Chemical BondingChapter 6

Sections 1, 2, and 5

Chemical Bonds

A chemical bond is the mutual electrical attraction between the nuclei and valence electrons of different atoms that bind the atoms together

Noble gases tend not to do this because of their filled s and p orbitals.

They have a stable octet: outer s and p orbitals are completely filled with e-’s ( 8 total)

Chemical Bonds

Atoms that don’t have a stable octet are more reactive because their potential energy is higher. They become more stable by decreasing their potential energy.

Octet Rule: chemical compounds tend to form so that each atom has an octet of e-’s in its highest occupied energy level

How to do this? Gain , lose or share electrons between atoms

Chemical Bonds

By forming a chemical bonds, atoms gain stability!

Chemical changes always involve energy

What type of bonds can be formed?Ionic bondCovalent bond

Nonpolar covalentPolar covalent

Chemical Bonds

Ionic bonding: bonds that result from electrical attractions between cations and anions

Covalent bonding: sharing of electron pairs between 2 or more atoms

*** In reality, bonding is often somewhere between the two extremes***

Two types of Covalent Bonds

Nonpolar –covalent: equal sharing of electron pairs

Polar-covalent: unequal attraction for the shared electrons

How can we determine the type of bond?

Knowing how strong an atom’s ability is to attract electrons (aka electronegativity), helps us determine if it will form a ionic or covalent bond with another atom.

A large difference in E.N. between atom’s will result in an ionic bond

A small difference between atom’s will result in a form of covalent bonding

NonpolarCovalentshare e-

Polar Covalent

partial transfer of e-

Ionic

transfer e-

Increasing difference in electronegativity

What type of Bond is it?

Electronegativity Difference Bond Type

0 to 0.3 Nonpolar Covalent

0.4 to 1.7 Polar Covalent

1.7 Ionic

Do you see any trends?

A metal and nonmetal tend to form ionic compounds

Nonmetal and nonmetal tend to form polar-covalent or nonpolar- covalent compounds

H F FH

Polar covalent bond or polar bond :covalent bond with greater electron density around one of the two atoms

electron richregion

electron poorregion e- riche- poor

d+ d-

Classify the following bonds as ionic, polar covalent,or covalent:

Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic

H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent

N – 3.0 N – 3.0 3.0 – 3.0 = 0 NonpolarCovalent

CsCl

H2S

N2

Properties of Molecular Covalent

CompoundsNot very soluble in water

Do not conduct electricity

Low melting points

Low boiling points

Can be solids, liquids and gases at room temperature

Comparison of Ionic and Covalent Compounds

Types of Crystals

Lewis Structures

Review: What are valence electrons?

Lewis Dot Diagrams− an electron-configuration notation with

only the valence electrons of an element are shown, indicated by dots placed around the element’s symbol.

− the inner core electrons are not shown.

• Eight electrons in the valence shell (filling s and p orbitals) make an atom STABLE

s2p6

This is called the octet rule

The Octet Rule

• Bond formation follows the octet rule: Chemical compounds tend to form so that each atom:

by gaining, losing, or sharing electrons, has an octet of electrons in its valence energy level.

Lewis Structures for Compounds • The pair of dots between two symbols

represents the shared pair of a covalent bond.

F F

• Each fluorine atom is surrounded by three pairs of electrons that are not shared in bonds.

• An unshared pair, also called a lone pair, is a pair of electrons that is not involved in bonding and that belongs exclusively to one atom.

Lewis Structures

• The pair of dots representing a shared pair of electrons in a covalent bond is often replaced by a long dash.

H H

F F

covalent bond : is a chemical bond in which two or more electrons are shared by two atoms.

Why should two atoms share electrons?

F F+

7e- 7e-

F F

8e- 8e-

F F

F F

Lewis structure of F2

lone pairslone pairs

lone pairslone pairs

single covalent bond

single covalent bond

Multiple Covalent Bonds

• double covalent bond or double bond :covalent bond in which two pairs of electrons are shared between two atoms

• shown by two side-by-side pairs of dots or by two parallel dashes

HC

H HC

Hor

HC

HC

H

H

Multiple Covalent Bonds

• triple covalent bond or triple bond :covalent bond in which three pairs of electrons are shared between two atoms.

