Chemical Bonding

Compounds are formed from chemically bound atoms or ions

Substances become more stable through chemical bonding, where

2 or more atoms are joined together by a simultaneous

attraction.

Valence electrons are electrons in the highest occupied energy level of an atom ( the last shell).

Bonding involves only valence electrons

i. Ionic Bonding

ii. Covalent Bonding

iii.Metallic Bonding

Ionic Bonds occur when the more electronegative element “steals” the electron pair away from the other atom.

The atom that has stolen the electron pair becomes a negativeion (anion) while the “victim” becomes a positive ion (cation).

The two atoms are held together by their opposite charges.

Na Cl

Can you predict which atoms will gain electrons and which will looseelectrons by looking at the trend in electronegativity?

Increase in Electronegativity

Inc

reas

e in

Ele

ctro

nega

tivi

ty

When you consider that for an ionic bond to form there must be agreat deal of difference in electronegativity between the atoms, can you predict what two types of atoms allow this to occur?

Metals Non-Metals

Ionic Properties

Why do most ionic compounds have similar properties?

We can hypothesis that it is due to the bonds formed between the

ions, holding them firmly in a rigid structure

Forming Ions

Metals tend to lose electrons and

become positive (cations)

Non-metals tend to gain electrons and

become negative (anions)

An ionic compound is formed when a

metal bonds to a non-metal. The total

charge of any compound is zero.

Ex:

Na and Cl which one will lose

electrons which one will gain

electrons

Write out Tin’s electron

configuration what will it do??

Lattice energy

measure of the energy released when

ions are combined to make a compound

Directly related to Coulombs Law, the

potential energy (E) between two ions is

directly proportional to the product of

their charges and inversely proportional

to the distance of separation between

them.

The anions and cations in an ionic compound are locked in a regular neutrally charged structure, held by the balance of attractive bonds and electrical repulsion.

The component ions in such crystals are arranged in repeating three-dimensional

(3-D) patterns.

Alkali metals combine with halogens in 1:1 ratios since alkali metals need to lose 1 e1- and halogens need to

gain 1e1-.

Alkaline earth metals combine with halogens in 1:2 ratios since alkaline earth metals need to lose 2 e1- and

halogens need to gain 1e1-.

LEWIS DOT STRUCTURES: Elements

Board Practice

Elements #1-20

Lewis Structures can be used to illustrate the formation of ionic

bonds.

Be 2 F+ [Be]2+F F1- 1-

Write an equation with electron dot diagrams to illustrate the

formation of aluminum chloride.

Al 3 Cl+ [Al]3+Cl Cl

1- 1-

Cl1-

Lewis Structures

Duet Rule = applies to H and He and states these two atoms are stable with 2 electrons in their outer shell

Octet Rule= elements are most stable with 8 electrons in their outer shell

1. Most are crystalline in structure

2. High melting/boiling points

3. Electrically neutral

4. Can conduct electricity when

melted or in aqueous solution

5. Hard/ Brittle

Electrostatic attraction force

between the cation and free

electrons.

Any successful bonding model for

metals must account for the typical

physical properties of metals:

malleability, ductility, and efficient and

uniform conduction of heat and

electricity in all directions.

Most metals are durable and have high

melting points.

These facts indicate that the bonding in

most metals are strong and non-

directional.

Metal atoms are arranged in very

compact and orderly patterns.

i) Body-centered cubic

ii) Face-centered cubic

iii) Hexagonal close-packed

1. Can conduct electricity (free electrons)

2. Malleable (put into shape)

3. Ductile ( made into wires)

4. Good conductors of heat

5. Metals are usually shiny

Covalent Bonding

Br + Br Br Br

O + O O O

Covalent Bonding

Electrons are shared by

nuclei Electrons can be shared equally (

non-polar covalent) or unequally

(polar covalent)

Types of Bonds1) Single bond – 1 pair of e- are shared

- weakest bond

-longest bond length

2) Double bond – 2 pairs of e- are shared

3) Triple bond- 3 pairs of e- are shared

- strongest bond

- shortest bond length

If the sharing is equal, this is called

Non – polar COVALENT BONDING

H HElectron pair

H H

when the sharing

is unequal, one

atom becomes

slightly positive

the other slightly

negative creating

a polar covalent

bond

The force of attraction of an element’s nucleus for electrons is called electronegativity. Atoms of different elements have different electronegativities. The higher the electronegativity, the stronger the attraction for electron pairs.

HFDifference

in En?

The bonding electrons are on the average closer to the fluorine than to the hydrogen atom.

The movement of the negatively charged electrons away from hydrogen toward fluorine, due to a difference in electronegativity, builds up a partial negative charge on the fluorine and a partial positive charge on the hydrogen.

This is not a complete transfer of an electron from hydrogen to fluorine; it is merely a drifting of electrons toward fluorine.

H

HWhen a charge separation of this

type is present, the molecule possesses an electric dipole, and

the bond is called a POLARCOVALENT BOND , or simply

a POLAR BOND.

