Chemical

Bonding I:

Basic

Concepts

Valence electrons are the outer shell electrons of an

atom. The valence electrons are the electrons that

particpate in chemical bonding.

1 1 ns1

2 2 ns2

13 3 ns2np1

14 4 ns2np2

13 5 ns2np3

16 6 ns2np4

17 7 ns2np5

Group # of valence e- e- configuration

Periodic table of elements

9.1

Lewis Dot Symbols for the Representative Elements &

Noble Gases

*Why do substances bond? *More stability

*Atoms want to achieve a lower energy state

*Between a metal and a non-metal (I, II group and 16, 17 group of periodic table, metal

hydride, between cation and acid residue)

*Metals lose electrons becoming a cations, while non-metals gain electrons becoming anions.

*An ionic bond is an electrostatic attraction between the oppositely charged ions.

Li + F Li+ F -

The Ionic Bond

1s22s1 1s22s22p5 1s2 1s22s22p6

[He] [Ne]

Li Li+ + e-

e- + F F -

F - Li+ + Li+ F -

Lattice energy (E) increases

as Q increases and/or as r decreases.

cmpd lattice energy

MgF2

MgO

LiF

LiCl

2957

3938

1036

853

Q= +2,-1

Q= +2,-2

r F- < r Cl-

Electrostatic (Lattice) Energy

E = k Q+Q-

r

Q+ is the charge on the cation

Q- is the charge on the anion

r is the distance between the ions

Lattice energy (E) is the energy required to completely separate

one mole of a solid ionic compound into gaseous ions.

Ionic Structures

*In an ionic compound (solid), the ions are packed together into a repeating array called a crystal lattice.

*The simplest arrangement is one in which the spheres in the base are packed side by side. Opposite charges are attracted to each other.

*Its called simple cubic packing (NaCl is an example)

http://www.avogadro.co.uk/structure/chemstruc/ionic/g-ionic.htm

*Depend on the forces between the particles *The stronger the bonding between the

particles, the higher the M.P and BP

*MP tends to depend on the existence of a regular lattice structure

*An impurity disrupts the regular lattice that its particle adopts in the solid state, so it weakens

the bonding.

*They always LOWER melting points

* Its often used to check purity of a known molecular covalent compound because its MP will

be off, proving its contamination

A 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 pairs lone pairs

lone pairs lone pairs

single covalent bond

single covalent bond

8e-

H H O + + O H H O H H or

2e- 2e-

Lewis structure of water

Double bond – two atoms share two pairs of electrons

single covalent bonds

O C O or O C O

8e- 8e- 8e- double bonds

Triple bond – two atoms share three pairs of electrons

N N

8e- 8e-

N N

triple bond triple bond

or

Lengths of Covalent Bonds

Bond Lengths

Triple bond < Double Bond < Single Bond

H F F H

Polar covalent bond or polar bond is a covalent

bond with greater electron density around one of the

two atoms (electrons are shared unequally)

electron rich

region electron poor

region e- rich e- poor

d+ d-

Electronegativity is the ability of an atom to attract

toward itself the electrons in a chemical bond.

Electronegativity - relative, F is highest

H F

electron poor

region

electron rich

region

The Electronegativities of Common Elements

Nonpolar Covalent

share e- equally

Polar Covalent

partial transfer of e-

Unequal sharing

Ionic

transfer e-

Increasing difference in electronegativity

Classification of bonds by difference in electronegativity

Difference Bond Type

0 Nonpolar Covalent

1.7 Ionic

0 < and

Classify the following bonds as ionic, polar covalent,

or covalent: The bond in CsCl; the bond in H2S; and

the NN bond in H2NNH2.

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 NonPolar Covalent

9.5

*A covalent bond that occurs between two atoms in which both electrons shared in the bond come from the same atom.

*Both electrons from the nitrogen are shared with the upper hydrogen

*Ammonium has 3 polar covalent bonds and 1 coordinate (dative) covalent bond.

Properties do not differ from those of a normal covalent bond.

