1. Chapter 7 – Chemical Formulas & Compounds

2. Chapter 7 – Section 2 – Oxidation Numbers
II. Oxidation Numbers
Oxidation numbers or state are assigned to atoms in compounds or ions to indicate their distribution of electrons; they are used to help name compounds and write formulas
Na+1 Mg+2 Al+3 N-3 O-2 Cl-1 Ne0
A. Assigning Oxidation Numbers
Usually shared electrons belong to the most electronegative atom in the compound

3. 1. The atoms in any pure element have an oxidation number of zero
2. In a binary molecular compound, the most electronegative atom is given an oxidation number equal to the charge it would have as an anion; less electronegative atoms are assigned numbers equal to the charge they would have as a cation
3. Fluorine is the most electronegative element, so its oxidation number is always -1

4. 4. Oxygen is usually a -2, unless it is a peroxide, like H2O2, when it will be -1 or combined with fluorine (OF2), when it will be +2
5. Hydrogen is +1 in all compounds where it is the least electronegative, but is -1 in compounds with metals
6. The sum of all the oxidation numbers in a neutral compound must equal zero
7. The sum of the oxidation numbers in a polyatomic ion will equal the charge of the polyatomic ion
8. A monatomic ion’s oxidation number is the same as its charge

5. B. Using Oxidation Numbers for Formulas & Names
1. Many nonmetals can have more than one oxidation number
2. Prefixes can be used to indicate the oxidation numbers (Carbon dioxide or Carbon monoxide)
3. The Stock system can also be used, which uses Roman numerals instead of prefixes [Carbon (I) oxide or Carbon (II) oxide]

6. Chapter 7, Section 3 – Using Chemical Formulas
III. A chemical formula can be used to calculate formula mass, molar mass, and percentage composition of mass for a compound
A. Formula Masses
1. The average atomic masses (mass on the periodic table) of the elements in a compound are added together to get the average atomic mass of the molecule or formula unit

7. 2. The mass of a molecule is referred to as its molecular mass
3. The mass of all the atoms in any unit of a formula (molecule, formula unit, or ion) is referred to as the formula mass
B. Molar Masses
1. A mole of any substance will be x 1023 particles of that substance
2. The molar mass of the substance will be the atomic mass of the substance in grams

8. 3. The molar mass and the formula mass will be numerically equal (the same), but the formula mass unit will be amu, and the molar mass unit will be grams
C. Molar Mass as a Conversion Factor
1. Molar mass can be used as a conversion factor to change moles to grams
2. Multiply moles by molar mass to get grams
3. To change grams to moles, divide grams by molar mass

9. D. Percentage Composition
1. To find the percentage of an element in a compound, divide the mass of the element by the mass of the compound and multiply by 100
2. The percentage of the element will be the same no matter how much of the compound you have
3. Divide the grams of the element in 1 mole of the compound by the molar mass of the compound and multiply by 100
4. The percentage of mass for each element in a compound is known as the percentage composition of the compound

10. Chp. 7 – Sec. 4 – Determining Chemical Formulas
IV. An empirical formula is made of element symbols and subscripts to show the smallest whole number mole ratio of each atom in the compound
For ionic compounds, the formula unit is usually the empirical formula
For covalent compounds, the empirical formula does not necessarily tell the actual number of atoms in the molecule
An example is diborane (B2H6), whose empirical formula is BH3, because the empirical formula is a more simple ratio than the molecular formula

11. A. Calculations of Empirical Formulas
1. Convert percentage composition to mass
2. Assume you have grams of the compound
3. Find the mass of each element in the 100 grams
4. Change the mass of each element to moles
5. Determine which element has the smallest number of moles
6. Divide both numbers of moles by the smallest one (to find the smallest ratio)
7. Round the answers to the nearest whole numbers, and these will be the subscripts in the empirical formula

12. 8. If mass is known instead of percentage composition, then change mass to moles and continue at Step 5
B. Calculation of Molecular Formulas
1. The molecular formula shows the actual number of atoms in a molecule
2. Two different compounds can have the same empirical formula, but have different molecular formulas [Examples: ethene (molecular – C2H4; empirical – CH2) and cyclopropane (molecular – C3H6: empirical – CH2)
3. To change the empirical formula to the molecular formula, divide the molecular mass by the empirical mass

