1. Chemical Bonding

2. Chemical Bond
The forces that hold groups of atoms together and make them function as a unit
Bonding involves only the valence electrons
There are 2 types of bonds:
Ionic: Transfer of electrons from a metal and to a nonmetal
Covalent: Sharing of electrons between 2 nonmetals
Note: When 2 metals bond an alloy is formed
Electrons are transferred or shared to give each atom a noble gas configuration (stable octet)
This is known as the octet rule

3. Lewis Diagrams
Valence electrons involved in bonding can be represented by Lewis dot diagrams
A chemical symbol represents the nucleus and the core electrons (not involved in bonding).
Dots around the symbol represent valence electrons.

4. Drawing Lewis DiagramsCl
Write the element symbol.
Draw dots, one for each valence electron
Dots should be spread over 4 sides
It does not matter what side the dots are placed, but do not start to pair dots until there is one on each side
The number of valence electrons is equal to the group number. With one exception.

5. Lewis diagrams for the first 20 elements

6. Ionic Bonding

7. Ionic Bonding Metals Nonmetals Electron donors
Donate their valence electrons to become a positive ion (cation)
Nonmetals
Electron acceptors
Accept valence electrons to become a negative ion (anion)

8. Ionic Bonding

9. Ionic Bonding
The two oppositely charged ions are attracted to each other by a force called an ionic bond

10. Properties of Ionic Compounds
Structure:
Crystalline solids
Melting point:
Generally high
Boiling Point:
Electrical Conductivity:
Excellent conductors, molten and aqueous
Solubility in water:
Generally soluble

11. NaCl Crystal Lattice Ionic compounds form solids at SATP.
Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions.
All lattices are arranged so that each ion has the greatest possible number of oppositely charged ions close by, while keeping similarly charged ions as far away as possible

12. Representing Ionic Compounds Lewis Diagrams
Formation of sodium chloride: Cl
· ·
Na+ [ ] Cl
· · Na +

13. Lewis Structures for Ionic CompoundsBa • O •• •• O Ba 2+ 2-
Ba and O
BaO Mg • Cl •• •• Cl Mg 2+ - 2
Mg and Cl
Binary ionic compounds.
Note the types of arrows used to move electrons – fishhooks for single e-.
Write the Lewis symbol for each atom Determine how many e- each atom must gain or lose.
Use multiples of one or both ions to balance the number of electrons.
MgCl2

14. Representing Ionic Compounds Criss-Cross Method
For monatomic ions:
Take the absolute value of the ionic charge for the cation and make it the subscript for the anion and vice versa.
Example: Al3+ and Cl-
The 3 becomes the subscript for the chloride ion and the 1 becomes understood for aluminum.
Forming aluminum chloride: AlCl3

15. Representing Ionic Compounds Criss-Cross Method
For polyatomic ions:
Additional step of including brackets around the polyatomic ion if it has a subscript other than one.
Example: Mg2+ and OH-
The 2 becomes the subscript for the hydroxide ion, but brackets are needed to indicate 2 of each the O and the H.
The 1 becomes the understood subscript for Mg.
Forming magnesium hydroxide: Mg(OH)2

16. Polyatomic Ions NICK the CAMEL ate a CLAM for SUPPER in PHOENIX
Underlined letter represents the symbol of the element.
The consonants represent the number of oxygen
The vowels represent the negative charge.
Eg. Underlined letter= N
Number of consanants= 3 represents oxygens
Number of vowels= 1 represents charge
NO3- Nitrate

17. Covalent Bonding Formation of hydrogen chloride: Cl ® Cl H H +
Lone pairs, valence electrons not involved in covalent bond
Formation of hydrogen chloride:
· · Cl
· · Cl
· · H H +
® H - Cl
· ·
· ·
Covalent bond, shared electrons
Structural Formula: H-Cl (lone pairs are not drawn)

18. + Lewis Structures H H2: ® H H or H H Cl + ® Cl Cl Cl2: or Cl Cl+
H2:
® H H
or H H Cl
· · +
® Cl Cl
· ·
Cl2:
or Cl Cl
· ·
Structural Formula: Cl-Cl

19. Double and Triple Bonds
Atoms can share 4 electrons to form a double bond or 6 electrons to form a triple bond. =
O O
· ·
O2:
· ·
N N 
N2:
The number of shared electron pairs (covalent bonds) that an atom can form is the
bonding capacity.

20. Multiple Covalent Bonds• • N • •• •• N • • N •• • • N • •• N ••

21. Multiple Covalent Bonds• • • C O • ••
• • O • C • • • O
• •
• • •
• • • C O •• C O ••

22. Drawing Lewis Structures
Arrange the element symbols.
Central atoms are generally those with the highest bonding capacity.
Carbon atoms are always central atoms
Hydrogen atoms are always peripheral atoms
Add up the number of valence electrons from all atoms.
Add one electron for each negative charge and subtract one for each positive charge.
Draw a skeleton structure with atoms attached by single bonds.
Complete the octets of peripheral atoms.
Place extra electrons on the central atom.
If the central atom doesn’t have an octet, try forming multiple bonds by moving lone pairs.

23. Structural Formula
From the Lewis structure, remove dots representing lone pairs
Replace bond dots with a dash
H can only accommodate two electrons
H and O are common exceptions to rule 2
Organic compounds are not compact nor symmetrical.

24. or H F HF: H F H O H or H O H H2O: H N H H or H N H H NH3: H C H H
Draw Lewis structures and the structural formula for:
· ·
or H F
· ·
HF:
H F
· ·
· ·
· ·
H O H
· ·
or H O H
· ·
H2O:
H N H H
· ·
or H N H H
· ·
NH3:
H C H H
· ·
or H C H H
CH4:

25. Coordinate Covalent Bonds
A covalent bond in which both of the shared electrons come from the same atom.
E.g. NH3 (ammonia) and H+ (hydrogen ion) to form NH4 (ammonium)

26. Drawing Lewis Structures
· ·
Cl C Cl O
· ·
· ·
COCl2
24 ve’s
· ·
· ·
· ·
· ·
· ·
· ·
· ·
· ·
HOCl
14 ve’s
H O Cl
· ·
· ·
· · 
· ·
O Cl O O
· ·
· ·
· ·
· ·
ClO3
26 ve’s
· ·
· ·
· ·
· ·
· ·
H C O H H
· ·
CH3OH
14 ve’s
· ·