1. UNIT 7: BONDING Why do elements form bonds?
How can we explain and draw ionic bonds?
How can we explain and draw covalent bonds?
What are metallic bonds and why are they good conductors?
What is the difference between bond polarity and molecule polarity?
How do we predict shapes of covalent molecules?
What are the different forces that hold molecules together?

2. Aim: What previous knowledge will help us understand the bonding unit?
1. Why do atoms become ions?
2.How do atoms become ions? 
3. How do metals form ions?
4. How do nonmetals form ions?
5. Conductivity:
6. Electricity:
7. Melting Point:
8. Boiling Point:

3. Aim # 1- Why do elements form bonds?
A chemical bond is
- the force of attraction between the atoms of a compound
- the proton of one atom is
attracted to the electron
of another atom

4. the combining atoms either lose, gain, or share electrons in order to complete their outer shells

5. Energy and Chemical Bonds
Endothermic- energy is absorbed, bonds break
Energy is consumed for the bond to break
Ex) AB + energy  A + B
Exothermic- energy is released, bonds form
Creating bonds creates stability
Ex) A + B  AB + energy
“BARF”

6. Types of bonds
They differ in the types of elements involved. Also, how the valence electrons are handled.

7. Types of Bonds Ionic Bond: Covalent Bond: electrons are shared
electrons are transferred from one atom to another
Covalent Bond: electrons are shared
Metallic
electrons are mobile
Polar Covalent:
Unequal sharing
Nonpolar Covalent: Equal sharing

8. Aim # 2: How can we explain and draw ionic bonds?
 An IONIC bond is
The Transfer of electrons
Attraction between oppositely charged ions
Bond between a metal and a nonmetal
For example when Na and Cl atoms
come together

9. When metal atoms react:
They lose electrons
They become + charged ions (cations)
They acquire a complete octet
Their radii decrease (become smaller)
“MELPS”

10. Ionic bond: When nonmetal atoms react:
They gain electrons
They become – charged ions (anions)
They acquire a complete octet
Their radii increase

11. Properties of Ionic solids
High MP and BP
Hard substances
Conduct electricity in the liquid phase and in solutions ONLY
Good solubility
Atoms have an electronegativity difference of 1.7 or greater

13.  DRAWING LEWIS DOT STRUCTURES FOR IONIC BONDS
There are several steps to follow in order to draw the Lewis dot structure for ion compounds.
Write the metal symbol with no dots in brackets (brackets are optional)
Place the charge at the top right of the bracket
Write the nonmetal symbol with 8 dots around it (except H!)
Draw brackets around the symbol and place the charge of the ion at the top right of the bracket

14. Drawing Lewis Dot Structures for Ionic Bonds
NaF

15. **The exception to the rule **
Polyatomic ions are composed of multiple atoms (table E)
They have both covalent and ionic bonds

16. Compound
Bond Type
Dot Structure KF
MgI2
BeS
AlBr3
BaCl2
SrI2

17. AIM# 3: How can we explain and draw covalent bonds?
These compounds are formed when two or more nonmetals share electrons
nonmetals have the ability to have single, double and triple bonds
Each bond contain 2 electrons

18. Types of Covalent bonding
Nonpolar Covalent Bond
Electrons are shared equally
Occurs between atoms of the same element
Little or no difference in electronegativity
Polar Covalent Bond
Electrons are shared unequally
One atom is pulling on electrons more strongly
Electronegativity difference is less than 1.7 but greater than 0

19. COVALENT BONDING (Nonpolar)
All Diatomic Molecules are covalent. They are
Br2, I2, N2, Cl2 ,H2, O2, F2,
N2 - Is the only diatomic molecule that has a triple bond at room temperature.

20. Coordinate Covalent Bonds
-A type of covalent bond in which BOTH electrons come from the same atom
-Exists in polyatomic ions
-Examples:
NH4+
H3O+

21. Properties exist in gas, liquid, or solid state
Good insulators
Poor conductors of electricity in any phase
Low melting points
Soft substances
Poor solubility

22. Network Solids
These are covalent compounds that are extremely hard and have very high melting and boiling points.
Strong bonds and strong intermolecular forces
Cannot conduct electricity
Ex. Diamonds, silicon dioxide (SiO2)

23.  CONSTRUCTING LEWIS DOT STRUCTURE FOR (SINGLE BONDS) COVALENT MOLECULAR COMPOUNDS
Determine valence electrons in total (add them up for each element in the compound)
Divide by 2 to determine the number of pairs of electrons in total for the compound
Place first pair between the two elements (use a dash – to represent the shared pair)
Place remaining pairs around each elements making sure not to violate the octet rule (Remember H can have a max of 2 electrons)

