1. Chemical Formulas & Naming (Nomenclature)
Chapter 7
Chemical Formulas
&
Naming (Nomenclature)

2. Chemical Formulas
Chemical formulas indicate the relative # of each type of atom in a compound
Ex: C8H18 (octane) Al2(SO4)3
3x everything in ( )
2 Aluminum
1 Sulfur & 4 Oxygen
8 Carbons
18 Hydrogens

3. Monatomic Ions Monatomic Ions are formed from a single atom
(+) positive ions = cations = lose e-
(-) negative ions = anions = gain e-
Atoms gain or lose electrons in an attempt to achieve a full valence shell (octet)

4. Naming Monatomic Ions Cations(+) K+ : potassium Mg2+ : magnesium
D-Block Cations
Cu+ : copper(I)
Fe3+ : iron(III)
* Use roman numerals to indicate the positive charge
Anions(-)
F- : fluoride
N3- : nitride
O2- : oxide
* Element name ending in ide

5. Binary Ionic Compounds
Binary ionic compounds are made from the combination of cations and anions
The charges of the positive and negative ions must balance out to zero (neutral).
Ex: Mg2+ + Br -  MgBr2

6. Crossing Over Method Write the cation followed by the anion
2) Cross over charges to form subscripts
3) Check to make sure charges are balanced

7. Naming Binary Ionic Compounds
Combine the names of both ions
Ex: Al2O3 = aluminum oxide
AgCl = silver chloride
CaBr2 = calcium bromide
NaCl = sodium chloride

8. Stock System of Nomenclature
Use Roman #’s to indicate the charge of the cation(+)
Ex: CuCl2 = copper (II) chloride
FeO = iron (II) oxide
Fe2O3 = iron (III) oxide
SnF4 = tin (IV) fluoride
ZnBr2 = zinc (II) bromide

9. Polyatomic Ions
Polyatomic ions are molecules with an overall positive or negative charge
Oxyanions- polyatomic ions containing oxygen
Ex: NO2- = nitrite
NO3- = nitrate
ClO2- = chlorite ClO- = hypochlorite
ClO3- = chlorate ClO4- = hyperchlorate

10. Naming Binary Molecular Compounds (covalent bonds)
Use prefixes to indicate the # of each type of element in the molecule
Write the element with the smaller group number 1st
 If both elements have the same group number, write the element with the greater period number 1st
Use a prefix is there is more than 1 of the 1st element

11. Naming Binary Molecular Compounds (covalent bonds) continued…
3) Write the second element always using a prefix to indicate the # of atoms present
Note: When an element begins with a vowel, the “o” or the “a” at the end of a prefix is dropped.
Example: monoxide, not monooxide

12. Prefixes 1 = mono 2 = di 3 = tri 4 = tetra 5 = penta 6 = hexa
7 = hepta
8 = octa
9 = nona
10 = deca

13. Molecular Compounds: Ex: P4O10 = tetraphosphorus decaoxide
N2O = dinitrogen monoxide
SO3 = sulfur trioxide

14. Acids & Salts Binary Acids: contain H & a halogen(17)
Ex: HCl = hydrochloric acid
Oxyacids: contain H, O, & a 3rd element
Ex: H2SO4 = sulfuric acid
HNO3 = nitric acid
Salts: formed from a cation and an anion of an acid
Ex: HCl + NaOH  NaCl + H2O
(Acid)
(Base)
(Salt)

15. Oxidation #’s (states)
Indicate the general distribution of electrons among the bonded atoms in a molecular compound or polyatomic ion
Rules:
The oxidation number of a free (single) element is always zero (Ex. O2, Na)
The more EN element has a – charge and the less EN element has a + charge
Fluorine (F) always has an oxidation # of -1 (most EN)
Hydrogen (H) has an oxidation # of +1 when combined with nonmetals, -1 when combined with metals
The oxidation # of oxygen (O) is usually -2
The oxidation # of other elements is based on the elements group #
The sum of the oxidation numbers of all of the atoms in a neutral compound is 0
The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion

16. Examples
UF6 H2SO4 ClO3-

17. Chemical Formulas Review
Formula Mass- sum of the average atomic masses of all elements in a compound, units (amu)
Ex: H2O
Molar Mass- mass of one mole of a substance, numerically equivalent to formula mass, different units (g/mol)
Ex: H2SO4

18. ***The Mole Road*** To convert from Moles  Mass
multiply (x)
multiply (x)
# of Atoms
Molar Mass(g/mol)
Mass(g)
Moles
6.02 x 1023
divide (/)
divide (/)
To convert from Moles  Mass
Multiply the moles by the molar mass of the compound
To convert from Moles  # of Atoms
Multiply the moles by 6.02 x 1023

19. % Composition
The percentage by mass of each element in a compound
Mass of Element
x 100 = % Composition
Mass of Compound
Ex: H2SO4

20. Deriving Chemical Formulas
Empirical Formula- the symbols of the elements in a compound w/ subscripts showing the smallest whole # ratio between the atoms
Ex: C6H12O6  CH2O
Chemical Formula
Empirical Formula

21. Determining Empirical Formulas
Assume g sample mass
(if no sample mass is given)
2) Convert % Composition to Mass for each element
71 % A = 71 g of “A”
29 % B = 29 g of “B”

22. Determining Empirical Formulas
3) Determine the # of Moles of each element using the Molar Mass
71.0 g A x 1 mol A/10.0 g A = 7.10 mol A
29.0 g of B x 1 mol B/2.01 g B = 14.5 mol B
4) Determine Mole Ratio by dividing both values from step 3 by the lower #
A: / 7.10 = 1.0
B: / 7.10 = 2.0
= AB2

23. *** If the mole ratio is NOT a whole # ratio after dividing then…
Multiply both #’s by the same factor to get a whole # ratio
Ex: A = x 2 = 3.0
B = x 2 = 2.0
A = x 4 = 9.0
B = x 4 = 4.0

24. Determining Empirical Formulas
Convert % of each Element to Grams(g)
Convert Grams(g) to Moles(mol)
Divide Moles by Lowest #
Determine Whole # Ratio

25. Practice Empirical Formulas
32.38% Sodium(Na)
22.65% Sulfur(S)
44.97% Oxygen(O)
2) % Carbon
10.6% Hydrogen
42.1% Oxygen

26. Calculating Molecular Formulas
Divide:
Actual Mass / Empirical Formula Mass
2) Multiply:
Multiply the # from step 1 by the subscript of each element in the empirical formula
Ex: BH3
Actual Mass = g/mol

27. END OF CHAPTER 7 NOTES!