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Molecules and Compounds
Compounds Display Constant
Composition
• Elements combine in fixed Proportions
• Law of Constant Composition
– H2O
– CO2
Molecules and Compounds
Compounds Display Constant
Composition
If we decompose water, we find 16.0 grams of
oxygen to every 2.00 grams of hydrogen. This
produces an Oxygen to Hydrogen mass ratio of
8.0.
Mass ratio = 16.0 grams = 8.0
2.00 grams
Water has a constant Mass Ratio of Oxygen to
Hydrogen of 8.0.
Molecules and Compounds
Chemical Formulas: How to
Represent Compounds
• Chemical Formula
– Elements present
– Relative numbers
• Subscripts
– Part of the definition
– Changing subscript changes compound
• Metals – Left-hand side of Periodic Table
– Listed first
• Nonmetals – Right-hand side of the Periodic Table
– Those to the left– More Metal-like
– Listed first
Significance of Chemical
Formulas
• 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
Molecules and Compounds
Writing Chemical Formulas
• Compound containing 2 aluminum atoms to
every 3 oxygen atoms
– Al2O3
• Compound containing 3 oxygen atoms to
every 1 sulfur atom
– SO3
Molecules and Compounds
Chemical Formulas: How to
Represent Compounds
• Water – H2O
– Hydrogen and
oxygen atoms
– 2:1 ratio
– Symbol and
subscript
– Dihydrogen
monoxide
Chemical Formulas and Compounds
Chemical Formulas: How to
Represent Compounds
• Carbon dioxide – CO2
–
–
–
–
–
Carbon and oxygen atoms
1:2 ratio
Symbol and subscript
Carbon dioxide
Two nonmetals
• Left-most listed first
• One carbon – not mono
• The chemical formula for an ionic
compound represents one formula unit—
the simplest ratio of the compound’s
positive ions (cations) and its negative
ions (anions).
– example: aluminum sulfate — Al2(SO4)3
– Parentheses surround the polyatomic ion
a unit. The subscript 3 refers to the unit.
SO42- to identify it as
Molecules and Compounds
Writing Chemical Formulas
Polyatomic Ions
• Contain several identical groupings of
atoms
• Polyatomic Ions
• Parenthesis
• Mg(NO3)2
• Magnesium nitrate
Molecules and Compounds
Writing Chemical Formulas
Polyatomic Ions
•
•
•
•
Mg(NO3)2
1 Mg
2N
6O
Molecules and Compounds
Molecular View of
Elements and Compounds
• Atomic Elements
– Single atoms
– Most Elements
• Molecular Elements
– Diatomic atoms
– Two atoms bonded together
– H2, N2, O2, F2, Cl2, Br2, I2
• Molecular Compounds
– Two or more Nonmetals
• Ionic Compounds
– Metal and nonmetal
Molecules and Compounds
Molecular View of
Elements and Compounds
Molecules and Compounds
Writing Formulas for
Ionic Compounds
•
•
•
•
Charge Neutral
Most elements form only one type of ion
Predictable charge
Sodium chloride
– Na  +1
– Cl  –1
– NaCl
• Magnesium chloride
– Mg  +2
– Cl  –1
– MgCl2
Common Monatomic Ions
(pg.221 in your textbook)
Common Monatomic Ions…
Molecules and Compounds
Writing Formulas for
Ionic Compounds
1.
2.
3.
4.
5.
Write the symbol for the metal and its charge
Write the symbol for the nonmetal and its
charge
Charge (without sign) becomes subscript for
other ion
Reduce subscripts to smallest whole number
ratio
Check that the sum of the charges of the cation
cancels the sum of the anions
Molecules and Compounds
Writing Formulas for
Ionic Compounds
Ionic compound between aluminum and oxygen
Al3+
O2–
Al2O3
2 (+3) = +6
3 (–2) = –6
Molecules and Compounds
Writing Formulas for
Ionic Compounds
Ionic compound between magnesium and oxygen
Mg2+
O2–
MgO
1 (+2) = +2
1 (–2) = –2
Smallest whole number ratio is 1:1
Molecules and Compounds
Naming Ionic Compounds
• Metal and Nonmetal
• Two Types
– Type I
• Metal has constant predictable charge
• Inferred from group number in Periodic Table
– Type II
• Charge NOT always the same
• Transition Metals
Molecules and Compounds
Naming Ionic Compounds
Type I Binary Ionic Compounds
• Binary compounds
– Two different kinds of elements
– Cation + anion with IDE
• Metal
– Predictable constant charge
– Name of metal
• Nonmetal
– Name of nonmetal
– Ending  IDE
Molecules and Compounds
Naming Ionic Compounds
Type I Binary Ionic Compounds
• Binary compounds
– Two different kinds of elements
– Cation + anion with IDE
• Metal
– Predictable constant charge
– Name of metal
• Nonmetal
– Name of nonmetal
– Ending  IDE
Molecules and Compounds
Naming Ionic Compounds
Type I Binary Ionic Compounds
• MgF2
– Metal  Magnesium
– Nonmetal  Fluoride
– Magnesium fluoride
• KBr
– Metal  Potassium
– Nonmetal  Bromine
– Potassium bromide
Molecules and Compounds
Naming Ionic Compounds
Type I Binary Ionic Compounds
• CaCl2
– Metal  Calcium
– Nonmetal  Chlorine
– Calcium chloride
• Na2O
– Metal  Sodium
– Nonmetal  Oxygen
– Sodium oxide
Molecules and Compounds
Naming Ionic Compounds
Type II Binary Ionic Compounds
• Binary compounds
– Two different kinds of elements
– Cation + anion with IDE
• Metal
– Charge NOT always the same
– Transition metals and roman
numeral for charge
– Name of metal
• Nonmetal
– Name of nonmetal
– Ending  IDE
Molecules and Compounds
Naming Ionic Compounds
Type II Binary Ionic Compounds
• FeCl3
–
–
–
–
Metal  Iron
Nonmetal  Chlorine
Charge on Iron must be +3
Iron(III) chloride
• CrO
–
–
–
–
Metal  Chromium
Nonmetal  Oxygen
Charge on Chromium must be +2
Chromium(II) oxide
Molecules and Compounds
Naming Ionic Compounds
Type II Binary Ionic Compounds
• PbCl4
– Metal  Lead
– Nonmetal  Chlorine
– Charge on Lead must be +4
– Lead(IV) chloride
• Must determine charge of Cation
from the Formula
Molecules and Compounds
Naming Ionic Compounds
Polyatomic Ions
• Group of atoms with an overall charge
• Common Polyatomic Ions on pg. 226 in
textbook. MEMORIZE (flashcards?)