N N or N N

C C or C CH H H H

8e-

H HO+ + OH H O HHor

2e- 2e-

Single Bond – two atoms share one pair of electrons

Double Bond – two atoms share two pairs of electrons

single covalent bonds

O C O or O C O

8e- 8e-8e-double bonds double bonds

Triple Bond – two atoms share three pairs of electrons

N N8e-8e-

N N

triple bondtriple bond

or

Bond Type

Bond Length(pm)

C-C 154

CC 133

CC 120

C-N 143

CN 138

CN 116

Lengths of Covalent Bonds

Bond Lengths

Triple bond < Double Bond < Single Bond

1. Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center.

2. Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.

3. Complete an octet for all atoms except hydrogen

4. If structure contains too many electrons, form double and triple bonds on central atom as needed.

Writing Lewis Structures

Write the Lewis structure of nitrogen trifluoride (NF3).

Step 1 – N is less electronegative than F, put N in center

F N F

F

Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)

5 + (3 x 7) = 26 valence electrons

Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms.

Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

Write the Lewis structure of the carbonate ion (CO32-).

Step 1 – C is less electronegative than O, put C in center

O C O

O

Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) -2 charge – 2e-

4 + (3 x 6) + 2 = 24 valence electrons

Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms.

Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

Step 5 - Too many electrons, form double bond and re-check # of e-

2 single bonds (2x2) = 41 double bond = 4

8 lone pairs (8x2) = 16Total = 24

resonance structure: one of two or more Lewis structures for a single molecule can be drawn to represent a molecule

O O O+ -

OOO+-

O C O

O

- -O C O

O

-

-

OCO

O

-

-

What are the resonance structures of the carbonate (CO3

2-) ion?

Exceptions to the Octet Rule

The Incomplete Octet

H HBeBe – 2e-

2H – 2x1e-

4e-

BeH2

BF3

B – 3e-

3F – 3x7e-

24e-

F B F

F

3 single bonds (3x2) = 69 lone pairs (9x2) = 18

Total = 24

Exceptions to the Octet Rule

Odd-Electron Molecules

N – 5e-

O – 6e-

11e-

NO N O

The Expanded Octet (central atom with principal quantum number n > 2)

SF6

S – 6e-

6F – 42e-

48e-

S

F

F

F

FF

F

6 single bonds (6x2) = 1218 lone pairs (18x2) = 36

Total = 48

Molecular GeometryChapter 6.5

VSEPR THEORYLewis Dot Diagrams are 2D but we live in a 3D world.

How are molecules actually arranged??Follows the Valance Shell Electron Pair Repulsion Theory or VSEPR

AB2 – LinearNumber of

Surround AtomsNumber of Lone

Pairs Bond Angle

2 0 180˚

Cl ClBe

AB3 – Trigonal Planar

Number of Surround Atoms

Number of Lone Pairs Bond Angle

3 0 120˚

AB2E1 – BentNumber of

Surround AtomsNumber of Lone

Pairs Bond Angle

2 1 <120˚

AB4 – TetrahedralNumber of

Surround AtomsNumber of Lone

Pairs Bond Angle

4 0 109.5˚

AB3E1 – Trigonal Pyramidal

Number of Surround Atoms

Number of Lone Pairs Bond Angle

3 1 107˚

AB2E2 – BentNumber of

Surround AtomsNumber of Lone

Pairs Bond Angle

2 2 104.5˚

Predicting Molecular Geometry

1. Draw Lewis structure for molecule.

2. Count number of lone pairs on the central atom and number of atoms bonded to the central atom.

3. Use VSEPR to predict the geometry of the molecule.

What are the molecular geometries of SO2 and SF4?

SO O

AB2E

bent

S

F

F

F F

AB4E

distortedtetrahedron

H F

electron richregion

electron poorregion

+d -d

The electronegativity of an atom will create a dipole,

or polar molecule.

Which of the following molecules have a dipole moment?H2O, CO2, SO2, and CH4

O HH

dipole momentpolar molecule

SO

O

CO O

no dipole momentnonpolar molecule

dipole momentpolar molecule

C

H

H

HH

no dipole momentnonpolar molecule

Intermolecular forces: attractive forces between molecules.

Intramolecular forces: hold atoms together, attractive forces within a molecule.

Generally, intermolecular forces are much weaker than intramolecular forces.

Properties of Ionic Compounds

Combination of ions (cation/anion)

Hard and Brittle

Tightly packed solids in a crystal lattice

Usually soluble in water

Conducts electricity when dissolved

High melting points

Breaking Ionic Bonds

Ionic Bonds are very tightly boundpositive and negative attraction

A LOT of energy needs to be put in to break an ionic bond

How does this affect the melting point?