1. Chemical Formulas and Chemical Compounds
Chapter 7
Chemical Formulas and Chemical Compounds

2. Section 1 – Chemical Names and Formulas

3. Significance of Chemical Formula
Chemical formula shows number and types of atoms in compound
C8H18
Subscript indicates that
there are 8 carbon atoms
in a molecule of octane.
Subscript indicates that
there are 18 hydrogen atoms
in a molecule of octane.

4. Ionic compound made of mixture of + and – ions held by attraction
Combination of cation and anion
Al2S3
Al3+ is the positively charged cation. S2- is the negatively charged anion

5. Monoatomic Ions Monoatomic ions – ions formed from single atom
Group 1 – lose one e-, form +1
Group 2 – lose two e-, form +2
Groups 15, 16, 17 gain electrons to form anions
-1, -2, -3

6. Not all main-group elements form ions easily
C and Si form covalent bonds
Transition metals can form +1, +2, +3, +4

7. Naming Monoatomic Ions
Cations named same as element name
K+ = potassium
Mg2+ = magnesium
Anions named by dropping ending and adding –ide
F- = fluoride
N3- = nitride

8. Binary Ionic Compounds
Binary compounds  compounds made of two different elements
Total positive and negative charge must be equal
Mg2+ Br-
= +1
(-1) = 0
Formula = MgBr2

9. Naming Binary Ionic Compounds
Nomenclature  naming system
Combine names of cations and anions
Cation name ALWAYS comes first
Then anion name
Al2O3
Aluminum oxide

10. Practice Problem 1
Write formulas for the binary ionic compounds formed between the following elements:
a. potassium and iodine KI
b. magnesium and chlorine
MgCl2
c. sodium and sulfur
Na2S

11. d. aluminum and sulfur
Al2S3
e. aluminum and nitrogen
AlN

12. Practice Problem 2 Name the binary ionic compounds: a. AgCl
Silver chloride
b. ZnO
Zinc oxide
c. CaBr2
Calcium bromide

13. d. SrF2
Strontium fluoride
e. BaO
Barium oxide
f. CaCl2
Calcium chloride

14. Stock System
Some elements form two or more cations with different charges
Stock system uses Roman numerals to show ion’s charge
Roman numeral included in () right after metal name
Fe2+ Iron (II)

15. CuCl2 Copper (II) chloride
Roman
numeral
indicating
charge
Name of anion
Name of
cation

16. Practice Problem 1
Write the formula and give the name for the compounds formed between the following ions:
a. Cu2+ and Br−
CuBr2, copper(II) bromide
b. Fe2+ and O2−
FeO, iron(II) oxide
c. Pb2+ and Cl−
PbCl2, lead(II) chloride

17. d. Hg2+ and S2−
HgS, mercury(II) sulfide
e. Sn2+ and F−
SnF2, tin(II) fluoride
f. Fe3+ and O2−
Fe2O3, iron(III) oxide

18. Polyatomic Ions Polyatomic ion  ion containing more than one atom
Most are oxyanions  polyatomic ions that contain oxygen
NO2- NO3-
Ion with more oxygen ends in –ate (nitrate)
Ion with less oxygen ends in –ite (nitrite)

19. Hypo- and Hyper- One oxyanion family has 4 members
ClO- ClO2- ClO3- ClO4-
Hypochlorite Chlorite Chlorate Perchlorate
Naming same as binary ionic compounds
Cation first, anion second

21. Practice Problem 1 Write formulas for the following ionic compounds:
a. sodium iodide
NaI
b. calcium chloride
CaCl2
c. potassium sulfide
K2S
d. lithium nitrate
LiNO3

22. e. copper(II) sulfate
CuSO4
f. sodium carbonate
Na2CO3
g. calcium nitrite
Ca(NO2)2
h. potassium perchlorate
KClO4

23. Practice Problem 2 Give the names for the following compounds: a. Ag2O
silver oxide
b. Ca(OH)2
calcium hydroxide
c. KClO3
potassium chlorate

