1. UNIT 7: BONDING 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?
What are the different forces that hold molecules together?

2. Review of Prior Knowledge
ENDOTHERMIC:
EXOTHERMIC:
POTENTIAL ENERGY:
Why do atoms become ions?
How do atoms become ions? 
How do metals form ions?
How do nonmetals form ions?

3. Aim # 1- Why do elements form bonds?
A chemical bond is
- force of attraction between the atoms of a compound

4. the electrons of one atom are attracted to the protons of another
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
Exothermic- energy is released, bonds form
Remember: “BARF”

6. Bonding and Stability
When bonds are formed their products are more stable
The bonded elements of a compound become stable because their valence shell is full
Atoms become ions to become stable
Remember metals lose electrons to become positively charged
Nonmetals gain electrons and become negatively charged

7. Types of bonds There are three types of bonds Ionic Covalent Metallic
They differ in the elements involved and how their handle their valence electrons

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

9. Aim # 2: How can we explain and draw ionic bonds?
An ionic bond is
Transfer of electrons
Attraction between oppositely charged ions
Bond between a metal and a nonmetal

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

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

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

14.  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

15. Drawing Lewis Dot Structures for Ionic Bonds
NaF

16. Drawing of aluminum oxide is MgCl2

17. PRACTICE

19. AIM# 3: How can we explain and draw covalent bonds?
Occurs between nonmetals
Nonmetals hold on to their valence electrons (don’t want to give them up)
They still want a noble gas configuration
They get it this sharing electrons

20. 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

21. 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.

22. Properties of covalent compounds (molecular solids)
Soft
Good insulators
Low MP
Poor conductors of electricity in any phase
May conduct electricity in solution if they are polar

23. Network Solids
Covalent compounds that are extremely hard and have very high melting and boiling points.
Cannot conduct electricity
Ex. Diamonds, silicon dioxide (SiO2- quartz), and silicon carbide (SiC)

24.  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)

25. H2
Cl2
Br2
HCl

26.  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)

27. NH3
CH4
CCl4

28.  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

29. CO2 O2 N2

30. Polyatomic Ions Table E
Polyatomic ions are held together by covalent bonds but form ionic bonds with other ions.
***Any compound containing a polyatomic ion has both covalent and ionic bonds!!!!!!

31. AIM #4 : What are metallic bonds and why are they good conductors?
Bonds between metals that form as a result of ions immersed in a “sea” of mobile electrons

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

34. Conductivity
Bond Type
MP and BP
Hardness
Solid
Liquid
Aqueous
Metallic
High
Hard
Yes
Covalent
Low
Soft No
Ionic

35. < 1.7 but not “0”, then polar covalent
Aim # 5 what is the difference between bond polarity and molecule polarity
BOND POLARITY (1.7 Rule)
1.7 rule is applied primarily to binary compounds.
Determine the electronegativities of all atoms in the bond.
Take the difference between the bonded atoms
The higher the difference the more polar the bond is
If the difference is:
>1.7 then ionic
< 1.7 but not “0”, then polar covalent
=0 non polar covalent

37. Nonpolar covalent bond
Ionic bond
Covalent bonds

38. Molecular Polarity
Molecular polarity can be determined by the shape of the molecule and the distribution of charge
Molecular shapes
Linear (X2 HX CO2)
Bent (H2O)
Pyramidal (NH3)
Tetrahedral (CH4 CCl4)
A polar molecule is called a dipole
It has a positive side and a negative side

39. Nonpolar
Polar H2
HCl
Symmetry:

40. MOLECULE POLARITY: Nonpolar/polar shapes
SNAP : Symmetrical Nonpolar Asymmetrical Polar

41. Compound
Polar or Nonpolar
NH3
H20

42. How can a molecule be both polar and non polar?
CX4 tetrahedral shape. They are non polar.
The individual ligands or bonds are polar
Conclusion – nonpolar/polar.

43. Aim # 6 What are the different forces that hold molecules together?
Hydrogen Bonding
VERY STRONG forces of attraction between specific polar molecules
Bond between an atom of H from one molecule with
FON(fluorine, oxygen, nitrogen) in another molecule
Sometimes called dipole-dipole forces

44. Dipole-Dipole Interactions
Molecules that have permanent dipoles are attracted to each other.
The positive end of one is attracted to the negative end of the other and vice-versa.
These forces are only important when the molecules are close to each other.

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

46. Intermolecular Forces
The stronger the intermolecular forces, the higher the boiling points and melting points
Ionic Solids
Molecules with Hydrogen bonds
H must be bonded to N, O, or F
Polar molecules
Non polar molecules
For non polar molecules, the greater the mass, the greater the
force of attraction

47. Intermolecular Forces
They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.
© 2009, Prentice-Hall, Inc.

48. Intermolecular Forces
These intermolecular forces as a group are referred to as van der Waals forces.
© 2009, Prentice-Hall, Inc.

49. Ion-Dipole Interactions
Ion-dipole interactions are important in solutions of ions.
The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.

50. Dipole- Dipole Interactions
Molecules that have permanent dipoles are attracted to each other.
The positive end of one is attracted to the negative end of the other and vice-versa.
These forces are only important when the molecules are close to each other.
The more polar the molecule, the higher is its boiling point.

51. London Dispersion Forces
London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.
© 2009, Prentice-Hall, Inc.

52. London Dispersion Forces
These forces are present in all molecules, whether they are polar or nonpolar.
The tendency of an electron cloud to distort in this way is called polarizability.
© 2009, Prentice-Hall, Inc.

53. Hydrogen Bonds
Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine.
Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

54. Summarizing Intermolecular Forces