1. Chemical Naming and Moles Chapter 9-10
Unit 6
Chemical Naming and Moles
Chapter 9-10

2. Naming Ions Positive Ions, cations, simply retain their name.
Na+  Sodium Ion
Mg2+  Magnesium Ion

3. Naming Ions Negative Ions, anions, change ending of element to –ide
Cl-  Chloride Ion
Br-  Bromide Ion

4. Unique Anions Names N3-  Nitride Ion O2-  Oxide Ion
S2-  Sulfide Ion
P3-  Phosphide Ion

5. Stock System
Used when the metal has more than one positive oxidation number
Use a roman numeral to indicate the charge of the ion

6. Roman Numerals Cation Charge Roman Numeral +1 I +2 II +3 III +4 IV +5+6 VI +7
VII +8
VIII

7. Stock System Examples Fe+2  Iron(II) Cu+  Copper(I)
Mn+7  Manganese(VII)
Au+3  Gold(III)
Cr+6  Chromium(VI)
Pb+4  Lead(IV)

8. Polyatomic Ions
Selected polyatomic ions are on Table E in the Reference Tables.
Polyatomic ions keep their names in most chemical names

9. Naming Systems Ionic System Binary Covalent System (Prefixes)
Stock System (Roman Numerals)
Binary Covalent System (Prefixes)

10. When to use ionic system
First element is a metal
NaCl
More than 2 elements, meaning polyatomic ion is involved
NH4Cl

11. Naming Ionic Compounds
Name positive ion first, then negative ion.
NaCl  Sodium chloride
Fe(OH)2  Iron(II) hydroxide

13. Binary Covalent Compounds
Use when compound is 2 nonmetals
Including metalloids
Use a prefix system to indicate the number of atoms for each element
Second element ends in -ide

14. Prefixes Number of atoms Prefix 1 mono- 2 di- 3 tri- 4 tetra- 5 penta-6
hexa- 7
hepta- 8
octa-

15. Example
N2Cl3
Dinitrogen Trichloride

16. Exceptions
When there is only one atom of the first element, do not use mono- prefix.
CO2
Carbon dioxide
OF2
Oxygen difluoride

17. Exceptions
When an element starts with a vowel, drop any o or a at the end of a prefix CO
Carbon Monooxide
Carbon Monoxide
P2O5
Diphosphorus Pentaoxide
Diphosphorus Pentoxide

20. Words representing Numbers
Dozen 12
Baker’s Dozen 13
Gross
Ream
Mole x 1023

21. Avogadro’s Number
6.02 x 1023
Number of representative particles in a mole
Mole number
1 mol He = 6.02 x 1023 atoms
1 mol H2 = 6.02 x 1023 molecules
1 mol H2 = x 1023 atoms

22. Mole-Mass relationship
1 mole of any element equals the atomic mass in grams
1 mol Carbon = 12.01g Carbon
1 mol Iron = 55.8g Iron

23. Gram Formula Mass Mass of the formula in g/mol
Simply add the atomic masses of each element in the formula together
H2O = = 18 g/mol
Also known as gram atomic mass, gram molecular mass, molar mass

24. Rounding Round most masses to the nearest whole gram Except:
Copper, Cu  63.5
Chlorine, Cl  35.5

25. Practice KNO3 = 39 + 14 + 16(3) = 101 g/mol
C6H14 = 12(6) + 1(14) = 86 g/mol
CuSO4 = (4) = g/mol

26. Mole - Mass Conversion
Example: 96 g of Oxygen gas = ? mol

27. Practice How many moles are there in 506g of ethanol, C2H6O?
What is the mass of 8 moles of CCl4?
11 mol C2H6O
1232g CCl4

28. Molar Volume At STP, 1 mol of any gas occupies 22.4L of space
Examples:
2 mol of He occupies how much space at STP?
11.2L will have how much H2 gas at STP?
44.8L
0.5 mol

29. Mole Road Map

31. Review How many moles are in 584g of SF6?
How many grams are in 6 mol of HCl?
584g
X =
= 4 mol
146 g/mol X
6 mol =
= 219 g
36.5 g/mol

32. Percent Composition

33. Example
H2O H
11.1% H O
88.9% O

34. Another Example
NH3 N H
82.4% N
17.6% H

35. Percent Composition
What is the percent composition of oxygen in H2SO3?
58.5%

36. Percent Composition
What is the percent composition of aluminum in Al2(SO4)3?
15.8%

37. Percent Composition
What is the percent composition of nitrogen in NH4NO3?
35%

38. Hydrates
Compounds that have a specific number of water molecules attached
Dot means plus (+)
gfm = (18) = 249.5g/mol
CuSO4·5H2O

40. O H Formulas Molecular Formula Structural Formula
Shows the total number of atoms in a molecule
H2O
Structural Formula
Shows the total number of atoms in a molecule, and how the bonds are arranged O H

41. Empirical Formula Simplest Whole-Number ratio of atoms in a compound
Examples
CO2
P4O10 P2O5
C6H12O6 CH2O

42. Empirical Formula
Molecular Formula is a multiple of the Empirical Formula

43. Examples Compound that has 4 carbon atoms for every 8 hydrogen atoms
C4H8 CH2
Compound that has 6 carbon atoms for every 18 hydrogen atoms
C6H18 CH3

44. Empirical Formula
A molecular formula has an empirical formula of CH2 and a molecular mass of 28 g/mol.
A molecular formula has an empirical formula of CH2 and a molecular mass of 42 g/mol.
C2H4
C3H6