1. Ionic, Covalent, Metallic

2. Physical Properties of Types of Compounds
IONIC
COVALENT
METALLIC
Attractive/force strength
Strong
Weak
Varies
Melting/Boiling point
High
Low
Vapor pressure
Electrical conductivity
Only in liquid or aquous solution
Never!
Always!
Malleable/Brittle
Brittle
Malleable
Type of bond
M + NM
NM + NM M
Example
NaCl,. ZnSO4
H2O, NH3,
C6H12O6
Zn, Cu, Ca, Mg,
Na, Fe

3. Naming Ionic Compounds

4. Naming Ionic Compound Metal is listed first, followed by nonmetal.
Change the name of the nonmetal to -ide.
Examples: nitride, sulfide, fluoride, oxide, bromide, iodide, chloride, telluride, phosphide.

5. The 5 Steps for writing an ionic compound formula
Write the symbols of the two elements.
Write the charge of each as superscripts.
Drop the positive and negative signs.
Crisscross the superscripts so they become subscripts. (1 is not written)
Reduce when possible.

6. Formula for boron oxide
Write the symbols of the two elements.
Write the charge for each element.
Drop the positive & negative sign.
4. Crisscross the superscripts so they become subscripts. (1 is not written)
5. Reduce subscripts when possible. (not possible here) +3 3 B O -2 2 2 3

7. Al S 2 3 3 +3 -2 2 Formula for Aluminum sulfide
Write the symbols of the two elements.
Write the charge for each element.
Drop the positive & negative sign.
4. Crisscross the superscripts so they become subscripts. (1 is not written)
5. Reduce subscripts when possible. (not possible here) 3 +3 Al S -2 2 2 3

8. Ca CaO O 2 2 2 +2 2 -2 Formula for Calcium oxide
Write the symbols of the two elements.
Write the charge for each element.
Drop the positive & negative sign.
4. Crisscross the superscripts so they become subscripts. (1 is not written)
5. Reduce subscripts when possible. Ca
CaO 2 +2 2 -2 O 2 2

9. Examples of Reduction of Subscripts:

10. Transition Metals Most Transition metals have two valences.
Roman numerals are used in the name of the transition metal in the compound to show the valence on the cation.

11. Cu Cl 2 2 1 -1 +2 Formula for Copper (II) Chloride
Write the symbols of the two elements.
Write the charge for each element.
Drop the positive & negative sign.
4. Crisscross the superscripts so they become subscripts. (1 is not written)
5. Reduce subscripts when possible. (not possible here) Cu 2 Cl 1 -1 +2 2

12. The 5 Steps for writing an ionic compound formula with polyatomic ions
Write the symbols of the cation and anion.
Look up in the polyatomic chart if needed
Write the charge of each as superscripts.
Drop the positive and negative signs.
Crisscross the superscripts so they become subscripts. (1 is not written)
If more than one of the polyatomic ion is needed, use paranthesis to represent as a group
Reduce when possible.

13. Mg MgCO3 CO3 ( ) 2 2 2 2 +2 -2 Formula for Magnesium carbonate
Write the symbols of the cation and anion.
Look up in the polyatomic chart if needed
Write the charge of each as superscripts.
Drop the positive and negative signs.
Crisscross the superscripts so they become subscripts. (1 is not written)
If more than one of the polyatomic ion is needed, use paranthesis to represent as a group
Reduce when possible. Mg
MgCO3 2
CO3 2 +2 -2
( ) 2 2

14. Formula for Zinc hydroxide
Write the symbols of the cation and anion.
Look up in the polyatomic chart if needed
Write the charge of each as superscripts.
Drop the positive and negative signs.
Crisscross the superscripts so they become subscripts. (1 is not written)
If more than one of the polyatomic ion is needed, use paranthesis to represent as a group
Reduce when possible. Zn
Zn(OH)2 2
OH3 +2 -1 1
( ) 2

15. Al AlPO3 PO4 ( ) 3 3 +3 3 3 -3 Formula for Aluminum Phosphate
Write the symbols of the cation and anion.
Look up in the polyatomic chart if needed
Write the charge of each as superscripts.
Drop the positive and negative signs.
Crisscross the superscripts so they become subscripts. (1 is not written)
If more than one of the polyatomic ion is needed, use paranthesis to represent as a group
Reduce when possible. +3 3 Al
AlPO3
PO4 3 -3
( ) 3 3

