Chemical Bonding

Forming compounds from atoms

Intramolecular Interactions

• Intramolecular = inside the molecules.

– The bonds that form between the atoms.• Metallic• Ionic• Covalent – Non polar and polar

Octet rule• Atoms form bonds in order to have a full outer shell.

i.e. they have the same electron configuration as the noble gases

• One atom can give electrons away (cation) and another formally receive electrons (anion) to form Ionic bonds

• Or they can share the electrons to form Covalent bonds

• Or the metals can live in a ‘sea of electrons’ where the electrons are free to move about. This give rise to Metallic bonding

WARNING

• I know you’re clever, however this is a beginner chemistry paper.

• There will be exceptions to the broad rules that will be used here as you move on in chemistry.

• It is too time consuming at this level to cover each and every exception to the general rules.

• (I’m happy to talk to you about them, but the exam will be based on the generic, not the exceptions)

Bonding

• The periodic table can be split into three parts based on electronegativity (how much they want another electron). High values, Medium values and Low values.

• The difference in these values dictate what kind of bonding will be shared between two atoms.

Bonding• High electronegativity

Bonding• Medium electronegativity

Bonding• Low electronegativity

Bonding• When two metals in the blue section bond,

Metallic Bonding is the result

Examples:

• Na (s)• Pd (s)• Ni (s)• etc etc…

Bonding• When an element from the blue and one from

the red area bond, Ionic bonding results

Examples:

• Na and Cl• Zn and F• Ni and O• etc etc…

Covalent Bonding

• The next two discussed are when the electrons are shared in a two electron bond

• The electrons can either be shared evenly (non-polar covalent) where the two atoms have similar electronegativity or unevenly (polar covalent) where one atom is more electronegative than the other

Bonding• When two elements from the orange area

bond, Non-Polar Covalent results

Examples:

• C and H• C and P• S and H• etc etc…

Bonding• When one element from the orange area

bonds with one from the red, a Polar Covalent results

Electronegativity

H2.1

Li

1.0

Be1.5

B

2.0

C

2.5

N

3.1

O

3.5

F

4.0

Na1.0

Mg1.3

Al

1.5

Si

1.8

P

2.1

S

2.4

Cl

3.0

K

0.9

Ca1.1

Ga1.8

Ge2.0

As2.2

Se2.5

Br2.8

Rb0.9

Sr

1.0

In

1.5

Sn1.7

Sb1.8

Te2.0

I

2.2

Cs0.9

Ba0.9

Tl

1.5

Pb1.6

Bi

1.7

Po1.8

At2.0

Fr

0.9

Ra0.9

The reason these bonds are polar is that their electronegativity differs by > 0.5

e.g. Si – Br2.8 – 1.8 = 1.0 therefore polar

Examples:

• Cl and H• O and P• S and F• C and O• etc etc…

What are these?Metallic, Ionic, Covalent (polar or non-polar)

• 1. Mg and Ca

• 2. K and Cl

• 3. C and H

• 4. C and O

• 5. As and Br

• 6. O and O

• 7. Ti and Cr

• 8. Na and F

• 9. Si and I

• 10. Sb and Cl

What are these?Metallic, Ionic, Covalent (polar or non-polar)

• 1. Mg and CaMetallic

• 2. K and Cl Ionic

• 3. C and H Covalent non-polar

• 4. C and O Covalent polar

• 5. As and Br Covalent polar

• 6. O and O Covalent non-polar

• 7. Ti and Cr Metallic

• 8. Na and F Ionic

• 9. Si and I Covalent polar

• 10. Sb and Cl Covalent polar

Lewis dot diagrams

• To work out the bonding in a covalent molecule, Lewis dot diagrams can be used• Most elements obey the octet rule, where

they want 4 pairs of electrons• By following some basic rules, structures can

be worked out

Instructions:1. Draw each atom with its valence electrons represented by dots around the symbol.

2. Underneath the symbol note how many electrons each atom needs to share to achieve the stable octet. (Remember hydrogen is an exception to the rule). 3. Work out how to fit the atoms together so each has the appropriate number of electrons. A good rule of thumb is that

the atom that needs to share the most electrons is probably the central atom (the atom the others are bonded to). 4. Draw a bond diagram of the molecule including lone pairs of electrons on the central atom.

Example: Cl2

1. Draw each atom with its valence electrons represented by dots around the symbol.

• Chlorine had 7 valence electrons from the periodic table. Therefore 3 pairs and one remaining

Example: Cl2

2. Underneath the symbol note how many electrons each atom needs to share to achieve the stable octet.

Needs on more electron

3. Work out how to fit the atoms together so each has the appropriate number of electrons. A good rule of thumb is that the atom that needs to share the most electrons is probably the central atom (the atom the others are bonded to).

These could pair

Example: Cl2

4. Draw a bond diagram of the molecule including lone pairs of electrons on the central atom.

This has formed a single, 2-electron bond:

Cl - Cl

Example: CH4

1. Draw each atom with its valence electrons represented by dots around the symbol.

• Carbon had 4 valence electrons from the periodic table. Therefore 4 individual electrons and H has one

Example: CH4

2. Underneath the symbol note how many electrons each atom needs to share to achieve the stable octet.

C needs 4 x1 more electronsH needs 1 x4 more electrons

Example: CH4

• 3. Work out how to fit the atoms together so each has the appropriate number of electrons. A good rule of thumb is that the atom that needs to share the most electrons is probably the central atom (the atom the others are bonded to).

These could pair

Example: CH4

4. Draw a bond diagram of the molecule including lone pairs of electrons on the central atom.

This has formed a four single, 2-electron bonds:

More examples:• CO2

• CH2O

• SiF4

• O2

• NH3

VSEPR

• Valence shell electron pair repulsion.– Gives the shape of the molecule that can have

important consequences for some properties (Intermolecular interactions)

– Have to work out how many electron density areas (‘steric number’) and lone pairs there are using Lewis dot diagrams

Molecular Polarity• Polar bonds can lead to molecules being polar

overall.• Must be careful that the polarities do not

cancel you by geometry– e.g. Label as polar or non-polar by determining

geometry:• CO2

• CCl4

• H2O

• PCl3

Molecular Polarity• Polar bonds can lead to molecules being polar

overall.• Must be careful that the polarities do not

cancel you by geometry– e.g. Label as polar or non-polar by determining

geometry:• CO2 Linear, therefore cancel out and non-polar

• CCl4 Tetrahedral, therefore cancel out and non-polar

• H2O Bent, not cancelling therefore polar

• PCl3 Trigonal pyramidal, not cancelling therefore polar

More examples

• CF2O

• CS2

• NCl3

• HBr• CH2CH2

• SCl2

• CH2F2