Chemical Formulas and Chemical Compounds

1
Chemical Formulas and Chemical Compounds

• Chapter 7

2
What Do You Think?

• CCl4 MgCl2
• Guess the name of each of the above compounds
  based on the formulas written.
• What kind of information can you discern from the
  formulas?
• Guess which of the compounds represented is
  molecular and which is ionic.
• Chemical formulas form the basis of the language
  of chemistry and reveal much information about
  the substances they represent.

3
Chemical Names and Formulas

• Section 1

4
Significance of a Chemical Formula

• A chemical formula indicates the relative number
  of atoms of each kind in a chemical compound.
• For a molecular compound, the chemical formula
  reveals the number of atoms of each element
  contained in a single molecule of the compound.
• example octane C8H18

The subscript after the C indicates that there
are 8 carbon atoms in the molecule.
The subscript after the H indicates that there
are 18 hydrogen atoms in the molecule.
5
Significance of a Chemical Formula

• The chemical formula for an ionic compound
  represents one formula unitthe simplest ratio of
  the compounds positive ions (cations) and its
  negative ions (anions).
• example aluminum sulfate Al2(SO4)3
• Parentheses surround the polyatomic ion
  to identify it as a unit. The subscript 3 refers
  to the unit.
• Note also that there is no subscript for sulfur
  when there is no subscript next to an atom, the
  subscript is understood to be 1.

6
Monatomic Ions

• Many main-group elements can lose or gain
  electrons to form ions.
• Ions formed form a single atom are known as
  monatomic ions.
• example To gain a noble-gas electron
  configuration, nitrogen gains three electrons to
  form N3 ions.
• Some main-group elements tend to form covalent
  bonds instead of forming ions.
• examples carbon and silicon

7
Naming Monatomic Ions

• Monatomic cations are identified simply by the
  elements name.
• examples
• K is called the potassium cation
• Mg2 is called the magnesium cation
• For monatomic anions, the ending of the elements
  name is dropped, and the ending -ide is added to
  the root name.
• examples
• F is called the fluoride anion
• N3 is called the nitride anion

8
The Stock System of Nomenclature

• Some elements such as iron, form two or more
  cations with different charges.
• To distinguish the ions formed by such elements,
  scientists use the Stock System of nomenclature.
• The system uses a Roman numeral to indicate an
  ions charge.
• examples Fe2 iron(II)
• Fe3 iron(III)

9
Common Monatomic Ions
10
Common Monatomic Ions
11
Name These

• Na1
• Ca2
• Al3
• Fe3
• Fe2
• Pb2
• Li1

• Cl-1
• N-3
• Br-1
• O-2
• Ga3

12
Write Formulas for These

• Potassium ion
• Magnesium ion
• Copper (II) ion
• Chromium (VI) ion
• Barium ion
• Mercury (II) ion

• Sulfide ion
• iodide ion
• phosphide ion
• Strontium ion

13
Binary Ionic Compounds

• Compounds composed of two elements are known as
  binary compounds.
• In a binary ionic compound, the total numbers of
  positive charges and negative charges must be
  equal.
• The formula for a binary ionic compound can be
  written given the identities of the compounds
  ions.
• example magnesium bromide
• Ions combined Mg2, Br, Br
• Chemical formula MgBr2

14
Binary Ionic Compounds

• A general rule to use when determining the
  formula for a binary ionic compound is crossing
  over to balance charges between ions.
• example aluminum oxide
• 1) Write the symbols for the ions.
• Al3 O2
• 2) Cross over the charges by using the absolute
  value of each ions charge as the
  subscript for the other ion.

15
Binary Ionic Compounds

• 3) Check the combined positive and negative
  charges to see if they are equal.
• (2 ? 3) (3 ? 2?) 0
• The correct formula is Al2O3
• Al3- O2-
• Al2O3

16
Writing the Formula of an Ionic Compound
17
Naming Binary Ionic Compounds

• The nomenclature, or naming system, for binary
  ionic compounds involves combining the names of
  the compounds positive and negative ions.
• The name of the cation is given first, followed
  by the name of the anion
• example Al2O3 aluminum oxide
• For most simple ionic compounds, the ratio of the
  ions is not given in the compounds name, because
  it is understood based on the relative charges of
  the compounds ions.

18
Example

• Write the formulas for the binary ionic compounds
  formed between the following elements
• a. zinc and iodine
• b. zinc and sulfur

19
Example II

• Write the formula and give the name for the
  compound formed by the ions Cr3 and F.