16. Steps to Ionic Formula to Naming
Identify the cation:
The cation is always written first in the formula for an ionic compound.
The only common polyatomic ion that you will encounter is the ammonium ion (NH4+).
Identify the anion:
Cover up the cation. Everything that is leftover will be the anion.
Write the name of the ionic compound by writing the name of the cation followed by the name of the anion.
If polyatomic, use the chart to find its charge
If transition metal, calculate the charge by balancing it out with anion’s charge

17. Write the correct name for (NH4)2CO3.
Identify the cation:
The cation will always be the first ion written.
2. Identify the anion:
The anion will be everything leftover once the cation has been identified: (NH4)2CO3
carbonate ion
3. The correct name for this compound is ammonium carbonate.
CO3
CO3
(NH4)2
(NH4)2
AMMONIUM CARBONATE

18. Tin is a transition metal, needs special charge written with the name
Write the name for SnO2.
Identify the cation:
The cation will always be the first ion written.
2. Identify the anion:
The anion will be everything leftover once the cation has been identified
3. If transition metal, calculate the charge by balancing it out with anion’s charge
The charges of the cation and the anions must exactly balance out.
Tin is a transition metal, needs special charge written with the name -4 +4 Sn Sn O2 O2 O2
Tin ?
Oxygen -2
Oxygen -2
Tin
Oxide
Tin (IV) Oxide

19. Write the name for Mn3(PO4)2.
Identify the cation:
The cation will always be the first ion written.
2. Identify the anion:
The anion will be everything leftover once the cation has been identified
3. If transition metal, calculate the charge by balancing it out with anion’s charge
The charges of the cation and the anions must exactly balance out. +2
Tin is a transition metal, needs special charge written with the name -6 +6
Manganese ?
Phosphate -3
Mn3
Mn3
(PO4)2
(PO4)2
(PO4)2 +2
Manganese ?
Phosphate -3 +2
Manganese ?
Manganese
Phosphate
Manganese (II) Phosphate

20. Write the name for Cu (OH)3.
Identify the cation:
The cation will always be the first ion written.
2. Identify the anion:
The anion will be everything leftover once the cation has been identified:
3. If transition metal, calculate the charge by balancing it out with anion’s charge
The charges of the cation and the anions must exactly balance out.
Hydroxide -1 -3 +3
Copper is a transition metal, needs special charge written with the name
Hydroxide -1 Cu Cu
(OH)3
(OH)3
(OH)3
Cupper ?
Hydroxide -1
Copper
Hydroxide
Copper (III) Hydroxide

21. Checking for understanding
Name these ionic compounds: KI
CaBr₂
Na₃N
Ca(OH)₂
Write formulas:
Barium oxide
Calcium phosphate
Sodium hydroxide
Magnesium sulfide

22. Click Below for the Video Lectures
Ionic Solids Ionic Bonding

23. Naming Covalent Molecules

24. Covalent Molecule Naming Rules
Compounds between two nonmetals
1. First element in the formula is named first.
Keeps its element name
Gets a prefix if there is a subscript on it
2. Second element is named second
Use the root of the element name plus the -ide suffix
Always use a prefix on the second element

25. List of Prefixes
1 = mon(o) 2 = di 3 = tri 4 = tetra 5 = penta
6 = hexa
7 = hepta
8 = octa
9 = nona
10 = deka

26. P2 P2 P2 O5 O5 O5 Naming P2O5 Diphosphorus Phosphorus pentoxide Oxide
1. First element in the formula is named first.
Keeps its element name
Gets a prefix if there is a subscript on it
2. Second element is named second
Use the root of the element name plus the -ide suffix
Always use a prefix on the second element P2 P2 P2 O5 O5 O5
Diphosphorus
Phosphorus
pentoxide
Oxide

27. Naming Covalent Compounds
diphosphorus pentoxide
CO2 =
carbon dioxide
CO =
carbon monoxide
N2O =
dinitrogen monoxide
9/16/2018

28. Checking for understanding
Name these covalent compounds: CO
PBr₅
H₂O
SeF₆
Write formulas:
Carbon Tetrachloride
Ammonium Sulfate
Dinitrogen Trifluoride
Pentaphosphorus Heptoxide