13. 4. When you divide, if your answer does not come out to be a whole number, then follow the 0-9 rule, which states that if the 1st decimal number is a 0, drop it and keep the whole number, but if the 1st decimal number is a 9, round up to the next whole number
5. If the first decimal number is not 0 or 9, then begin multiplying your answer by 2 (or 3 or 4 or 5, etc. if necessary) until you get a number that is 0 or 9
6. The answer you get when you divide (and multiply, if necessary) will then be multiplied by each subscript in the empirical formula

14. 7. After multiplying, this will be the molecular formula
C. Hydrates
1. A hydrate is a molecule that has one or more water molecules chemically attached to the compound
2. To show this chemical attachment, the formula shows a dot between the compound and the water (CuSO4 · 5H2O)
3. The coefficient in front of the water molecule shows how many water molecules are attached
4. In calculating the molecular formula for a hydrate, the numbers obtained from dividing and multiplying will be used as coefficients, rather than subscripts

15. Section Review Questions, pp. 251-252
1. a) Ions formed from a single atom
b) Examples include Na+1, Mg+2, and Cl-1
2. The nitrate ion, NO3-1, has one more oxygen atom than the nitrite ion, NO2-1, does
3. a) K+1 b) Ca+2 c) S-2 d) N-3 e) Ba+2 f) Br-1
4. a) Na+1 b) Al+3 c) Cl-1 d) N-3 e) Fe+2 f) Fe+3
5. a) potassium ion (or potassium cation)
b) magnesium ion (or magnesium cation)
c) Aluminum ion (or aluminum cation)
d) chloride ion (or chloride anion)
e) oxide ion (or oxide anion)
f) calcium ion (or calcium cation)

16. 6. a) NaI b) CaS c) ZnCl2 d) BaF2 e) Li2O
7. a) potassium chloride b) calcium bromide
c) lithium oxide d) magnesium chloride
8. a) CrF4, chromium (II) fluoride
b) NiO, nickel (II) oxide
c) Fe2O3, iron (II) oxide
9. The less electronegative element is written first
10. a) carbon dioxide b) carbon tetrachloride
c) phosphorus pentachloride
d) selenium hexafluoride
e) diarsenic pentoxide

17. 11. a) CBr4 b) SiO2 c) P4O10 d) As2S3
12. Binary acids, such as HCl and HBr, contain only two elements, usually hydrogen and a halogen. Oxyacids, such as H2SO4 and HNO3, contain hydrogne, oxugen, and a third element
13. a) an ionic compound composed of a cation and the anion from an acid
b) Examples include NaCl and MgSO4
14. a) hydrofluoric acid b) hydrobromic acid
c) nitric acid d) sulfuric acid
e) phosphoric acid

18. 15. a) H2SO3 b) HClO3 c) HCl d) HClO
e) HClO4 f) H2CO3 g) HC2H3O2
16. a) NaF b) CaO c) K2S d) MgCl2 e) AlBr3
f) Li3N g) FeO
17. a) ammonium ion b) chlorate ion
c) hydroxide ion d) sulfate ion e) nitrate ion
f) carbonate ion g) phosphate ion
h) acetate ion i) hydrogen carbonate ion (or
j) chromate ion bicarbonate ion)
18. a) NH4+1 b) C2H3O2-1 c) OH-1 d) CO3-2
e) SO4-2 f) PO4-3 g) Cu+2 h) Sn+2 i) Fe+3
j) Cu+1 k) Hg2+2 l) Hg+2

19. 19. a) iron (II) ion b) iron (III) ion c) lead (II) ion
d) lead (IV) ion e) tin (II) ion f) tin (IV) ion
20. a) carbon (IV) bromide b) silicon (IV) oxide
c) phosphorus (V) oxide d) arsenic (III) sulfide
21. a) PI3 b) SCl2 c) CS2 d) N2O5
22. a) numbers assigned to bonded atoms in molecular compounds or polyatomic ions to indicate the general distribution of electrons
b) they aid in naming compounds, writing formulas, balancing chemical equations, and studying chemical reactions

20. 23. a) sodium chloride b) potassium fluoride
c) calcium sulfide d) cobalt (II) nitrate
e) iron (III) phosphate f) mercury (I) sulfate
g) mercury (II) phosphate
24. a) +1, -1 b) +3, -1 c) +4, -2 d) +1, -1
e) +5, -2 f) +1, +5, -2
25. a) +5, -2 b) +7, -2 c) +5, -2 d) +6, -2
e) +4, -2