20
Your Turn

• Write the Formulas for the following binary
  compounds
• Potassium and iodine
• Magnesium and chlorine
• Aluminum and nitrogen
• Sodium and sulfur
• Name the following binary compounds
• AgCl
• ZnO
• CaBr2
• CaCl2

21
Your Turn II

• Write the formula for the following compounds
• Cu2 and Br
• Fe2 and O-2
• Pb2 and Cl
• Hg2 and Sn2-
• Give the names for the following compounds
• CuO
• CoF3
• SnI4
• FeS

22
Compounds Containing Polyatomic Ions

• Many common polyatomic ions are
  oxyanionspolyatomic ions that contain oxygen.
• Some elements can combine with oxygen to form
  more than one type of oxyanion.

• example nitrogen can form or
  .

• The name of the ion with the greater number of
  oxygen atoms ends in -ate. The name of the ion
  with the smaller number of oxygen atoms ends in
  -ite.

nitrate nitrite
23
Naming Binary Compounds

• Some elements can form more than two types of
  oxyanions.
• example chlorine can form , ,
  or
• In this case, an anion that has one fewer oxygen
  atom than the -ite anion has is given the prefix
  hypo-.
• An anion that has one more oxygen atom than the
  -ate anion has is given the prefix per-.

hypochlorite chlorite
chlorate perchlorate
24
Polyatomic Ions
25
Naming Compounds with Polyatomic Ions
26
Understanding Formulas for Polyatomic Ionic
Compounds
27
Example

• Write the formula for tin(IV) sulfate.

Sn(SO4)2
28
Your Turn III

• Write the names for the following compounds
• Ag2O
• Ca(OH)2
• KClO3
• Fe3(CrO4)2
• NH4OH
• Write the Formulas for the following compounds
• Sodium iodide
• Copper (II) sulfate
• Potassium perchlorate
• Lithium nitrate
• Sodium carbonate

29
Naming Binary Molecular Compounds

• Unlike ionic compounds, molecular compounds are
  composed of individual covalently bonded units,
  or molecules.
• As with ionic compounds, there is also a Stock
  system for naming molecular compounds. (more on
  this later)
• The old system of naming molecular compounds is
  based on the use of prefixes.
• examples
• CCl4 carbon tetrachloride (tetra- 4)
• CO carbon monoxide (mon- 1)
• CO2 carbon dioxide (di- 2)

30
Prefixes for Naming Covalent Compounds
31
Rules for the Prefix System

• The element with the smaller group number is
  usually given first. If both are in the same
  group the one with the greater period goes first.
• The first element is given a prefix only if it
  contributes more than one atom.
• The second element is named by combining three
  things
• A prefix indicating the number of atoms
• The root of the name
• The ending ide.

32
Example

• Give the name for As2O5.
• Diarsenic pentoxide
• Write the formula for oxygen difluoride.
• OF2

33
Your Turn IV

• Write the formulas for the following compounds
• Carbon tetraiodide
• Dinitrogen trioxide
• Sulfur hexafluoride
• Phosphorus trichloride
• Xeon tetrafluoride
• Name the following binary molecular compounds
• SO3
• ICl3
• PBr5
• N2O4
• PF4

34
Covalent-Network Compounds

• Some covalent compounds do not consist of
  individual molecules.
• Instead, each atom is joined to all its neighbors
  in a covalently bonded, three-dimensional
  network.
• Subscripts in a formula for covalent-network
  compound indicate smallest whole-number ratios of
  the atoms in the compound.
• examples
• SiC, silicon carbide
• SiO2, silicon dioxide
• Si3N4, trisilicon tetranitride

35
Acids and Salts

• An acid is a certain type of molecular compound.
  Most acids used in the laboratory are either
  binary acids or oxyacids.
• Binary acids are acids that consist of two
  elements, usually hydrogen and a halogen.
• Oxyacids are acids that contain hydrogen, oxygen,
  and a third element (usually a nonmetal).

36
Acids and Salts

• In the laboratory, the term acid usually refers
  to a solution in water of an acid compound rather
  than the acid itself.
• example hydrochloric acid refers to a water
  solution of the molecular compound hydrogen
  chloride, HCl
• Many polyatomic ions are produced by the loss of
  hydrogen ions from oxyacids.

sulfuric acid H2SO4 sulfate
nitric acid HNO3 nitrate
phosphoric acid H3PO4 phosphate
37
Acids and Salts

• An ionic compound composed of a cation and the
  anion from an acid is often referred to as a
  salt.
• examples
• Table salt, NaCl, contains the anion from
  hydrochloric acid, HCl.
• Calcium sulfate, CaSO4, is a salt containing the
  anion from sulfuric acid, H2SO4.
• The bicarbonate ion, , comes from
  carbonic acid, H2CO3.