29. Click Below for the Video Lectures
Covalent Bonding Covalent Network Solids Molecular Solids

30. Drawing Covalent Molecules

31. The Octet Rule: The Diatomic Fluorine Molecule1s 2s 2p
Each has seven valence electrons F 1s 2s 2p F F

32. The Octet Rule: The Diatomic Oxygen Molecule1s 2s 2p
Each has six valence electrons O 1s 2s 2p O O

33. The Octet Rule: The Diatomic Nitrogen Molecule1s 2s 2p
Each has five valence electrons N 1s 2s 2p N N

34. Drawing Covalent Molecules
3. CH3NH2
(14 e- )
2. NH3 H
1. HBr
(8 e- )
(8 e- ) H H N H N H C Br H H H H
Needs 2 e-
Has 2 e-
1. Count the total number of valence electrons
2. Place the atom that makes most bonds in the middle. (least electronegative other than hydrogen) H H H C N
3. Draw single bonds to the other atoms off of the central atom. H H
4. Place electrons around peripheral atoms to fill octet, then to central atom
5. if electrons run out, start making double, triple bonds

35. 1. Count the total number of valence electrons
2. Place the atom that makes most bonds in the middle. (least electronegative other than hydrogen)
3. Draw single bonds to the other atoms off of the central atom.
4. Place electrons around peripheral atoms to fill octet, then to central atom
5. if electrons run out, start making double, triple bonds
6. CO2
(16 e- )
4. CH4
(8 e- )
5. O3
(18 e- ) H O O O O C O H C H
2 electrons deficient
4 electrons deficient H
Use these two to make a double bond O C O O O O
9/16/2018

36. Molecular Force

37. Nonpolar molecules
form LONDON DISPERSION FORCE attractions. Since there are no permanent positive or negative ends, these attractions are extremely weak.
The attractions are a combination of temporary poles due to electron movement around the molecule or, in the case of huge molecules, they actually get tangled up with each other like sticky strands of spaghetti or yarn.
London dispersion forces generally get stronger as the size of the molecule increases.

39. Polar molecules
form DIPOLE attraction: the simple attraction of the oppositely charged ends of two molecules.
The partially positive end of one molecule attracts to the partially negative end of a different molecule.
This attraction allows these substances to exist as solids and liquids at higher temperatures than are possible for nonpolar molecules of equivalent size.

40. H - Cl H - Cl H - Cl H - Cl δ - δ + δ + δ - δ + δ - 3.2 2.1 3.2 2.1

41. Hydrogen Bond Some polar molecules form HYDROGEN BONDS between them.
Bonding between hydrogen and more electronegative neighboring atoms; Nitrogen, Oxygen and Flourine (F,O,N)
An extremely high melting point, boiling point, heat of fusion and heat of vaporization for a molecule its size

42. Hydrogen Bonding in Water
The attraction between the two molecule is nearly that of ionic attraction.
Not only that but the H end of one molecule actually forms temperary covalent bonds with the N,O or F that makes up the end of the other molecule.
This gives extra strength that allows water to be liquid at room teperature despite its small size.

43. Hydrogen bonding in water

44. So, what should you be able to do now?
1) Identify whether a compound is molecular, ionic, metallic or network based on its properties 2) Draw dot diagrams of simple molecules 3) Draw structural formulas of simple molecules 4) Determine the shape of simple molecules 5) Determine if simple molecules are polar or nonpolar 6) 7) Determine the attractive force type that attracts specific simple molecules to each other.

45. Checking for understanding
Explain the difference between polar and nonpolar molecules
Explain London dispersion force
Explain dipole attraction forces including hydrogen bond

46. Click Below for the Video Lectures
Intermolecular forces Dipole Forces London Disperson forces

47. Chemical Formula

48. Click Below for the Video Lectures
Molecule and Elements

49. Interpreting chemical formula
The subscript tells you how many moles of that particular ion you have in one mole of the compound.
A Mole is to chemistry what a “dozen” is to donuts or eggs. Chemical formulas and reactions are written by the mole. 1 mole = 6.02 X 1023 of anything.
If there is no subscript, then the number of moles of that ion is 1.
When Interpreting Formulas, if there are (parentheses) around the element, any subscript outside the parentheses multiplies all of the elements inside the parentheses by that amount.