24. H2
Cl2
Br2
HCl

25.  CONSTRUCTING LEWIS DOT STRUCTURES FOR (SINGLE BONDS) COVALENT MOLECULAR COMPOUNDS (more than two elements involved)
Determine valence electrons in total (add them up for each element in the compound
Divide by 2 to determine the number of pairs of electrons in total for the compound
Determine the most electronegative element and place it in the middle
Place the other elements around it
Start placing pairs (as dash lines) between the central atom and the terminal atoms
Place remaining around each elements making sure not to violate the octet rule (Remember H can have a max of 2 electrons)

26. NH3
CH4
CCl4

27.  CONSTRUCTING LEWIS DOT STRUCTURES FOR MULTIPLE COVALENT MOLECULAR COMPOUNDS (more than two elements involved)
Determine valence electrons in total (add them up for each element in the compound
Divide by 2 to determine the number of pairs of electrons in total for the compound
Determine the most electronegative element and place it in the middle
Place the other elements around it
Start placing pairs (as dash lines) between the central atom and the terminal atoms
Place remaining around each elements making sure not to violate the octet rule (Remember H can have a max of 2 electrons)
If octet rule is not yet reached you can make additional pairs of electrons into double or triple bonds until octet rule is obeyed by all elements
*can only be done with CNOPS

28. CO2 O2 N2

29. AIM #4 : What are metallic bonds and why are they good conductors?
Metallic atoms lose their valence electrons easily (remember they have low ionization energy)
Arranged in fixed positions called a crystal lattice. Therefore their bonds are a force of attraction between the negatively charged electrons and positively charged nucleus. The electrons are not attached to a particular nucleus
mobile see of electrons (MSOME)

31. Properties of Metals Malleability High MP Good conductors in any phase
Insoluble
Luster

32. Aim # 5 what is the difference between bond polarity and molecule polarity
BOND POLARITY (1.7 Rule)
To determine the bond type, this is based on the difference in electronegativity
The higher the difference the more polar the bond is
Ionic bonds will have an electronegativity difference of greater then 1.7
Remember “BONDED”

33. Polar Covalent Bonds HCl H2O NH3
Have electronegativity differences of 0.4 and 1.7
HCl
H2O
NH3

34. Nonpolar covalent bond
Have electronegativity difference of 0 and 0.3
Cl2
Br2

35. Nonpolar covalent bond
Ionic bond
Covalent bonds

36. Molecular Polarity (remember “SNAP”)
- asymmetrical
- unequal sharing of electrons
- asymmetrical distribution of charge
- lone electrons on an atom
HCl H2O
NH3

37. Nonpolar Molecules - symmetrical - equal sharing of electrons
- 2 atoms are the same, no lone electrons
Cl CH4

38. Compounds with polar bonds and nonpolar shapes
CCl4

39. Aim # 6 How do we predict shapes of covalent molecules?
 MOLECULAR SHAPES
To determine the shapes of molecules we use Lewis Structures and the VSEPR theory
VSEPR: Valence shell electron pair repulsion will predict the shapes of molecules from electron pairs on the central atom
Linear – line, “HX”
HCl
CO2 N2

40. Bent (angular) – Hydrogen and Group 16
H2O H2Se
Pyramidal- going to have one side unoccupied (NX3, PX3)
NF3 PI3
Tetrahedral- CX4
CCl4 CH3Br

41. Aim # 7: What are the different forces that hold molecules together?
 INTERMOLECULAR FORCES (IMF)
The attractions are ONLY found in covalent compounds
IMF’s between molecules are not nearly as strong as the intramolecular attraction that hold compounds together
These are considered weak forces in comparison to bonds
They are, however, strong enough to control physical properties such as boiling and melting points and vapor pressures
There are four types of IMF’s

42. 1. London Dispersion Forces
The weakest of all of the IMF’s
Only important for the NONPOLAR molecules, such as the diatomics and noble gases.
This repulsion creates brief dipoles in atoms

43. 2. Dipole-Dipole
Stronger version of LDF
Dipole: are interactions between molecules that have a positive end and a negative end
- this is a result from the unequal sharing of electrons
(think dipole, polar)
These forces are important when the molecules are close to each other
The more polar the molecule, the higher is its boiling point

44. Dipole-Dipole Interactions
The more polar the molecule, the higher is its boiling point.
© 2009, Prentice-Hall, Inc.

45. 3. Hydrogen Bonds Not a BOND, but a force of attraction
The strongest type of IMF
Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen and fluorine
“FON”
These attractions are responsible for the high boiling point of water

46. 4. Molecule- Ion Attraction
How ionic compounds dissolve in water
Example: NaCl in H2O

47. Bond Type
Conductivity
BP/MP
Examples
Ionic
Only in liquids and aqueous states
(l) (aq)
HIGH!
Strong bonds and strong IMF
CuCl2
NaCl
Covalent
(Molecular)
None!
Low!
Weak Bonds and low IMF’s
CO2
H2O
Metallic
Solid and Liquid Phase Only!
(s) (l)
Strong Bonds and Strong IMF’s Al K Ca
*Network Solids
NONE!
Strong bonds and IMF’s
SiO2 and Diamond