38
Naming Acids

• If the anion attached to hydrogen ends in -ide,
  put the prefix hydro- and change -ide to -ic acid
• HCl - hydrogen ion and chloride ion
• hydrochloric acid
• H2S hydrogen ion and sulfide ion
• hydrosulfuric acid

39
Naming Acids

• If the anion has oxygen in it
• it ends in -ate of -ite
• change the suffix -ate to -ic acid
• HNO3 Hydrogen and nitrate ions
• Nitric acid
• change the suffix -ite to -ous acid
• HNO2 Hydrogen and nitrite ions
• Nitrous acid

40
Your Turn V

• Name these acids
• HF
• H3P
• H2SO4
• H2SO3
• HCN
• H2CrO4

41
Your Turn V

• Write the formulas for the following acids
• Perchloric acid
• Hydrobromic acid
• Chlorous acid
• Phosphoric acid
• Carbonic acid

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43
Oxidation Numbers

• Section 2

44
Oxidation Numbers

• The charges on the ions in 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
  are assigned to the atoms composing the compound
  or ion.
• Unlike ionic charges, oxidation numbers do not
  have an exact physical meaning rather, they
  serve as useful bookkeeping devices to help
  keep track of electrons.

45
Assigning Oxidation Numbers

• In general when assigning oxidation numbers,
  shared electrons are assumed to belong to the
  more electronegative atom in each bond.
• More-specific rules are provided by the following
  guidelines.
• The atoms in a pure element have an oxidation
  number of zero.
• examples all atoms in sodium, Na, oxygen, O2,
  phosphorus, P4, and sulfur, S8, have oxidation
  numbers of zero.

46
Assigning Oxidation Numbers

• The more-electronegative element in a binary
  compound is assigned a negative number equal to
  the charge it would have as an anion. Likewise
  for the less-electronegative element.
• Fluorine has an oxidation number of 1 in all of
  its compounds because it is the most
  electronegative element.

47
Assigning Oxidation Numbers

• Oxygen usually has an oxidation number of 2.
• Exceptions
• In peroxides, such as H2O2, oxygens oxidation
  number is 1.
• In compounds with fluorine, such as OF2, oxygens
  oxidation number is 2.
• Hydrogen has an oxidation number of 1 in all
  compounds containing elements that are more
  electronegative than it it has an oxidation
  number of 1 with metals.

48
Assigning Oxidation Numbers

• The algebraic sum of the oxidation numbers of all
  atoms in an neutral compound is equal to zero.
• The algebraic sum of the oxidation numbers of all
  atoms in a polyatomic ion is equal to the charge
  of the ion.
• Although rules 1 through 7 apply to covalently
  bonded atoms, oxidation numbers can also be
  applied to atoms in ionic compounds similarly.

49
Assigning Oxidation Numbers

• Monatomic ions have a oxidation number equal to
  the charge of their ion.
• Na - oxidation number 1
• Ca2 - oxidation number 2
• Cl- - oxidatin number -1

50
Example

• Assign oxidation numbers to each atom in the
  following compounds
• UF6
• H2SO4
• ClO3-

51
Solution

• Fluorine always has -1, so U must have 6 to
  equal the -6 from the fluorines.
• Hydrogen is 1 so SO4 is -2.
• Overall charge must be -1. Oxygen has an
  oxidation number of -2 which give a total of -6,
  so Cl must be 5 to give a total of -1.

52
Your Turn VI

• Assign oxidation numbers to the following
  compounds or ions
• HCl
• CF4
• PCl3
• SO2
• HNO3
• KH
• P4O10
• HClO3

53
Using Oxidation Numbers for Formulas and Names

• As shown in the table in the next slide, many
  nonmetals can have more than one oxidation
  number.
• These numbers can sometimes be used in the same
  manner as ionic charges to determine formulas.
• example What is the formula of a binary compound
  formed between sulfur and oxygen?
• From the common 4 and 6 oxidation states of
  sulfur, you could predict that sulfur might form
  SO2 or SO3.
• Both are known compounds.

54
Common Oxidation States of Nonmetals
55
Using Oxidation Numbers for Formulas and Names

• Using oxidation numbers, the Stock system,
  introduced in the previous section for naming
  ionic compounds, can be used as an alternative to
  the prefix system for naming binary molecular
  compounds.

