1. Ch 7.1 Forming Ions

2. Review… Cations are Groups 1A, 2A, and 3A
They have positive charges.
Anions are Groups 5A, 6A, and 7A
They have negative charges
They end in “ide”
Majority of elements in Groups 4A and 8A do not usually form ions.
These are called monatomic ions!

3. Ions of Transition Metals
Some transition metals in Groups 1B-8B form more than one cation with different charges.
Examples:
Iron forms Fe2+ and Fe3+
Two methods of naming:
Stock System: Iron (II) ion and Iron (III) ion
Classical name: Ferrous ion and Ferric ion
*We will use the stock system.

4. Common Metal Ions with More than One Ionic Charge
Symbol
Stock Name
Classical Name
Cu+
Copper(I) ion
Cuprous ion
Cu2+
Copper(II) ion
Cupric ion
Fe2+
Iron(II) ion
Ferrous ion
Fe3+
Iron(III) ion
Ferric ion
*Hg22+
Mercury(I) ion
Mercurous ion
Hg2+
Mercury(II) ion
Mercuric ion
Pb2+
Lead(II) ion
Plumbous ion
Pb4+
Lead(IV) ion
Plumbic ion

5. Common Metal Ions with More than One Ionic Charge
Symbol
Stock Name
Classical Name
Sn2+
Tin(II) ion
Stannous ion
Sn4+
Tin(IV) ion
Stannic ion
Cr2+
Chromium(II) ion
Chromous ion
Cr3+
Chromium(III) ion
Chromic ion
Mn2+
Manganese(II) ion
Manganous ion
Mn3+
Manganese(III) ion
Manganic ion
Co2+
Cobalt(II) ion
Cobaltous ion
Co3+
Cobalt(III) ion
Cobaltic ion

6. Other Transition Metal Ions
Some transition metals have only one charge and do not use a Roman numeral.
Examples (Write on your periodic table!)
Silver: Ag+
Cadmium: Cd2+
Zinc: Zn2+

7. Polyatomic Ions
The names of most polyatomic anions end in “-ite” or “-ate”
The “-ite” ending indicates one less oxygen atom than the “-ate” ending
Remember to use parentheses if more than one is needed.

8. Writing a Formula
Write the formula that will form between Ba and Cl Solution: 1. Write the positive ion of metal first, and then negative ion Ba2+ Cl 2. Do the charges equal zero? NO!! 3. Use Criss-Cross method – write subscripts Ba2+ Cl1 BaCl2 8

9. Learning Check
Write the correct formula for the compounds containing the following ions: 1. Na+, S2- 2. Al3+, Cl- 3. Mg2+, N3- 4. Al3+, S2- 9

10. Solution
1. Na+, S2- Na2S 2. Al3+, Cl- AlCl3 3. Mg2+, N3- Mg3N2 4. Al3+, S2- Al2S3 10

11. Binary Compound: is composed of two elements and can be either ionic or molecular.

12. Naming Binary Ionic Compounds
To name a binary ionic compound, place the cation name first, followed by the anion name.
Remember the anion ends in “-ide”
Examples:
Cs2O
NaBr
CuO
Cesium Oxide
Sodium Bromide
Copper(II) Oxide

13. Naming Compounds with Polyatomic Ions
State the cation first and then the anion just as you did in naming binary ionic compounds.
KNO3
Mg(ClO2)2
Potassium Nitrate
Magnesium Chlorite

14. Naming Binary Molecular Compounds
Binary molecular compound: must be composed of two nonmetals
Use prefixes to indicate the number and kind of atom in the compound
Use the following general format:
1st name: prefix + element name
2nd name: prefix + element name + “ide”
If there is only 1 of the 1st element, no prefix.

15. Prefixes in Covalent Compounds pg 228
Number of atoms
Prefix 1
mono- 6
hexa- 2
di- 7
hepta- 3
tri- 8
octa- 4
tetra- 9
nona- 5
penta- 10
deca-

16. Anyone want a cold glass of dihydrogen monoxide?

17. Examples Name the following CO Carbon Monoxide CO2 N2O Carbon Dioxide
Cl2O8
Carbon Monoxide
Carbon Dioxide
Dinitrogen Monoxide
Dichlorine Octoxide

18. Writing Formulas for Binary Molecular Compounds
Use the prefixes in the name to tell you the subscript of each element in the formula.
Then write the correct symbols for the two elements with the appropriate subscripts.
The least electronegative element is written first
Dinitrogen Tetroxide - N2O4

19. Examples Write formulas for the following: Nitrogen Monoxide
Carbon Tetrachloride
Diphosphorous Pentoxide NO
CCl4
P2O5

20. Naming Acids
An acid is a compound that contains one or more hydrogen atoms and produces hydrogen ions (H+) when dissolved in water.
All acids begin with hydrogen
General Format: HnX
“X” represents a monatomic or polyatomic anion.
“n” represents the number of hydrogen ions

21. 3 Rules for Naming Common Acids
If the name of “X” ends in -ate:
____________-ic acid
If the name of “X” ends in -ite:
____________-ous acid
If the name of “X” ends in -ide:
hydro-__________-ic acid

22. Name these acids H2SO4 HCl H2S HNO3 HClO2 Sulfuric Acid
Hydrochloric Acid
Hydrosulfuric Acid
Nitric Acid
Chlorous Acid

23. Writing Formulas for Acids
If the acid ends in –ic, then “X” ends in –ate
If the acid ends in –ous, then “X” ends in –ite
If the acid has hydro-______-ic, then “X” ends in –ide.
The subscript on hydrogen is equal to the charge of “X”.

