1. Bonding
By Mr. M

2. Types of Bonds
There are bonds between compounds: intermolecular forces
London Dispersion
Van Der Waals
Dipole-dipole
Hydrogen bonds
There are bonds within a single compound: intramolecular forces
Metallic
Ionic
covalent

3. What you will have to know
You must be able to identify a chemical as a metallic bond, ionic bond, or covalent bond
You must be able to describe chemicals that have metallic, ionic, or covalent bonds
You must be able to name, write formulas for, draw diagrams of, and draw lewis structures of ionic and covalent chemicals

4. How to tell which is which
Metallic = metal bonded to a metal
Ionic = metal bonded to a nonmetal
Covalent = nonmetal bonded to a nonmetal
You can also use electronegativity to tell the difference…

5. C. Bond Polarity
Most bonds are a blend of ionic and covalent characteristics.

6. C. Bond Polarity Nonpolar Covalent Bond e- are shared equally
symmetrical e- density
usually identical atoms

7. + - C. Bond Polarity Polar Covalent Bond e- are shared unequally
asymmetrical e- density
results in partial charges (dipole) + -

8. C. Bond Polarity
Nonpolar
Polar
Ionic

9. C. Bond Polarity Electronegativity
Attraction an atom has for a shared pair of electrons.
higher e-neg atom  -
lower e-neg atom +
This gets more complex, so we will come back to polarity later on

10. IONIC COVALENT METALLIC
e- are transferred from metal to nonmetal
e- are shared between two nonmetals
e- are delocalized among metal atoms
Bond Formation
Type of Structure
crystal lattice
true molecules
“”electron sea”
liquid or gas
solid
Physical
State
solid
very high
Melting
Point
high
low
Solubility in
Water
yes
usually not no
yes (solution or liquid)
yes (any form)
Electrical Conductivity no
Other
Properties
malleable, ductile, lustrous

11. Metallic Bonds
Metallic bonds are used to make alloys such as brass, bronze, or steal
Electrons in metals can move to different atoms (electron sea)

12. Ionic Bonds Used to make salts
Made up of an anion (negative charge) and a cation (positive charge)
Electrons are stolen

13. Lattice Structures Crystal shapes made from an ion cage
Water can get stuck between the ions, making a hydrate

14. Ionic Nomenclature Binary nomenclature Polyatomic nomenclature
Hydrate nomenclature

15. Binary nomenclature
Use only two types of elements but it can be in different amounts
Step 1: name the first element (the cation)
Step 2: name the second (the anion) but change the ending to “ide”
NaCl = sodium chloride
CaBr2= calcium bromide

16. Anion Names Chlorine -> chloride Fluorine -> fluoride
Bromine -> bromide
Iodine -> iodide
Nitrogen -> nitride
Oxygen ->oxide
Sulfur -> sulfide
Phosphorus ->phosphide

17. Turning the name into a formula
You first need to find the symbol for the elements
Then you find the charges
Use your ion sheet to look for roman numerals
Use I, II, III, IV, V, VI, VII, VIII for roman numerals
Then you must balance the charges

18. Charges 1+ 2+ 3+ NA 3- 2- 1-

19. Balancing Charges Magnesium Chloride
MgCl2
Subscripts tell you the number of atoms
You need 2 Cl – because Mg has charge of + 2
Sum of charges must equal zero (neutral)
= 0
Charge of one chlorine
Charge of Mg
Charge of second chlorine

20. Mg +2 + O -2 -> Magnesium Oxide = MgO
Na + + O -2 -> Sodium oxide = Na2O
K + + N -3 -> potassium nitride = K3N
Al +3 + O -2 -> aluminum oxide = Al2O3
There is two aluminum atoms in this compound
There is three oxygen atoms in this compound
Charges add up to zero
= = 0

21. Roman numerals
Transition metals can have different charges, iron can be + 2 or + 3
FeCl2 is iron II chloride
FeCl3 is iron III chloride
If you have lead IV oxide, you are combing Pb+4 with O-2
This makes PbO2 because it takes two O to cancel the Pb charge

22. Poly Atomic Nomenclature
Every ionic compound is two items and normally the first item is the first element only
The exception is NH4+ which is ammonium
The rest of the compound is an anion on your ion sheet
The rules are the same as binary except you use parenthesis when adding subscripts to polyatomic ions