24. d. NH4OH
ammonium hydroxide
e. FeCrO4
iron(II) chromate
f. KClO
potassium hypochlorite

25. Naming Binary Molecular Compounds
Molecular compounds made of covalently bonded molecules
Usually happens between 2 NON-metals

26. Prefix System Uses prefixes to show how many atoms present Number1
Mono- 2
Di- 3
Tri- 4
Tetra- 5
Penta- 6
Hexa- 7
Hepta- 8
Octa- 9
Nona- 10
Deca-

27. Examples CCl4  carbon tetrachloride P4O10  tetraphosphorous decoxide
There are rules for determining which ion comes first

28. Rules
The less-electronegative element is given first. If only 1 atom, no prefix used
Second element named by combining (a) prefix, (b) element root name, (c) –ide
The o or a at the end of prefix dropped when word following beings with vowel

29. Practice Problem Name the following binary molecular compounds: a. SO3
sulfur trioxide
b. ICl3
iodine trichloride
c. PBr5
phosphorous pentabromide

30. Acids Most acids either binary acids or oxyacids
Binary acids  acids that are made of 2 elements
Usually H and one of the halogens
Oxyacids  acids that contain H, O, and a 3rd element (usually nonmetal)

31. Naming Acids Binary acids – hydro(element root)ic acid
Ex. HCl – hydrochloric acid
Oxyacids
– with less oxygen, (element root)ous acid
– with more oxygen, (elementroot)ic acid
Ex. HNO2 = nitrous acid
HNO3 = nitric acid

32. Formula
Name HF
Hydrofluoric acid
HCl
Hydrochloric acid
HBr
Hydrobromic acid HI
Hydroiodic acid
H3PO4
Phosphoric acid
HNO2
Nitrous acid
HNO3
Nitric acid
H2SO3
Sulfurous acid
H2SO4
Sulfuric acid
CH3COOH
Acetic acid
HClO
Hypochlorous acid
HClO2
Chlorous acid
HClO3
Chloric acid
HClO4
Perchloric acid
H2CO3
Carbonic acid

33. Salts
Salt  an ionic compound made of a cation and an anion from an acid
Table salt – NaCl – contains anion from HCl
Some salts have anions where one or more H atoms from acid are kept
Named by adding hydrogen OR prefix bi- to anion name
HCO hydrogen carbonate ion
bicarbonate ion

34. Practice Problem 1
Write formulas for the compounds formed between the following:
a. aluminum and bromine
AlBr3
b. sodium and oxygen
Na2O
c. magnesium and iodine
MgI2
d. Pb2+ and O2−
PbO

35. e. Sn2+ and I−
SnI2
f. Fe3+ and S2−
Fe2S3
g. Cu2+ and NO3−
Cu(NO3)2
h. NH4 + and SO42−
(NH4)2SO4

36. Practice Problem 2
Name the following compounds using the Stock system:
a. NaI
sodium iodide
b. MgS
magnesium sulfide
c. CaO
calcium oxide

37. d. K2S
potassium sulfide
e. CuBr
copper (I) bromide
f. FeCl2
iron (II) chloride

38. Practice Problem 3 Write formulas for each of the following compounds:
a. barium sulfide
BaS
b. sodium hydroxide
NaOH
c. lead(II) nitrate
Pb(NO3)2
d. potassium permanganate
KMnO4

39. e. iron(II) sulfate
FeSO4
f. diphosphorus trioxide
P2O3
g. disulfur dichloride
S2Cl2
h. carbon diselenide
CSe2

40. i. acetic acid
CH3COOH
j. chloric acid
HClO3
k. sulfurous acid
H2SO3
l. phosphoric acid
H3PO4

41. Section 2 – Oxidation Numbers
The charges on the ions making an ionic compound reflect the electron distribution of the compound. In order to indicate the general distribution of electrons among the bonded atoms in a molecular compound or a polyatomic ion, oxidation numbers, also called oxidation states, are assigned to the atoms composing the compound or ion.
Unlike ionic charges, oxidation numbers do not have an exact physical meaning. However, oxidation numbers are useful in naming compounds, in writing formulas, and in balancing chemical equations. And they are helpful in studying certain types of chemical reactions.

