Chemistry 100 Chapter 8

Chemical Bonding Basic Concepts

Chemical Bonds

Three basic types of bonds Ionic

Electrostatic attraction between ions.

Covalent Sharing of electrons.

Metallic Metal atoms bonded

to several other atoms.

The Valance Electrons

When atoms interact to form chemical bonds, only the outer (valance) electrons take part.

Need a tool for keeping track of valence electrons, e.g., The Lewis dot symbol

1 v.E. 7 v.E’s

These two elements combine to form an ionic compound

2 Na (s) + Cl2 (g) ® 2 NaCl (s)

+ NaClNa ...Cl: :.

What’s Happening?

[Ne]3s1 [Ne]3s23p5

(g) ® Na+ (g) + e- (ionizes, loses e-) an electron configuration of [Ne]

(g) + e- ® Cl- (g) an electron configuration of [Ar]

In the crystal lattice, Na+ and Cl- ions; strong electrostatic

attractions

Na.

+ NaClNa ...Cl: :.

Cl::..

.

The NaCl Crystal

Ionic Bonding

Electrostatic attractions that hold ions together in an ionic compound.

The strength of interaction depends on charge magnitude and distance between them.

rqkq

E 21ionic

q1 magnitude of charge 1q2 magnitude of charge 2r distance between the ionic centres

Stability of Ionic Compounds

The stability of ionic compounds depends on two main factors

1. The electron affinity of one of the elements2. The ionization energy of the other

Note electron affinities and ionization potentials

are gas-phase reactions? How are they related to the stability of solid

materials?

The Lattice Energy

A quantitative measure of just how strong the interaction is between the ionic centres (i.e., a measure of the strength of the ionic bond)

For the reaction KCl (s) K+ (g) + Cl- (g) H = 718 kJ/mol

Lattice energy (DlatH). The energy required to completely separate

one mole of the solid ionic compound into its gas-phase ions.

Lattice Energies of Various Ionic Compounds

Determined using a thermochemical cycle -

the Born-Haber cycle (a Hess’s Law application)

Covalent Bonding

In a wide variety of molecules, the bonding atoms fulfill their valance shell requirements by sharing electrons between them.

Covalent bonds - a bond in which the electrons are shared by two atoms.

H2 ® H-H, F2 ® F-F, Cl2 Cl-Cl For many electron atoms (like F and Cl), we

again to worry only about the outermost (valence) electrons.

Covalent Bonding

In covalent bonds, atoms share electrons.

There are several electrostatic interactions in these bonds: Attractions between

electrons and nuclei, Repulsions between

electrons, Repulsions between

nuclei.

Examples of Covalent Bonding

Let’s look at the Cl2 example. Each Cl atom has 7 valence shell

electrons 3 Lone pairs and one unpaired electron

Cl::..

.Lone pairs

Unshared electron

The Cl2 Molecule

The structure we have just drawn are called Lewis structures.

The dash between the atomic centres represents bonding electrons

Redraw F2

Cl Cl:..

. . :....

lone pairs (non bonding)bonding electrons

F F:..

. . :....

Note both Cl2 and F2 satisfy their valence shell requirements by the formation of a single bond.

What about O2? How can we satisfy the octet rule for 2 O atoms?

O O... .

....

Valence shell requirements are satisfied by the formation of a double bond.

check out N2 :NºN: (triple bond) Note that the octet rule works

mainly for the second row elements. Filled valence shells can have more

than 8 electrons after Z=14 (Si). This is generally termed octet expansion.

Covalent Compounds

Compounds that contain only covalent bonds are called covalent compounds.

There are two main of covalent compounds, Molecular covalent compounds (CO2, C2H4) Network covalent compounds (SiO2, BeCl2). The network covalent compound are

characterized by an extensive “3-D” network bonding

Comparison between Ionic and Covalent Compounds

Ionic Compounds usually solids with very

high melting points conduct electricity when

molten (melted) usually quite water

soluble and they are electrolytes in aqueous solution

NaCl

Covalent Compounds usually low melting solids,

gases or liquids don’t conduct electricity

when molten aren’t very soluble in water

and are non electrolytes CCl4

The Filled Valence Shell rule

Filled Valence Shell rule Atoms participate in the formation of bonds

(either ionic or covalent) in order to satisfy their valence shell requirements.

Atoms other than H tend to form bonds until they end up being surrounded by 8 valence electrons (the noble gas configuration). Your text calls this the “octet” rule.

Electronegativity

Electronegativity is defined as the ability of an atom to attract electrons towards itself in a molecule ( (pronounced ‘chi’))

Examine the H-F covalent bond+H-F

denotes a partial “+” charge on the H atom - denotes a partial “-“ charge on F atom

Electronegativity is related to the electron affinity and the ionization energy.

Compare the following elements. Na low I1, small negative E.A. low F high I1, large, negative E.A., high

Trends in the Values

Across a row The values generally increase as we

proceed from left to right in the periodic table.