HPolar covalent bond (polar bond)

covalent bond joins two atoms of

different elements and the bonding

electrons are shared unequally

Non-polar covalent bond

bonding electrons are shared

equally

1. Soft , brittle solids

2. Low boiling/melting points

3. Tend to be more flammable

4. Do not conduct electricity

5. Usually non-soluble in water

Lewis Structures for

Molecular Compounds

N N

Bonding capacity is the number of covalent bonds (shared electron pairs) that an atom can form.

Covalent molecules often consist of atoms of different elements, with

different bonding capacities.

Exceptions to the Octet Rule

1) B and Be usually have less than 8

electrons

2) Elements in the 3rd energy level and

above can have more than 8 electrons

in their outer shell

1. Determine the central atom.

The least electronegative element

The atom with the smallest number of

valence electrons

The “oddball” element

It is NOT Hydrogen

2. Count the number of valence electrons

in each atom.

Ex: NO3- N = 5 O = 6

3. Include the number of electrons lost or

gained due to charge.

Ex: NO3- N = 5 O = 6 -ve = 1

4. Total the number of valence electrons.

Ex: NO3-

N 5

3O 18

-ve 1

24

5. Draw the skeletal structure with the

central atom in the middle.

6. Draw the first bonds connecting all

atoms to the central atom.

7. Subtract the number of electrons

involved in these bonds (two electrons for

each bond) from the total number of

electrons.

24 – 6 = 18

8. Finish adding electrons around the

atoms up to the number of electrons left.

Work from the outside in.

9. Make sure all atoms satisfy the octet

rule. If they don’t move electrons around to

form multiple bonds (double or triple) so

that everything has a full octet

10. If it is an ion, put [square brackets]

around the diagram and put the charge

outside.

[ ]-

Resonance

= occurs when more than

one valid Lewis structure

can be written for a

particular molecule

Ex. NO3-1

Formal Charge

charge assigned to an atom in a molecule,

assuming that electrons in all chemical

bonds are shared equally between atoms,

regardless of relative electronegativity,

helps determine the most correct structure

of a molecule when there is more than one

option.

Formal Charge

FC= # of valence electrons – (# of covalent bonds) –

(# of electrons in lone pairs)

The total formal charge should add up to

the charge on the molecule

The most correct structure will be the one

with the lowest formal charge on each

element in the molecule

Draw Lewis Structures for the following:

H2O, NF3, Cl2, SnCl2, PCl5, SO3,

BeCl2, C2H6, C2H2, ClF3, CHCl3, ICl,

O2, N2, SF6, CO2, BF3, C2H4, O3, IF7

Compounds are arranged

in many different shapes

The VSEPR Theory states that

because electron pairs repel,

molecular shape adjusts so the

valence-electron pairs are as far

apart as possible.

VSEPR model

Valence Shell Electron Pair

Repulsion

Used to predict the geometry

of molecules

The structure will minimize

electron pair repulsions

Types of Molecular Shapes

There are several classifications of

molecular geometry based on the number

of electron domains:

2 - Linear

3 - Trigonal planar

4 - Tetrahedral

5 - Trigonal bipyramidal

6 - Octahedral

Intermolecular forces play a key

role in determining the physical

and chemical properties of

covalent compounds.

Van Der Waals consists of 2 possible types of forces:

1. London Dispersion Forces

2. Dipole-Dipole Forces

-This is the only type of force present in non-polar covalent molecules.

-It is the weakest of the intermolecular interactions caused by the motion of the electrons.

-The strength of dispersion forces generally increases as the number of electrons in a molecules increases because the molecule is more polarizable.

Ex. Halogen diatomic molecules.

- This occurs when polar covalent bonds are attracted to one another.

- Electrostatic attractions occur between oppositely charged regions. (partially (–) and partially (+)).

- Dipole interactions are similar to but much weaker than ionic bonds.

Dipolar molecules

Have a center of positive charge and a center of negative charge.

aka: dipole moment

Ex. HF

Dipole moment in NH3

Dipole cancels out in CO2

This is found in polar covalent molecules that

have hydrogen that is bonded to a very

electrostatic element (N, F, O) and attracted to a

(N, O, F) in another adjacent molecule.

Hydrogen bonds are the strongest of the

intermolecular forces.

Hydrogen > dipole-dipole > London Dispersion

Bonds interactions forces

Hydrogen bonds are extremely important in

determining the properties of water and biological

molecules such as proteins.

The water molecule has a bent shape (105°) and

is considered to be polar and the universal solvent.

The attraction in water results from the

intermolecular hydrogen bonds.

Surface tension: the inward force, or pull that tends to minimize the surface area of a liquid

- this surface tension tends to hold a drop of liquid in a spherical shape

The higher the surface tension, the more nearly spherical is the drop of that particular.

Because of hydrogen bonding, water absorbs

a large amount of heat as it evaporates or

vaporizes.

The hydrogen bonds must be broken before

water changes from the liquid to vapor state.

Vapor Pressure the force exerted due to the

gas above the liquid

Boiling Point: occurs when the

temperature at which the vapor pressure

of the liquid is just equal to the external

pressure.

Boiling leads to evaporation of a liquid.

In the case of water, hydrogen bonds

break in order for the liquid to vaporize.