Chemical Chemical BondingBonding

• Bonding within a molecule is called intramolecular attraction– Ionic bonds– Covalent

bonds– Polar covalent

bonds

  Electronegativity

• Electronegativity is the tendency of an atom to attract electrons.

• Group 1 and 2 lose electron easily, these elements have a low electonegativity

• Group 17 and 16 can attain electrons easily, these elements have high electronegativities

• Delta represents a partial charge or fractional charge

0 : H

+-

N : H

+-

Ionic Bonds• Steals an e-, this causes one

atom to have a positive charge and the other a negative charge. This is the reason for the attraction

• Opposites attract!!• Group 1 and 2 loose

electrons to group 16 and 17 (usually)

• When the difference in electonegativities between atoms is greater than 1.7 the molecules is ionic

• Example NaCl

• Na = .9

• Cl = 3.0

• 3.0 - .9 = 2.1

• 2.1 > 1.7 = ionic bond

• Cl would be the more negative atom because it has a higher electronegativity (it steals the e- from Na)

Covalent Bonds

• Atoms share electrons because they have similar electronegativities

• Atoms share electron to fill octet rule

• Hydrogen form stable molecules were it shares two electrons, this is the duet rule

• Single bond formed when atoms share one pair of e-

• Double bond formed when atoms share two pair of e-

• Triple bond formed when atoms share three pair of e-

• The electonegativites differences are 0 - .3

• Example O2

• O = 3.5• 3.5 – 3.5 = 0• 0 falls between 0

and .3 therefore O2 is a covalent bond

More examples of covalent bonds

Ø     

Polar covalent bonds• Electrons not always shared equally. • The atom with the higher

electronegativity attract the shared electron pair more strongly pulling it away from the other atom.

• The shared pair is shifted from the center between the two participating atoms making one end of the molecule positive and the other end negative. The bond is polarized. (Dipole – one side of molecule is slightly negative and one part slightly positive)

• The difference in electronegativity among the atoms is .4 to 1.7

Now practice with worksheet #55

• Example PO3

• P = 2.1• O = 3.5• 3.5 – 2.1 = 1.4• 1.4 falls between .4 and

1.7 therefore PO3 is a polar covalent bond

• Oxygen would be the more negative atom because of the greaterelectonegativity

Lewis Structures

Represents individual valence electrons

Represents a nonbonding pair or lone pair of electrons

Represents a pair of e-

B. Lewis Structures

• Electron Dot Diagrams– show valence e- as dots– distribute dots like arrows

in an orbital diagram– Show a single line for a single bond, double line for a double bond,

triple line for a triple bond– 4 sides = 1 s-orbital, 3 p-orbitals– EX: oxygen

2s 2pO

X

Lewis Structures• 1) Count all valence electrons in all

the atoms in the molecule; it doesn’t matter which atoms they come from.

• 2) If the compound has more than 2 atoms, the least electronegative atom is the central atom, often a single atom. If carbon is present, it is almost always the central atom. Hydrogen is never a central atom since it can only form one bond.)

• 3) Make a bond between the central atom and the other atoms using a dash (-) to show a pair of shared electrons.

• 4) Place the remaining valence electrons around each of the atoms so that they have an octet.

Lewis structure

• When more than one Lewis structure can be drawn for a particular molecule this molecule exhibits resonance

• Example of an electron dot diagram for CCl4

• First step determine the number of valence electrons– C has 4 (remember 2 from 2s

and 2 from 2p)– Cl has 7 (4) (remember the 7

comes from 2e- from the 3s and 5ee from 3p) the 4 comes from Cl4

– 4 + 7 (4) = 32 valence electrons

CLO O

O

Draw the e- dot diagram

Now count to see if you have 32 valence e-

• Your turn, draw the e- dot diagram for CH3I• Step 1 – count valence e-• Step 2 – draw e- dot diagram remember (If the compound has more

than 2 atoms, the least electronegative atom is the central atom, often a single atom. If carbon is present, it is almost always the central atom. Hydrogen is never a central atom since it can only form one bond.)

• Step 3 count to see if you have the correct number of Valence e-

Did you remember to:

• Check to see if all of the valence electrons got used?

• Check to make sure every atom has an octet of electrons (or 2 for Hydrogen)

• There are 14 Valence e- (4 +3 + 7)• H only needs two e- to satisfy the duet rule• Iodine and Carbon needs 8 e- to satisfy the octet rule

CH

I

H

HHow did you do? Now practice with worksheet #56 and #57

Metallic Bonding

• Metallic bonding is the attraction between metal atoms and a sea of surrounding valence electrons

• A result of this bonding is mobile electons which gives rise to the excellent electrical conductivity of most metals.

Intermolecular Forces

• Intermolecular forces – bond that holds molecules together.– Effect boiling and

freezing points

– http://www.bcpl.net/~kdrews/interactions/interactions.html

Intermolecular Forces

• Dipole – Dipole– Molecules are attracted to

each other as a result of partial charges of dipole molecules this is called a dipole-dipole intermolecular force

– In general, intermolecular forces are about 1% as strong as intramolecular forces.

Intermolecular Forces

• Hydrogen Bonding– Especially strong dipole-

dipole are called hydrogen bonding, example = water

– When F,O and N are attached to hydrogen they will form hydrogen bonds with other molecules

– Common molecules that form hydrogen bonds are HF, H2O and NH3.

– Molecules with an O-Hbond, like alcohols also exhibit hydrogen bonding

Hydrogen Bonding – type of attraction that holds two water

molecules together.• Cohesion – attractive force between

particles of the same kind

• Adhesion – attractive force between unlike substances (meniscus)

Intermolecular Forces• London Dispersion Forces

(LDF)– The force that holds noble gas

atoms and nonpolar molecules together

– No permanent dipoles however, at any given time, the e- can be mostly on one side of the molecule creating a slightly negative charge on one side compared to the other

– Weak and easily broken force call LDF

– Larger the molecule the more e- it has therefore the stronger LDF

Bonding forces and boiling and freezing points

• Strongest bonding force is a hydrogen bond, followed by a dipole-dipole force, and then LDF.

• H-Bonds have highest boiling, melting and freezing points

• Dipole-dipole force have intermediate boiling, freezing and melting points

• LDF have lowest boiling, freezing and melting points

Now practice with worksheet #58

VSEPR model• Valance Shell Electron Pair

Repulsion (VSEPR) is a 3D model of a molecule

1. Draw the Lewis structure for the molecule

2. Count the electron pairs and arrange them in the way that minimizes repulsion(that is, put the pairs as far apart as possible)

3. Determine the positions of the atoms from the way the electron pairs are shared

4. Determine the name of the molecular structure from the positions of the atoms.

5. http://wunmr.wustl.edu/EduDev/Vsepr/table1.html

Linear

• Number of bonding pairs around central atom = 2

• Number of lone pairs around central atom = 0

• 180 bond• Example- BeCl2, CO2

Triangle (Trigonal planar structure)

• Number of bonding pairs around central atom = 3

• Number of lone pairs around central atom = 0

• Examples- BF3, SO3

tetrahedral

• Number of bonding pairs around central atom = 4

• Number of lone pairs around central atom = 0

• 109.5 bond angle• Example- CH4

                          

Trigonal pyramid

• Number of bonding pairs around central atom = 3

• Number of lone pairs around central atom = 1

• 120 bond angle• Example-NH3

                            

Bent or v-shaped

• Number of bonding pairs around central atom = 2

• Number of lone pairs around central atom = 1 or 2

• 106 bond angle• Example H2O, SO2