42. Assigning Oxidation Numbers
Shared electrons assumed to belong to more- electronegative atom in each bond

43. Rules 1. Atoms in a pure element have Ox# of zero.
Na, O2, P4 = 0
2. More-electronegative element in binary molecular compound assigned the number equal to negative charge it has as ion. Less electronegative assigned equal to positive charge it has as ion.
3. Fluorine always is -1 because it is most electronegative.

44. 4. Oxygen has -2 in almost all compounds, except
Peroxides (H2O2, O = -1)
Compounds with halogens (O = +)
5. H = +1 in all compounds with more electronegative element, H = -1 in compounds with metals
6. Sum of Ox# of all atoms in neutral compound = zero
7. Sum of Ox# of all atoms in polyatomic ion = charge of ion
8. Ox# of ions in ionic compounds = its charge

45. Sample Problem 1 Assign Ox# to each atom in UF6.
Start by placing known Ox#s above appropriate elements.
From rules, we know F is always -1

46. Multiply known Ox#s by appropriate number of atoms and place totals underneath matching elements.
There are 6 F atoms, 6 x -1 = -6

47. UF6 is molecular
According to guidelines, sum of Ox#s must be zero
Total positive Ox#s must be +6

48. Divide total positive Ox#s by number of atoms
+6 ÷ 1 = +6

49. H2SO4
H2 S O4

50. Practice Problem
Assign oxidation numbers to each atom in the following compounds or ions:
a. HCl
+1, -1
b. CF4
+4, -1
c. PCl3
+3, -1

51. d. SO2
+4, -2
e. HNO3
+1, +5, -2
f. KH
+1, -1
g. P4O10
+5, -2

52. h. HClO3
+1, +5, -2
i. N2O5
+5, -2
j. GeCl2
+2, -1

53. Stock System Remember there were two ways to name covalent compounds
Prefix system
Some nonmetals can have more than one oxidation number
Listed in Table A-15 in handout

54. Group 14
Carbon
-4, +2, +4
Group 15
Nitrogen
Phosphorous
-3, +3, +5
Group 16
Sulfur
-2, +4, +6
Group 17
Chlorine
Bromine
Iodine
-1, +1, +3, +5, +7

55. Can use Roman numerals to show oxidation number
Prefix System
Stock System
PCl3
Phosphorous trichloride
Phosphorous(IV) chloride
PCl5
Phosphorous pentachloride
Phosphorous(V) chloride
N2O
Dinitrogen oxide
Nitrogen(I) oxide NO
Nitrogen monoxide
Nitrogen(II) oxide
PbO2
Lead dioxide
Lead(IV) oxide
Mo2O3
Dimolybdenum trioxide
Molybdenum(III) oxide

56. Practice Problem 1
Assign oxidation numbers to each atom in the following compounds or ions:
a. HF
+1, -1
b. CI4
+4, -1
c. H2O
+1, -2
d. PI3
+3, -1

57. e. CS2
+4, -2
f. Na2O2
+1, -1
g. H2CO3
+1, +4, -2
h. NO2 −
+3, -2

58. i. SO42−
+6, -2
j. ClO2 −
+3, -2
k. IO3 −
+5, -2

59. Practice Problem 2
Name each of the following binary molecular compounds according to the Stock system:
a. CI4
carbon(IV) iodide
b. SO3
sulfur(VI) oxide
c. As2S3
arsenic(III) sulfide
d. NCl3
nitrogen(III) chloride

60. Section 3 – Using Chemical Formulas
As you have seen, a chemical formula indicates the elements as well as the relative number of atoms or ions of each element present in a compound. Chemical formulas also allow chemists to calculate a number of characteristic values for a given compound. In this section, you will learn how to use chemical formulas to calculate the formula mass, the molar mass, and the percentage composition by mass of a compound.