24. Write the Formula for the Following Acids
Hydrobromic Acid
Carbonic Acid
Phosphoric Acid
Sulfurous Acid
HBr
H2CO3
H3PO4
H2SO3

25. Homework
7.1 pg 251 #14, 15(a-d), 41

26. Ch 7.3 Using Chemical Formulas

27. The Mass of a Mole of an Element
Molar Mass: is the atomic mass of an element expressed in grams/mole (g/mol).
Carbon = g/mol
Hydrogen = 1.01 g/mol
When dealing with molar mass, round off to two decimals g/mol  g/mol

28. The Mass of a Mole of a Compound
You calculate the mass of a molecule by adding up the molar masses of the atoms making up the molecules.
Example: H2O
H = 1.01 g x 2 atoms = 2.02 g/mol
O = g x 1 atom = g/mol
Molar Mass of H2O = 2.02 g/mol g/mol = g/mol
This applies to both molecular and ionic compounds

29. Find the molar mass of PCl3
P = g x 1 atom = g/mol
Cl = g x 3 atoms = g/mol
PCl3 = g g = g/mol
What is the molar mass of Sodium Hydrogen Carbonate (NaHCO3) ?
Na = g x 1 atom = g/mol
H = 1.01 g x 1 atom = 1.01 g/mol
C = g x 1 atom = g/mol
O = g x 3 atoms = g/mol
NaHCO3 =
= g/mol
137.5 g/mol
84.0 g NaHCO3

30. Converting Moles to Mass
You can use the molar mass of an element or compound to convert between the mass of a substance and the moles of a substance.
Mass (g) = # of moles x mass (g)
1 mole
Example: If molar mass of NaCl is g/mol, what is the mass of 3.00 mol NaCl?
Mass of NaCl = 3.00 mol x 58.44g =
1 mol
175 g NaCl

31. Example 2: Moles to Mass
What is the mass of 9.45 mol of Aluminum Oxide (Al2O3)?
Find molar mass of Al2O3
= g/mol
Mass = 9.45 mol Al2O3 x g Al2O3
1 mol Al2O3
= 964 g Al2O3

32. Converting Mass to Moles
You can invert the conversion factor to find moles when given the mass.
Moles = mass (g) x 1 mole
mass (g)
Example: If molar mass of Na2SO g/mol, how many moles is 10.0 g of Na2SO4?
Moles of Na2SO4 = 10.0 g x 1 mol = g
= mol Na2SO4

33. Example 2: Mass to Moles
How many moles are in 75.0 g of Dinitrogen Trioxide?
Find molar mass of N2O3
= g/mol
Moles = 75.0 g N2O3 x 1 mole =
76.02 g
N2O3
0.987 mol N2O3

34. Percent Composition
Percent Composition: the relative amount of the elements in a compound.
Also known as the percent by mass
It can be calculated in two ways:
Using Mass Data
Using the Chemical Formula
% mass of element= mass of element x100% mass of compound

35. Example
When a g sample of a compound containing Mg and O is decomposed, 5.40 g O is obtained. What is the % composition of this compound?
Mass of compound: g
Mass of oxygen: 5.40 g O
Mass of magnesium: g g = 8.20 g Mg
% Mg = 8.20 g Mg x 100% =
13.60 g
% O = 5.40 g O x 100% =
60.3%
39.7%

36. Find the percent composition of Cu2S. Find mass of Cu and S
Cu = x 2 = g
S = g
Find mass of Cu2S
g g = g
% Composition
Cu = g x 100% = g
S = g x 100% =
79.85%
20.15%

37. Find the percentage by mass of water in the hydrate Na2CO310H2O
Find the mass of 10H2O
H2O = 10 mol x g/mol = g
Find the mass of Na2CO310H2O
Na2CO3 = 1 mol x g/mol = g
g g = g
% By Mass of Water
H2O = g H2O x 100% =
286.2 g Na2CO310H2O
62.96%

38. Homework
7.3 pg 253 #30-33

39. Ch 7.4 Determining Chemical Formulas

40. Empirical Formulas
Empirical Formula: shows the smallest whole-number ratio of the atoms of the elements in a compound.
Example:
The Empirical Formula for Hydrogen Peroxide (H2O2) is HO with a 1:1 ratio.
The Empirical Formula for Carbon Dioxide (CO2) is CO2 with a 1:2 ratio.

41. Determining the Empirical Formula of a Compound
A compound is found to contain 25.9% Nitrogen and 74.1% Oxygen. What is the Empirical Formula of the compound?
25.9 g N x 1 mol N = g N
74.1 g O x 1 mol O = g O
N1.85O4.63 = N1O2.5 = N2O5
1.85 mol N
4.63 mol O

42. Molecular Formulas
Molecular Formula: tells the actual number of each kind of atom present in a molecule of a compound
Example:
The Molecular Formula for Hydrogen Peroxide is H2O2.
The Molecular Formula for Carbon Dioxide is CO2
It is possible to find the Molecular Formula using the Empirical Formula if you know the molar mass of the compound.

43. Finding the Molecular Formula
Calculate the molecular formula of a compound whose molar mass is 60.0 g/mol and empirical formula is CH4N
Step 1: Find the empirical formula molar mass
(4 x 1.01) = g/mol
Step 2: Divide molar mass by EF molar mass
60.0 g/mol = 1.99  g/mol
Step 3: Multiply empirical formula by 2
CH4N x 2 = C2H8N2

44. Homework
7.4 pg 253 #36-38