Bonding Chapter 8

Two general types of bonds Ionic: the transfer of electrons from a metal to a non metal. Covalent: the sharing of electrons between atoms. Valence electrons: the electrons being transferred or shared

Valence Electrons The electrons used to form bonds

Valence electrons are depicted using Lewis Dot Structures

Ionic Bonding

Covalent Bonding A type of bonding between non-metals. This involves the sharing of electrons by 2 or more atoms. The electrons are shared, not transferred as with ionic bonds.

When nonmetals bond together a covalent bond is created and we call them molecules or molecular compounds!

Molecules Molecules are neutral atoms that are joined together by covalent bonds Molecular formula - shows you how many atoms of each element is in a substance Example: CO2 , NH4

Octet Rule and Covalent Bonding An octet is 8 valence electrons around an atom to achieve a noble gas configuration! Molecules want the same thing, but they share their valence electrons to achieve the octet rule.

EXAMPLES INCLUDE THE BONDS BETWEEN: H2 F2 Br2 Cl2 HCl H2O Other than hydrogen these elements are on the right side of the periodic table.

Cl Cl Lewis Dot Structures Formulas in which atomic symbols represent the element and all inner-shell electrons, dots represent valence electrons and dashes between two atomic symbols represent electron pairs in covalent bonds. Cl Cl

Single Covalent Bonds When atoms share one pair of electrons they form a single covalent bond

Chlorine forms a covalent bond with itself Cl2

How will two chlorine atoms react? Cl Cl

Cl Cl Each chlorine atom wants to gain one electron to achieve an octet

Cl Cl what can they do to achieve an octet? Share unpaired electrons!

Cl Cl

Cl Cl

Cl Cl

Cl Cl

Cl Cl octet

Cl Cl octet

Cl Cl The octet is achieved by each atom sharing the electron pair in the middle

Cl Cl The octet is achieved by each atom sharing the electron pair in the middle

Cl Cl This is the bonding pair (shared pair of electrons)

Cl Cl It is a single bonding pair

Cl Cl It is called a SINGLE BOND

Single bonds are abbreviated Cl Cl Single bonds are abbreviated with a dash Normally in the final structure the valence electrons are not drawn

This is the chlorine molecule, Cl Cl This is the chlorine molecule, Cl2

Neutral Molecules 4 bonds 3 bonds 2 bonds 1 bond

Double Bonds Sharing of two pairs of electrons between two atoms

O2 Oxygen is also one of the diatomic molecules

O How will two oxygen atoms bond?

O Each atom has two unpaired electrons

O

O

O

O

O

O

O Oxygen atoms are highly electronegative. So both atoms want to gain two electrons.

O Oxygen atoms are highly electronegative. So both atoms want to gain two electrons.

O

O O

O O

O O

Both electron pairs are shared.

O O 6 valence electrons plus 2 shared electrons = full octet

O O 6 valence electrons plus 2 shared electrons = full octet

O O two bonding pairs, making a double bond

O O = For convenience, the double bond can be shown as two dashes.

This is the oxygen molecule, = This is the oxygen molecule, O2

Triple Bonds Atoms that share three pairs of electrons: Example: N2

Molecular Shapes

bond polarity Covalent bonds can be polar or non-polar. If a bond is non-polar, that means that there is an equal sharing of electrons between atoms (Cl2). If a bond is polar, that means the electrons are not shared equally, making one side of the bond more negative and the other side more positive. The electrons are more strongly attracted to the more electronegative atom.

The dipole’s direction is from the positive pole to the negative pole. Dipole – is created by opposite charges that are separated by a short distance. The dipole’s direction is from the positive pole to the negative pole. The dipole (magnet) is created due to electronegativity differences between atoms. H Cl (H= 2.1; Cl = 3.0)

Electronegativity 0.0-0.3 Non polar covalent 0.3-1.7 polar covalent >/= 1.8 Ionic

Polar Compounds Need to look at all the bonds in the compound. Are individual bonds polar or non polar. What is the sum of all the bonds. atom. electronegativity

Bond Lengths and Bond Strengths

Multiple Bonds Double bonds are stronger and shorter than single bonds. Triple bonds are even stronger and shorter than both double and single bonds. Common between carbon, oxygen and nitrogen atoms.

Bond Lengths and Bond Energies