Chapter 7 – Chemical Formulas
& Compounds
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
• 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. 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
• 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]
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
• 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 6.022 x
1023 particles of that substance
• 2. The molar mass of the substance will be
the atomic mass of the substance in grams
• 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
• 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
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
• A. Calculations of Empirical Formulas
• 1. Convert percentage composition to mass
• 2. Assume you have 100.00 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
• 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
• 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
• 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
Section Review Questions, pp. 251-252
• 1. a) Ions formed from a single atom
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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)
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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
• 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
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b) Examples include NaCl and MgSO4
• 14. a) hydrofluoric acid b) hydrobromic acid
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c) nitric acid d) sulfuric acid
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e) phosphoric acid
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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
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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
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b) they aid in naming compounds, writing
formulas, balancing chemical equations, and
studying chemical reactions
• 23. a) sodium chloride b) potassium fluoride
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c) calcium sulfide d) cobalt (II) nitrate
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e) iron (III) phosphate f) mercury (I) sulfate
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g) mercury (II) phosphate
• 24. a) +1, -1 b) +3, -1 c) +4, -2 d) +1, -1
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e) +5, -2 f) +1, +5, -2
• 25. a) +5, -2 b) +7, -2 c) +5, -2 d) +6, -2
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e) +4, -2