23. NH4Br = ammonium bromide
KNO3 = potassium nitrate
Ca(NO3)2 = calcium nitrate
Fe2(SO4)3 = iron III sulfate
FeSO3 = iron II sulfite
FePO4 = iron III phosphate

24. Hydrate Nomenclature
Hydrates are named the exact same way as other ionic compounds except they get a third word
The third word is the word “hyrdate” with a prefix to say how much water is attached
Mono = 1, di = 2, tri =3, tetra = 4, penta = 5, hexa = 6, hepta = 7, octa = 8, nona = 9, deca = 10

25. CaBr∙2H2O = calcium bromide dihydrate
SnF4 ∙ 5H2O = tin IV fluoride pentahydrate
Fe(HCO3)3 ∙ 3H2O= iron III bicarbonate trihydrate

26. Covalent Bonds The atoms share electrons and appear to overlap
Diatomics = H, N, O, Cl, Br, I, F

27. Covalent Nomenclature
PREFIX
mono-
di-
tri-
tetra-
penta-
hexa-
hepta-
octa-
nona-
deca-
NUMBER 1 2 3 4 5 6 7 8 9 10

28. First element gets a prefix if there is more than one
The second element always gets a prefix
Second element changes its ending to “ide”
No roman numerals

29. CCl4
N2O
SF6
carbon tetrachloride
dinitrogen monoxide
sulfur hexafluoride

30. arsenic trichloride
dinitrogen pentoxide
tetraphosphorus decoxide
AsCl3
N2O5
P4O10

31. Lewis Dot Diagrams
A method of drawing ionic and covalent compounds using dots to represent valence electrons
The total compound should have dots equal to the sum of valence electrons

32. Ionic

33. Covalent
Octet rule – most atoms will try to get eight valence electrons
There are exceptions for more or less than eight
The most common exception is H and He only needing 2

34. Process
Step 1: Write down the first nonhydrogen element, this is your central atom
Step 2: Add dots one per side at a time until all of the central atom’s dots have been added
Step 3: Spread the remaining atoms around the central atom
Step 4: Give the remaining atoms their dots starting with the dot needed to complete the bond
Step 5: Fill any holes by moving dots into the bond position to make double and triple bonds
Step 6: Negative charges add dots, positive charges remove dots

35. CF4 F
F C F

36. CO2
O C O

37. ClO4-1 O
O Cl O

38. Resonance Structures
SO3 O
O S O O
O S O O
O S O

39. Ozone

40. VSEPR
Lewis dots and quantum mechanics can be combined to make the valence shell electron pair repulsion theory
It describes the shape of molecules
The shapes are named based on number of bonds and number of lone pairs on the central atom (lone pairs are not involved in bonding)

41. Shape 1
2 total
2 bond
0 lone
BeH2
LINEAR
180°

42. Shape 1 alternative
CO2
2 total
2 bond
0 lone
LINEAR
180°

43. Shape 2
3 total
3 bond
0 lone
BF3
TRIGONAL PLANAR
120°

44. Shape 3
3 total
2 bond
1 lone
SO2
BENT
<120°

45. Shape 3 alternative
4 total
2 bond
2 lone
H2O
BENT
104.5°

46. Shape 4
4 total
4 bond
0 lone
CH4
TETRAHEDRAL
109.5°

47. Shape 5
4 total
3 bond
1 lone
NH3
TRIGONAL PYRAMIDAL
107°

48. Shape 6
5 total
5 bond
0 lone
PCl5
TRIGONAL BIPYRAMIDAL
120°/90°

49. Shape 7
6 total
6 bond
0 lone
SF6
OCTAHEDRAL
90°

50. Shape 8
F P F F
PF3
4 total
3 bond
1 lone
TRIGONAL PYRAMIDAL
107°

51. + - Shapes with Polarity H Cl
Dipole moment is the direction of the polar bond in a molecule.
Arrow points toward the more e-neg atom.
H Cl + -

52. Determining Molecular Polarity
Depends on:
dipole moments
molecular shape

53. B. Determining Molecular Polarity
Nonpolar Molecules
Dipole moments are symmetrical and cancel out.
BF3 F B

54. B. Determining Molecular Polarity
Polar Molecules
Dipole moments are asymmetrical and don’t cancel .
H2O H O
net
dipole
moment

55. B. Determining Molecular Polarity
Therefore, polar molecules have...
asymmetrical shape (lone pairs) or
asymmetrical atoms
CHCl3 H Cl
net
dipole
moment