61. Formula Masses
Like atoms, a molecule, formula unit, or ion has an average mass (amu)
Simply add amu of all atoms in molecule
H2O
Amu H = 2 x 1.01 amu
Amu O = 1 x amu
Amu H2O = amu

62. Mass of water molecule can be named molecular mass
Mass of 1 NaCl is not molecular mass b/c it’s not a molecule, it’s an ionic compound
Formula mass  sum of the average atomic masses of all the atoms represented in its formula

63. Practice Problem Find the formula mass of each of the following:
a. H2SO4
b. Ca(NO3)2
c. PO43−
d. MgCl2

64. a. H2SO amu
b. Ca(NO3) amu
c. PO43− amu
d. MgCl amu

65. Molar Masses
Molar mass is the same are formula mass, but units are in grams/mole (g/mol)
Molar Mass H2O = ?
H = 2 x 1.01 g/mol
O = 1 x g/mol
H2O = g/mol

66. Practice Problem Find the molar mass of each of the compounds a. Al2S3
b. NaNO3
c. Ba(OH)2

67. a. Al2S g/mol
b. NaNO g/mol
c. Ba(OH) g/mol

68. Molar Mass as Conversion Factor
g/mol can be used as a conversion factor to change from moles to mass
What is the mass in grams of 2.50 mol of oyxgen gas?

69. Practice Problem 1
Ibuprofen, C13H18O2, is the active ingredient in many nonprescription pain relievers. Its molar mass is g/mol.
a. If the tablets in a bottle contain a total of 33 g of ibuprofen, how many moles of ibuprofen are in the bottle?
b. How many molecules of ibuprofen are in the bottle?
c. What is the total mass in grams of carbon in 33 g of ibuprofen?

70. a mol C13H18O2
b. 9.6 x 1022 molecules C13H18O2
c. 25 g C

71. Practice Problem 2
How many moles of compound are there in the following?
a g (NH4)2SO4
b. 4.5 kg Ca(OH)2
a mol
b. 61 mol

72. Practice Problem 3 How many molecules are there in the following?
a g H2SO4
b. 125 g of sugar, C12H22O11
a × 1023 molecules
b × 1023 molecules

73. Percent Composition
Percentage composition  percentage by mass of each element in a compound
Divide mass of element in sample of compound by total mass of sample, multiply by 100

74. Mass percentage of element in compound same regardless of sample’s size
Simpler way to calculate….
Determine how many grams of element are present in 1 mole of compound
Then divide this value by molar mass of compound, multiply by 100

75. Sample Problem
Find the percentage composition of copper(I) sulfide, Cu2S.

76. 1. Analyze Given: formula, Cu2S Unknown:
percentage composition of Cu2S

77. 2. Plan formula → molar mass → mass percentage of each element
The molar mass of the compound must be found
Then the mass of each element present in one mole of the compound is used to calculate the mass percentage of each element.

78. 3. Compute
Molar mass: Cu = 2 x g Cu = g Cu S = 1 x g S = g S
159.2 g/mol Cu2S

79. Sample Problem 2
As some salts crystallize from a water solution, they bind water molecules in their crystal structure. Sodium carbonate forms such a hydrate, in which 10 water molecules are present for every formula unit of sodium carbonate. Find the mass percentage of water in sodium carbonate decahydrate, Na2CO3•10H2O, which has a molar mass of g/mol.

80. 1. Analyze Given: chemical formula, Na2CO3•10H2O
molar mass of Na2CO3•10H2O
Unknown:
mass percentage of H2O

81. 2. Plan chemical formula → mass H2O per mole of Na2CO3•10H2O → % water
The mass of water per mole of sodium carbonate decahydrate must first be found.
This value is then divided by the mass of one mole of Na2CO3•10H2O.