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57
Using Chemical Formulas

• Section 3

58
Remember

• A chemical formula indicates
• the elements present in a compound
• the relative number of atoms or ions of each
  element present in a compound
• Chemical formulas also allow chemists to
  calculate a number of other characteristic values
  for a compound
• formula mass
• molar mass
• percentage composition

59
Formula Masses

• The formula mass of any molecule, formula unit,
  or ion is the sum of the average atomic masses of
  all atoms represented in its formula.
• exampleformula mass of water, H2O
• average atomic mass of H 1.01 amu
• average atomic mass of O 16.00 amu

average mass of H2O molecule 18.02 amu
60
Formula Masses

• The mass of a water molecule can be referred to
  as a molecular mass.
• The mass of one formula unit of an ionic
  compound, such as NaCl, is not a molecular mass.
• The mass of any unit represented by a chemical
  formula (H2O, NaCl) can be referred to as the
  formula mass.

61
Example

• Find the formula mass of potassium chlorate, KClO3

formula mass of KClO3 122.55 amu
62
Your Turn VII

• Find the formula mass of each of the following
• H2SO4
• Ca(NO3)2
• PO43-
• MgCl2

63
Molar Masses

• The molar mass of a substance is equal to the
  mass in grams of one mole, or approximately 6.022
  ? 1023 particles, of the substance.
• example the molar mass of pure calcium, Ca, is
  40.08 g/mol because one mole of calcium atoms has
  a mass of 40.08 g.
• The molar mass of a compound is calculated by
  adding the masses of the elements present in a
  mole of the molecules or formula units that make
  up the compound.

64
Molar Mass

• One mole of water molecules contains exactly two
  moles of H atoms and one mole of O atoms. The
  molar mass of water is calculated as follows.
• molar mass of H2O molecule 18.02 g/mol
• A compounds molar mass is numerically equal to
  its formula mass.

65
Calculating Molar Masses for Ionic Compounds
66
Example

• What is the molar mass of barium nitrate,
  Ba(NO3)2?

molar mass of Ba(NO3)2 261.35 g/mol
67
Your Turn VIII

• For each of the following compounds tell how many
  mole of each atom there are and determine the
  molar mass.
• Al2S3
• NaNO3
• Ba(OH)2
• K2SO4
• (NH4)2CrO4

68
Molar Mass as a Conversion Factor

• The molar mass of a compound can be used as a
  conversion factor to relate an amount in moles to
  a mass in grams for a given substance.
• To convert moles to grams, multiply the amount in
  moles by the molar mass
• Amount in moles ? molar mass (g/mol) mass in
  grams

69
Mole-Mass Calculations
70
Example

• What is the mass in grams of 2.50 mol of oxygen
  gas?
• moles O2 grams O2
• amount of O2 (mol) ? molar mass of O2 (g/mol)
  mass of O2 (g)

71
Solution
Use the molar mass of O2 to convert moles to mass.
72
Converting Between Amount in Moles and Number of
Particles
73
Your Turn IX

• What is the mass in grams of 3.04 moles of NH3?
• Calculate the mass of 0.257 mol of Ca(NO3)2
• How many grams are there in 4.33 mol of H2SO4?

74
Example

• Ibuprofen, C13H18O2, is the active ingredient in
  many nonprescription pain relievers. Its molar
  mass is 206.31 g/mol.
• If the tablets in a bottle contain a total of 33
  g of ibuprofen, how many moles of ibuprofen are
  in the bottle?
• How many molecules of ibuprofen are in the
  bottle?
• What is the total mass in grams of carbon in 33 g
  of ibuprofen?

75
Solution
a.
b.
c.
76
Your Turn X

• How many moles of compound are there in the
  following
• 6.60 g (NH4)2SO4
• 4.5 kg Ca(OH)2
• How many molecules are there in the following
• 25.0 g H2SO4
• 125 g of C12H22O12
• What is the mass in grams of 6.25 mol of copper
  (II) nitrate?

77
Percentage Composition

• It is often useful to know the percentage by mass
  of a particular element in a chemical compound.
• To find the mass percentage of an element in a
  compound, the following equation can be used.

• The mass percentage of an element in a compound
  is the same regardless of the samples size.

78
Percentage Composition

• The percentage of an element in a compound can be
  calculated by determining how many grams of the
  element are present in one mole of the compound.

• The percentage by mass of each element in a
  compound is known as the percentage composition
  of the compound.

79
Percentage Composition of Iron Oxides
80
Percentage Composition Calculations
81
Example

• Find the percentage composition of copper(I)
  sulfide, Cu2S.

Molar mass of Cu2S 159.2 g
82
Solution
83
Your Turn XI

• Find the percentage of water in the hydrate
  Na2CO3?10H2O.
• Find the percent composition in the following
• PbCl2
• Ba(NO3)2
• Find the percent of water in ZnSO4?7H2O
• Magnesium hydroxide is 54.87 oxygen by mass.
  How many grams of oxygen are in 175 g of the
  compound? How many moles of oxygen is this?

