Network Covalent Not soluble