82. 3. Compute
1 mol Na2CO3•10H2O contains 10 mol of H2O

83. 4. Evaluate
Checking shows that the arithmetic is correct and that units cancel as desired.

84. Section 4 – Determining Chemical Formulas

85. When a new substance is synthesized or is discovered, it is analyzed to show its percentage composition
From this data, the empirical formula is then determined
An empirical formula consists of the symbols for the elements combined in a compound, with subscripts showing the smallest whole- number mole ratio of the different atoms in the compound

86. For an ionic compound, the formula unit is usually the compound’s empirical formula
For a molecular compound, however, the empirical formula does not necessarily indicate the actual numbers of atoms present in each molecule
For example, the empirical formula of the gas diborane is BH3, but the molecular formula is B2H6
In this case, the number of atoms given by the molecular formula corresponds to the empirical ratio multiplied by two.

87. Calculation of Empirical Formulas
Determine empirical formula from percent composition
Change percent composition to grams
Ex. 29.9% H … g H
Use molar mass to change grams to mole
Divide all moles by smallest mole amount
Result is the subscript for the empirical formula

88. Example Percent composition of diborane is 78.1% B and 21.9% H
In g sample of diborane, there is 78.1 g B and 21.9 g H

89. Mass composition of each element converted to composition in moles by dividing by molar mass

90. 7.22 mol B to 21.7 mol H
Not a ratio of smallest whole numbers
Divide each number of mol by smallest number in ratio
Rounding, empirical formula is BH3

91. Sample Problem 1
Quantitative analysis shows that a compound contains 32.38% sodium, % sulfur, and 44.99% oxygen. Find the empirical formula of this compound.

92. 1. Analyze Given: percentage composition: 32.38% Na, 22.65% S, and
Unknown: empirical formula

93. 2. Plan percentage composition mass moles
smallest whole-number mole ratio of atoms

94. 3. Compute Mass composition (mass of each element in 100.0 g sample):
32.38 g Na,
22.65 g S,
44.99 g O

96. Divide each mole by smallest mole in ratio
Na2SO4

97. 4. Evaluate
Calculating the percentage composition of the compound based on the empirical formula determined in the problem reveals a percentage composition of 32.37% Na, 22.58% S, and % O
These values equal 100%

98. Practice Problem 1
Analysis of a g sample of a compound known to contain only phosphorus and oxygen indicates a phosphorus content of g. What is the empirical formula of this compound?
P2O5

99. Practice Problem 2
A compound is found to contain 63.52% iron and 36.48% sulfur. Find its empirical formula.
FeS

100. Practice Problem 3
Find the empirical formula of a compound found to contain 26.56% potassium, 35.41% chromium, and the remainder oxygen.
K2Cr2O7

101. Practice Problem 4
Analysis of 20.0 g of a compound containing only calcium and bromine indicates that 4.00 g of calcium are present. What is the empirical formula of the compound formed?
CaBr2

102. Calculating Molecular Formulas
Empirical formula contains smallest whole number ratio
Molecular formula is actual formula for compound
Empirical  CH
Molecular  C2H4 (ethene), C3H6 (cyclopropane)

103. x(empirical formula) = molecular formula
Relationship between empirical and molecular:
x(empirical formula) = molecular formula
x is whole-number multiple
Must know molecular mass
Ex. Experiment shows molar mass of diborane (BH3) is g/mol
Empirical formula mass = g/mol

104. Practice Problem 1
The empirical formula of a compound of phosphorus and oxygen was found to be P2O5. Experimentation shows that the molar mass of this compound is g/mol. What is the compound’s molecular formula?
P4O10

105. Practice Problem 2
Determine the molecular formula of the compound with an empirical formula of CH and a formula mass of g/mol.
C6H6

106. Practice Problem 3
A sample of a compound with a formula mass of g/mol is found to consist of g H and 6.92 g O. Find its molecular formula.
H2O2