Chemical Bonding

• Although we have talked about atoms

and molecules individually, the world

around us is almost entirely made of

compounds and mixtures of compounds.

• We are going to take an in depth look at

these compounds and the interactions

of the atoms that hold them together

and make up the compounds

Bonds• Bonds are the force that holds

groups of two or more atoms

together and makes them function as

a unit.

• Bond Energy is the energy required

to break the bond between two

atoms.

Ionic Bonding

• Generally occurs between a Metal and a

Non-metal

• Cations lose electrons, Anions gain electrons

• Electrons are transferred

• Opposite charges on atoms attracts them to

one another

Covalent Bonding• Generally occurs between a nonmetal and a

nonmetal

• Both atoms share electrons to achieve a

lower energy state.

• Electrons are “shared”. More like a tug of

war.

• Same charges, lower energy is responsible

for bonds

Nonpolar Covalent Bonding

• Covalent bonds where electrons are shared

equally between two atoms.

• Atoms must have the same values of

electronegativity

• If a covalent bond is like a tug of war a

nonpolar covalent bond would be a

stalemate.

Polar Covalent Bonding

• During the “tug of war” in covalent bonding

electrons aren’t always shared equally.

• Some atoms have a stronger attraction for

electrons and pull them closer than other atoms.

• This unequal sharing of electrons causes one

atom to have a small positive charge and one to

have a small negative charge

Electronegativity

• Attraction of shared electrons to an

atom.

• Determines the type of bond

• Can calculate a value of

electronegativity based on relative

values for each element.

Electronegativity• For differences in electronegativity,

generally:

O = nonpolar covalent examples: Cl-

Cl, C-S

.1-1.5 = polar covalent examples: C-F,

P-S

1.6-3.3 = Ionic examples: Na-O, K-I

Practice• Identify the following as Ionic, polar

covalent, or nonpolar covalent bonds:

S-F Mg-Cl

Br-Br B-F

B-N N-Cl

P-I Mn-S

Dipole Moments

• Polar covalent bonds that do not share

electrons equally are said to have a dipole

moment.

• The atom pulling the electrons the strongest

or with the higher electronegativity will have

a partial negative charge.

• The atom with the weaker pull on electrons

will have a partial positive charge.

Lewis Dot Structures

• Representations of atoms or molecules which show

the valence electrons around an atom or molecule

• Hydrogen follows a duet rule – two valence

electrons give it the same electron configuration

as helium

• Most other atoms follow a octet rule – eight

valence electrons will give each atom the same

number of valence electrons as a noble gas

Lewis Dot Structures - Ionic• Metals lose electrons, nonmetals gain electrons

• Rules for LDS for ionic compounds:

1. Write each element symbol

2. Determine the number of valence electrons

3. Add the valence electrons to each atom.

Clockwise- 12,3,6,9 one at a time.

4. Show the electron transfer from metal(s) to

nonmetal(s)

Lewis Dot Structures-Practice• Draw the Lewis Dot Structure for the

following atoms:

• Ca

• F

• Se

• Al

• P

• Si

Lewis Dot Structures - Practice

• Draw the Lewis Dot Structures for the

following ionic compounds:

• Na + Cl

• Mg + Br

• Al + O

• B + F

Lewis Dot Structures -Covalent

• Two nonmetals share electrons to achieve a

lower energy

• Rules:

1-3 same as ionic

4. Circle electrons that will pair together

5. Rearrange the compound so shared electrons

are aligned correctly(between atoms).

Lewis Dot Structures - Practice

• Draw the LDS for the following covalent

compounds:

• CCl4

• NBr3

• Phosphorus + Iodine

• Silicon + Fluorine

Bond Strength

• Of the three covalent bonds:

• A triple bond is the strongest followed by a

double bond and then a single bond which

is the weakest of the three

• A triple bond has the highest bond energy,

then a double bond, followed by a single

bond

Bond Length

• Of the three covalent bonds:

• A triple bond is the shortest followed

by a double bond and then a single

bond which is the longest

• Why?

LDS- Covalent Compounds

• So far we have looked at simple

covalent compounds and how they

will share valence electrons

• Now we will look at more complex

covalent compounds and how to

determine the Lewis Dot Structure.

• Try drawing the Lewis Dot Structure for

SO2 using the rules for covalent

compounds that you have learned

• The actual structure of the compound is:

• Let’s take a look at how to draw LDS

when we are given the formula and the

compound is more complex

Rules for Complex Covalent LDS• 1. Determine the total number of valence

electrons in the compound

• 2. Begin by putting a single bond between each

atom (Choose appropriate middle atom if

necessary)

• 3. Fill in lone pair electrons to fulfill duet/octet

rule

• 4. Add a double bond (or triple bond) if

necessary to insure the duet/octet rules are

fulfilled and the total number of valence

electrons are correct.

Resonance• Resonance occurs when several equally correct Lewis

Dot structures can be assigned to compounds.

• Double arrows are used to show options for

compounds with resonance structures.

• Resonance structures for SO2:

Practice:

• Draw the LDS for the following

compounds with the new rules you

have been given:

HF N2 NH3 CH4

NF3 O2 CO PH3

LDS for Ions

• For ions the rules for drawing Lewis Dot Structures

are the same except the total number of valence

electrons will either increase or decrease

depending on the charge

• For a positive charge subtract a valence electron

• For a negative charge add a valance electron

• After Drawing the LDS brackets are added and the

charge is added outside the brackets- top right

Example: NH4+

• Add up valence electrons for each atom:

N-5 H-1 total = 9

• Because of the +1 charge we assume a

valence electron has been lost

• Our new total of valence electrons is 8

• Draw the LDS using the same rules.

Don’t forget the brackets and the charge

Practice

• Complete the LDS for the following

ions. Show resonance structures if

they exist

NO+ NO3- SO4

-2

ClO3-

PO4-3 SCN-

Lone Pair Electrons• The unshared valence electrons

represented in LDS are called lone

pair electrons.

• Example: In CF4 each fluorine has six

lone pair electrons and

carbon has zero for a total

of twenty four.

VSEPR

• Stands for Valence Shell Electron Pair

Repulsion

• This theory states that electrons pairs

around an atom will spread out as far as

possible

• This repulsion is due to the same charges

on electrons

Molecule Polarity• We have already discussed a polar

covalent bond in terms of dipole

moments caused by differences in

electronegativity.

• We will now use this knowledge to

determine whether a molecule is

polar or non polar

Molecule Polarity

• In order for a molecule to be considered

polar it needs to have a concentrated

partial positive charge on one end and a

concentrated partial negative charge on

the other.

• Molecules that have polar bonds will not

necessarily be polar molecules

Molecule Polarity

• Symmetrical molecules can have

dipole moments cancel each other

out causing them to be nonpolar.

• Examples CH4, BF3, CO2

Molecule Polarity

• A molecule with a lone pair of

electrons in place of a bond will

always be polar (bent and trigonal

pyramidal)

• Examples: NH3, H2O, NO2-

Molecule Polarity

• Molecules that have a tetrahedral,

trigonal planar, and linear geometry

can be either polar or nonpolar.

• The central atom would have to have

different atoms bonded to it to be a

polar molecule. Example: HCN, HCO2-