1. Bonding

2. Why do substances bond? More stability
Atoms want to achieve a lower energy state

3. What are bonds?
Forces that hold groups of atoms together and make them function as a unit.
Four types of bonds
1. Ionic
2. Covalent
3. Polar Covalent
4. Metallic

4. What are Ionic Bonds?
Ionic bonds involve transfer of electrons from one atom to another.
This transfer causes ions to form.
The ions (with opposite charges) are then held together by electrostatic attraction.
Metal with a non-metal OR metal with a polyatomic anion

5. Why do some things form ionic bonds and some form covalent bonds?
Because some elements form ions more easily than other elements.
For example, when a potassium ion forms, energy is released. This is exothermic. (This means K likes being an ion better than it likes being a neutral atom!)

6. How to Read Chemical Formulas
Subscripts = # of atoms
Symbols = types of atoms
Ex: NaCl
Ex: CaCl2

7. Ex: H3PO4
three hydrogen atoms
one phosphorus atom
four oxygen atoms
Ex: Na3C6H5O7
3 sodium
6 carbon
5 hydrogen
7 oxygen

8. Recognizing Ions
You must learn to recognize ions in chemical formulas so that you can name the chemicals.
Monoatomic ions have only one atom
Ex: H+ or Ca2+
Polyatomic ions are two or more atoms covalently bonded. They act like a single ion
Ex: SO OH NH C2H3O2-1

9. Using the back of your periodic table, identify the ions in the following compounds:
HF Ca(NO3)2
KI Mg3(PO4)2
NaOH NH4OH
CaCl AgClO3
H2SO BaCr2O7

10. How do Ions Bond? Ions bond so that the compound has no net charge.
Positive charges are balanced by negative charges. (The charges algebraically add to zero)
Ex: NaCl
Ex: FeCl2

11. To write formulas for ionic compounds:
Write the cation (positive ion) first and the anion (negative ion) second.
Treat polyatomic ions as one atom.
If the charges of the two ions are opposite but equal, you’re done (the charges are balanced, the compound has a neutral charge and the formula is correct)

12. Ex: sodium bicarbonate is made from two ions Na+1 and HCO3-1
writing in proper order gives:
NaHCO3
The magnitudes of the charges are equal, therefore the formula is correct as is (no subscripts are required)

13. Ex: Write the formula for ammonium hydroxide
Ex: Write the formula for calcium sulfate.

14. If the anion and cation don’t have charges of the same magnitude,
Use the positive charge as the subscript of the negative ion
Use the negative charge as the subscript of the positive ion
Drop the negative and positive signs (subscripts are always positive)
This is called “crisscrossing the charges”.

15. Example: Write the formula for the compound made of chlorine and barium.

16. Special tips
NEVER write the number 1 as a subscript.
If an ion is polyatomic and you need to add a subscript after it, you MUST PUT PARENTHESES around the polyatomic before adding the subscript.
If you end up with two subscripts and one is a multiple of the other, you must “reduce” the subscripts on both elements.

17. Ex: Write the formula using iron(3+) and hydroxide.
Ex: Write the formula for aluminum sulfate.

18. Write the formula for sulfate and lead (+4)
Write the formula containing aluminum and oxygen

19. Electronegativity
The ability of an atom in a molecule to attract shared electrons to itself.
Linus Pauling

20. Table of Electronegativities

21. Nonpolar-Covalent bonds
Ionic Bonds
Electrons are transferred
Electronegativity difference is greater than 1.7
Polar-Covalent bonds
Electrons are unequally shared
Electronegativity difference between .3 and 1.7
Nonpolar-Covalent bonds
Electrons are equally shared
Electronegativity difference of 0 to 0.3

22. Keeping Track of Electrons
Atoms in the same column have the same outer electron configuration.
And the same number of valence electrons.
For 1A -8A, number of valence equals the
group number

23. Electron Dot diagrams Lewis Structures
A way of keeping track of valence electrons.
Write the element symbol.
Put one dot for each valence electron
Electrons don’t pair up until they have to X

24. Write the electron dot diagram for
1. Na
2. Mg
3. C
4. O
5. F
6. Ne
7. He

25. Stable Electron Configurations
All atoms react to achieve noble gas configuration.
Noble gases have 2 s and 6 p electrons.
8 valence electrons .
Also called the Octet Rule. Ar

26. Ionic Bonding Na Cl

27. Ionic Bonding
Na+
Cl-

28. Ionic Bonding
All the electrons must be accounted for! Ca P

29. Ionic Bonding Ca P

30. Ionic Bonding
Ca+2 P

31. Ionic Bonding
Ca+2 P Ca

32. Ionic Bonding
Ca+2
P-3 Ca

33. Ionic Bonding
Ca+2
P-3 Ca P

34. Ionic Bonding
Ca+2
P-3
Ca+2 P

35. Ionic Bonding Ca
Ca+2
P-3
Ca+2 P

36. Ionic Bonding Ca
Ca+2
P-3
Ca+2 P

37. Ionic Bonding
Ca+2
Ca+2
P-3
Ca+2
P-3

38. Ionic Bonding
Ca3P2
Formula Unit

39. Properties of Ionic Compounds
Crystalline structure.
A regular repeating arrangement of ions in the solid.
Ions are strongly bonded.
Structure is rigid.
High melting points- because of strong forces between ions.

40. Sodium Chloride Crystal Lattice
Ionic compounds form solids at ordinary temperatures.
Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions.

41. The Octet Rule
Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level.
Diatomic Fluorine

42. Formation of Water by the Octet Rule

43. Comments About the Octet Rule
2nd row elements C, N, O, F observe the octet rule.
2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive.
3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals.
When writing Lewis structures, satisfy octets first, then place extra electrons around elements having available d orbitals.

44. Completing a Lewis Structure -CH3Cl
Make carbon the central atom
Add up available valence electrons:
C = 4, H = (3)(1), Cl = 7 Total = 14
Join peripheral atoms
to the central atom
with electron pairs. H .. ..
Complete octets on
atoms other than
hydrogen with remaining
electrons H .. C .. Cl .. .. .. H

45. Multiple Covalent Bonds: Double bonds
Two pairs of shared electrons

46. Multiple Covalent Bonds: Triple bonds
Three pairs of shared electrons

47. VSEPR Model
(Valence Shell Electron Pair Repulsion)
The structure around a given atom is determined principally by minimizing electron pair repulsions.
(That means the electrons are as far apart as possible)

48. Predicting a VSEPR Structure
Draw Lewis structure.
Put pairs as far apart as possible.
Determine positions of atoms from the way electron pairs are shared.
Determine the name of molecular structure from positions of the atoms.

49. Table – VSEPR Structures

50. VSEPR and the water molecule

51. Ammonia has a pyramidal shape

52. Methane
Has a tetrahedral shape

53. Relative magnitudes of forces
The types of bonding forces vary in their strength as measured by average bond energy.
Strongest
Weakest
Covalent bonds (400 kcal)
Hydrogen bonding (12-16 kcal )
Dipole-dipole interactions (2-0.5 kcal)
London forces (less than 1 kcal)

54. Hydrogen Bonding Bonding between
hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen
Hydrogen bonding in Kevlar, a strong polymer used in bullet-proof vests.

55. Hydrogen Bonding in Water

56. Hydrogen Bonding between Ammonia and Water

57. Dipole-Dipole Attractions
Attraction between oppositely charged regions of neighboring molecules.

58. The water dipole

59. The ammonia dipole

60. London Dispersion Forces
The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors.
Fritz London
London forces increase with the size of the molecules.

61. London Forces in Hydrocarbons