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85
Determining Chemical Formulas

• Section 4

86
Empirical and Actual Formulas

• 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.
• For an ionic compound, the formula unit is
  usually the compounds empirical formula.
• For a molecular compound, however, the empirical
  formula does not necessarily indicate the actual
  numbers of atoms present in each molecule.
• example the empirical formula of the gas
  diborane is BH3,
• but the molecular formula is B2H6.

87
Empirical and Actual Formulas
88
Calculation of Empirical Formulas

• To determine a compounds empirical formula from
  its percentage composition, begin by converting
  percentage composition to a mass composition.
• Assume that you have a 100.0 g sample of the
  compound.
• Then calculate the amount of each element in the
  sample.
• example diborane
• The percentage composition is 78.1 B and 21.9
  H.
• Therefore, 100.0 g of diborane contains 78.1 g of
  B and 21.9 g of H.

89
Calculation of Empirical Formulas

• Next, the mass composition of each element is
  converted to a composition in moles by dividing
  by the appropriate molar mass.

• These values give a mole ratio of 7.22 mol B to
  21.7 mol H.

90
Calculation of Empirical Formulas

• To find the smallest whole number ratio, divide
  each number of moles by the smallest number in
  the existing ratio.

• Because of rounding or experimental error, a
  compounds mole ratio sometimes consists of
  numbers close to whole numbers instead of exact
  whole numbers.
• In this case, the differences from whole numbers
  may be ignored and the nearest whole number taken.

91
Example

• Quantitative analysis shows that a compound
  contains 32.38 sodium, 22.65 sulfur, and 44.99
  oxygen. Find the empirical formula of this
  compound.

92
Solution
93
Solution
Smallest whole-number mole ratio of atoms The
compound contains atoms in the ratio 1.408 mol
Na0.7063 mol S2.812 mol O.
Rounding yields a mole ratio of 2 mol Na1 mol
S4 mol O. The empirical formula of the compound
is Na2SO4.
94
Your Turn XII

• Determine the empirical formula of the compound
  that contains 17.15 carbon, 1.44 hydrogen, and
  81.41 fluorine.
• Determine the empirical formula of a compound
  that contains 36.70 potassium, 33.27 chlorine,
  and 30.03 oxygen.
• Analysis of a 10.150 g sample of a compound known
  to contain only phosphorus and oxygen indicates a
  phosphorus content of 4.433 g. What is the
  empirical formula of this compound?
• An 20.0 g sample contains only calcium and
  bromine, of that 4.00 g of calcium are present.
  What is the empirical formula of the compound?

95
Calculation of Molecular Formulas

• The empirical formula contains the smallest
  possible whole numbers that describe the atomic
  ratio.
• The molecular formula is the actual formula of a
  molecular compound.
• An empirical formula may or may not be a correct
  molecular formula.
• The relationship between a compounds empirical
  formula and its molecular formula can be written
  as follows.
• x(empirical formula) molecular formula

96
Calculation of Molecular Formulas

• The formula masses have a similar relationship.
• x(empirical formula mass) molecular formula
  mass
• To determine the molecular formula of a compound,
  you must know the compounds formula mass.
• Dividing the experimental formula mass by the
  empirical formula mass gives the value of x.
• A compounds molecular formula mass is
  numerically equal to its molar mass, so a
  compounds molecular formula can also be found
  given the compounds empirical formula and its
  molar mass.

97
Comparing Empirical and Molecular Formulas
98
Example

• 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
  283.89 g/mol. What is the compounds molecular
  formula?
• x(empirical formula) molecular formula

99
Solution

• Molecular formula mass is numerically equal to
  molar mass.
• molecular molar mass 283.89 g/mol
• molecular formula mass 283.89 amu
• empirical formula mass
• mass of phosphorus atom 30.97 amu
• mass of oxygen atom 16.00 amu
• empirical formula mass of P2O5
• (2 ? 30.97 amu ) (5 16.00 amu ) 141.94 amu

100
Solution
2 x (P2O5) P4O10 The compounds molecular
formula is therefore P4O10.
101
Your Turn XIII

• Determine the molecular formula of the compound
  with an empirical formula of CH and a formula
  mass of 78.110 amu.
• A sample of a compound with a formula mass of
  34.00 amu is found to consist of 0.44 g H and
  6.92 g of O. Find its molecular formula.
• The empirical formula for trichloroisocyanuric
  acid, the active ingredient in bleach, is OCNCl.
  The molar mass of this compound is 232.41 g/mol.
  What is the molecular formula of
  trichloroisocyanuric.

102
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