Down a group The values generally decrease as we

descend the group. Transition metals

Essentially constant values

Plot of Values

Electronegativity and Bond Type

Can we use the electronegativity values to help us deduce the type of bonding in compounds?

values bond type

0.0 < < 0.5 non-polar covalent

0.5 1.9 polar covalent

2.0 3.3 Ionic bond

An Outline for Drawing Lewis Structures

Predict arrangement of atoms (i.e., predict the skeletal arrangement of the molecule or ion). · The H is always a terminal atom,

bonded to ONE OTHER ATOM ONLY. A halogen atom is usually a terminal atom.

· Note that the central atom usually has the least negative electron affinity.

Count total number of valence shell electrons (include ionic charges).

Place 1 pair electrons (sigma bond, ) between each pair of bonded atoms (i.e., the central atom and each one of the terminal atoms).

Place remaining electrons around the terminal atoms to satisfy the filled valence shell rule. (lone pairs).

All remaining electrons are assigned to the central atom. Atoms in the 3rd or higher row can have more than eight electrons around them. If a central atom does not have a filled

valence shell, use a lone pair of electrons from a terminal atom to make a pi () bond.

Formal Charges

Definition: formal charge on atom = number of valence electrons – number of non-bonding - ½ the number of bonding electrons.

Formal charge in a Lewis Structure is a bookkeeping “device” keeps track of the electrons “associated”

with certain atoms in the molecule vs. the valence e-‘s in the isolated atom!

How does it work?

Rules for Formal Charges

Neutral molecules ® S formal charges = 0 Ions ® S formal charges = charge of ion For molecules where the possibility of multiple

Lewis Structures with different formal charges exist Neutral molecule - choose the structure with

the fewest formal charges. Structures with large formal charges are less

likely than ones with small formal charges Two Lewis Structures with similar formal

charge distribution ® negative formal charges on more electronegative atom

Resonance

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

Resonance

Note the true, observed structure of ozone… …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 like 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.

Experimental Evidence for Resonance.

The resonance structures for benzene C6H6

We would expect to find two different bond lengths in benzene (C=C and C-C bonds). C= C ® bond length = 133 pm = 0.133 nm C-C ® bond length = 0.154 nm

Experimentally, all benzene carbon-carbon bond lengths are equivalent at 0.140 nm

Exceptions to the Filled Valence Shell Rule

Be compounds BeH2, BeCl2, Boron and Al compounds BF3, AlCl3, BCl3 BF3 is stable Þ The B central atom has a tendency

to pick up an unshared e- pair from another compound

BF3 + NH3 ® BF3NH3 the B-N bond is an example of a coordinate

covalent bond, or a “dative” bond ® i.e. a bond in which one of the atoms donates both bonding electrons.

Odd e- molecules

These molecules have uneven numbers of electrons \ no way that they can form octets.

Examples NO and NO2. These species have an odd

number of electrons.

N OO.

..

......:N O

.....

:

Look at the dimerization reaction of NO2

.

2 NO2 (g) ⇄ N2O4 (g) Keq = 210

N N

O

OO

O: ..: : :

: : : :..

Valence Shells having more than 8 Electrons (Expanded Octets)

A central atom having more than 8 valance shell electrons is possible with atomic number 14 and above.

Cl

Cl ClPCl Cl

Reason - elements in this category can use the energetically low-lying d orbitals to

accommodate extra electrons

Look at HClO3

High formal charge on the electronegative Cl atom (f.c.(Cl) = 7-2-1/2 (6) = +2)

This resonance structure would make a very small contribution to the overall resonance hybrid.

Cl O

O

O H......:

..: :

..

..

With the possibility of using the low lying d-orbitals on the Cl atom to accommodate extra electron pairs, we may write other Lewis structures

Note: the final three structures reduce the formal charges

Cl O

O

O H......:

..: :

..

.. Cl O

O

O H......

..: :

..

..Cl O

O

O H......

: :....:

Cl O

O

O H......

: :....

Bond Energies and Thermochemistry

Look at the energy required to break 1 mole of gaseous diatomic molecules into their constituent gaseous atoms.

H2 (g) ® H (g) + H (g) DH° = 436.4 kJCl2 (g) ® Cl (g) + Cl (g) DH° = 242 kJ

These enthalpy changes are called bond dissociation energies. In the above examples, the enthalpy changes are designated D (H-H) and D (Cl-Cl).

For Polyatomic Molecules.

CO2 (g) ® C (g) + 2 O (g) DH = 745 kJ Denote the DH of this reaction D(C=O) What about dissociating methane into C + 4

H’s?CH4 (g) ® C(g) + 4 H (g) DH° = 1650 kJ

Note 4 C-H bonds in CH4 \ D (C-H) = 412 kJ/mol

H2O (g) ® 2 H (g) + O (g) DH° = 929 kJ/mol H2O It takes more energy to break the first O-H

bond.H2O (g) ® H (g) + OH (g) DH° = 502 kJ/mol H2OHO (g) ® H (g) + O (g) DH = 427 kJ/mol H2O Note: we realize that all chemical reactions

involve the breaking and reforming of chemical bonds. Break bonds add energy. Make bonds energy is released.

rxnH° S D(bonds broken) - S D(bonds formed)

These are close but not quite exact. Why? The bond energies we use are averaged bond

energies, i.e., This is a good approximate for equations

involving diatomic species. We can only use the above procedure for GAS

PHASE REACTIONS ONLY.