50. Ca(NO3)2 Ca(NO3)2 1 mol Ca , 2 mol N, 6 mol O 6.02x1023 atoms of Ca
2 outside the parentheses, so double the number of atoms inside to get the total number of atoms of that element (2 N’s and 6 O’s).
Since Ca is not inside the parentheses, it is not doubled. There is only 1 Ca in the formula.
1 mol Ca , 2 mol N, 6 mol O
6.02x1023 atoms of Ca
2 x (6.02x1023 ) atoms of N
6 x (6.02x1023 ) atoms of O

51. Checking for understanding
Complete the chart bellow, an example was done for you
BaCO3
1mol Ba
1mol C
3mol O
Na3PO4
Al2(SO4)3

52. Molar Mass

53. Molar Mass The molar mass
is the mass of one mole of an element or compound.
Expressed in g/mol

54. Molar Mass from Periodic Table
1 mole Ag mole C mole S
= g = g = g

55. Molar Mass of an Element
To calculate molar mass use the atomic masses of all elements present in a compound
Molar Mass of CO2
= 1 x g/mol
44.01 g/mol
+ 2 x g/mol =

56. Molar Mass of a Compound
The molar mass of a compound is the sum of the molar masses of the elements in the formula.
Example: Calculate the molar mass of CaCl2.
Element
Number of Moles
Atomic Mass
Total Mass Ca 1
40.1 g/mole
40.1 g Cl 2
35.5 g/mole
71.0 g
CaCl2
111.1 g

57. Molar Mass of K3PO4 Calculate the molar mass of K3PO4. Element
Number of Moles
Atomic Mass
Total Mass in K3PO4 K 3
39.1 g/mole
117.3 g P 1
31.0 g/mole
31.0 g O 4
16.0 g/mole
64.0 g
K3PO4
212.3 g

58. Mole Conversion

59. Mole A mole is a unit of measurement and it represents
6.02 x of anything
In chemistry, it usually represents 6.02 x of particles, atoms or molecules
A mole is also called an Avogadro’s number
A mole is used to convert between particles/atoms/molecules and grams
1 mole = 6.02 x particles, atoms, molecules

60. Mole Map moles grams atoms Circle the numbers
Molar
Mass
atoms
6.02x1023
Circle the numbers
Underline what you are looking for
Start with what you circled

61. Mass to Mole Conversion
How many moles of carbon is
26 g of carbon?
26 g C
1 mol C
12.01 g C
= 2.2 mol C
Molar
Mass
moles
grams

62. How many moles is 5.69 g of Na? 5.69 g Na 1 mol Na 22.99 g Na moles
Molar
Mass
moles
grams 62

63. How many grams are in 9.45 mol of nitrogen atoms?
9.45 mol N
14.01 g N
1 mol N =
=132 g N
Molar
Mass
moles
grams 63 63

64. Remember… 1 Mole = 6.02 x 1023 atoms or molecules

65. Molar Conversion Examples
How many atoms are in 2.50 moles of C?
6.02  1023 atoms
1 mol
2.50 mol
1.51  1024
=atoms C
6.02x1023
atoms
moles

66. Calculate the number of atoms in 0.500 mol of Al.
3.01  1023
=atoms Al
6.02x1023
atoms
moles

67. Calculate the number of moles of 1.80 x 1024 Na atoms.
1.80 x 1024 atoms
= 2.99 moles
6.02x1023
atoms
moles

68. dinitrogen trioxide, N2O3 ?
How many grams are in
1.20 x1024 molecules of
dinitrogen trioxide, N2O3 ?
1.20  1024
molecules
1 mol
6.02  1023
molecules
76.02 g
1 mol
= g
=152 g N2O3
Molar
Mass
6.02x1023
atoms
moles
grams 68 68 68 68

69. Find the mass of 2.1  1024 molecules of NaHCO3. 2.1  1024 molecules
6.02  1023
molecules
84.01 g
1 mol
= 290 g NaHCO3
Molar
Mass
6.02x1023
atoms
moles
grams 69 69 69 69 69

70. Checking for understanding
1. How many atoms are in 5g of Ca?
2. How many atoms are in 1.2 mol K?
3. Calculate the mass of 1.4x molecules
of CaCO3

71. Click Below for the Video Lectures
The Mole