*Hydronium (H3O+) *Carbon monoxide (CO)

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 number 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 number 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) = 4 1 double bond = 4

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

Exceptions to the Octet Rule

The Incomplete Octet

H H Be Be – 2e-

2H – 2x1e-

4e-

BeH2

BF3

B – 3e- 3F – 3x7e-

24e-

F B F

F

3 single bonds (3x2) = 6

9 lone pairs (9x2) = 18

Total = 24

9.9

Exceptions to the Octet Rule

Odd-Electron Molecules (radicals -very reactive)

N – 5e- O – 6e-

11e-

NO N O

Polarity

*The shape of the molecule directly influences the overall polarity of the molecule.

*If there is symmetry the charges cancel each other out, making the molecule non-polar

*If there is no symmetry, then its polar

http://www.mhhe.com/physsci/chemistry/chang7/esp/folder_structure/bo/m4/s2/index.htm

*Polar bonds do not guarantee a polar molecule

*Ex: CCl4 and CO2 both have polar bonds, but both are NON-POLAR molecules. They have a

dipole moment of zero

*The greater the dipole moment, the more polar the molecule

*Polar molecules have higher melting and boiling points (for example the BP of HF is

19.5° C, and the BP of F2 is –188° C).

*Polar solvents dissolve ionic and polar molecules more efficiently than non-polar

solvents

Dipole Moments and Polar Molecules

H F

electron rich

region electron poor

region

d+ d-

The bent shape creates an

overall positive end and negative end

of the molecule = POLAR The symetry of the molecule

Cancels out the “charges”

Making this NON-POLAR

No overall DIPOLE

* Linear:

*When two atoms attached to central atom are the same, the molecule will be Non-Polar (CO2)

*When the two atoms are different the dipoles will not cancel, and the molecule will be Polar

(HCN)

* Bent:

* The dipoles created from this molecule will not cancel creating a net dipole moment and the

molecule will be Polar (H2O)

* Pyramidal:

* The dipoles created from this molecule will not cancel creating a net dipole and the molecule

will be Polar (NH3)

* Trigonal Planar:

*When the three atoms attached to central atom are the same, the molecule will be Non-Polar

(BF3)

*When the three atoms are different the dipoles will not cancel, resulting in a net dipole, and the

molecule will be Polar (CH2O)

* When the four atoms attached to the central atom are the same the molecule will be Non-Polar

* When three atoms are the same, and one is different, the dipoles will not cancel, and the molecule will be Polar

Intermolecular Forces

Dipole-Dipole Forces

Attractive forces between polar molecules

Orientation of Polar Molecules in a Solid

Intermolecular Forces

Ion-Dipole Forces

Attractive forces between an ion and a polar molecule

Ion-Dipole Interaction

Intermolecular Forces

Hydrogen Bond

The hydrogen bond is a special dipole-dipole interaction

between the hydrogen atom in a polar N-H, O-H, or F-H bond

and an electronegative O, N, or F atom.

Why is the hydrogen bond considered a

“special” dipole-dipole interaction?

Types of Crystals Metallic Bonds- electron sea model - metal cations in a sea of valence

electrons

• the valence electrons do not “belong” to a single cation, but

are delocalized and may move about

• Good conductors of heat and electricity

Cross Section of a Metallic Crystal nucleus &

inner shell e-

mobile “sea”

of e-

Metallic bond

*Occurs between atoms with low

electronegativities

*Metal atoms pack close together in 3-

D, like oranges in a

box.

Conductivity

*Delocalised electrons are free to move so when a

potential difference is

applied they can carry the

current along

*Mobile electrons also mean they can transfer

heat well

*Their interaction with light makes them shiny

(lustre)

*Depend on the forces between the particles *The stronger the bonding between the

particles, the higher the M.P and BP

*It takes energy to break the bonds between the C12H22O11 molecules in sucrose crystal structure.

*It also takes energy to break the hydrogen bonds in water so that one of these sucrose molecules can fit into solution.

*In order for sugar to dissolve, there must be a greater release of energy when the dissolution occurs than when the breaking of bonds occur.

*The energy needed to break the ionic bond must be less than the energy that is released

when ions interact with water.

*The intermolecular ion-dipole force is stronger than the electrostatic ionic bond

*Breaks up the compound into its ions in solution.

*Soluble salt in water breaks up as NaCl (s) Na

+ (aq) + Cl

- (aq)