Molecules and Compounds
Naming Ionic Compounds
Polyatomic Ions
• KNO3
– Metal  Potassium
– Nonmetal  Nitrate
– Potassium nitrate
• FeSO4
– Metal  Iron
– Nonmetal  Sulfate
– Iron(II) sulfate
Ionic Compounds Containing
Polyatomic Ions….
• Write the formula for tin(IV) sulfate
• Sn(SO4)2
Molecules and Compounds
Naming Molecular Compounds
• Molecular Compounds
•Prefixes
• Two or more nonmetals
Mono – 1
• Right of the Periodic Table
Di – 2
Tri – 3
• Binary Molecular Compounds
Tetra – 4
Penta – 5
• Two elements
Hexa – 6
Hepta – 7
• Most “metal-like” first
Octa – 8
• Prefix [Element 1] Prefix [Element 2]
Nona – 9
Deca - 10
• Do not begin with “mono” and if
prefix is two syllables (& element
starts with vowel) then drop the
ending vowel of prefix.
Molecules and Compounds
Nomenclature: Naming
Compounds
• Systematic ways to name compounds
• Common names
• Water
– H2O
– Common name  Water
– Systematic name  Dihydrogen monoxide
Molecules and Compounds
Naming Molecular Compounds
• CO2
• Most “metal-like” – Carbon
• Nonmetal – Oxygen
• Carbon dioxide
• N2O
• Most “metal-like” – Nitrogen
• Nonmetal – Oxygen
• Dinitrogen monoxide
• Give name for As2O5:
• Write formula for oxygen difluoride:
Molecules and Compounds
Naming Acids
•Acids
–Molecular Compounds
–Form H+
–Sour Taste
•Binary Acids- acids that consist of
two elements (usually hydrogen
and a halogen)
•Oxyacids – acids that contain
hydrogen, oxygen and a third
element (usually a nonmetal)
Binary Acids…
If anion ends in “ide” then acid name is
“hydro—ic”
• HCl hydrochloric acid
• HBr hydrobromic acid
• HI hydroiodic acid
Naming Oxyacids
• No “hydro” in name
• If anion ends in “ate” then acid name is “—
ic”
• If anion ends in “ite” then acid name is “—
ous”
Oxyacids
Anion
sulfate SO42sulfite SO32nitrate NO3phosphate PO43-
Acid _____
sulfuric acid H2SO4
sufurous acid H2SO3
nitric acid HNO3
phosphoric acid H3PO4
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.
Molecules and Compounds
Nomenclature Summary
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.
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.
1. 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.
Assigning Oxidation Numbers,
continued..
2. 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.
3. Fluorine has an oxidation number of –1
in all of its compounds because it is
the most electronegative element.
Assigning Oxidation Numbers,
continued..
4. Oxygen usually has an oxidation number of –2.
– Exceptions:
• In peroxides, such as H2O2,
oxygen’s oxidation number is –1.
• In compounds with fluorine, such as OF2, oxygen’s oxidation
number is +2.
5. 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.
Assigning Oxidation Numbers,
continued..
6. The algebraic sum of the oxidation numbers
of all atoms in an neutral compound is equal
to zero.
7. The algebraic sum of the oxidation numbers
of all atoms in a polyatomic ion is equal to
the charge of the ion.
8. Although rules 1 through 7 apply to
covalently bonded atoms, oxidation
numbers can also be applied to atoms in
ionic compounds similarly.
Assigning Oxidation Numbers,
continued..
Assign oxidation numbers to each atom
in the following compounds or ions:
a. UF6
b. H2SO4
c. ClO3-
Molecules and Compounds
Formula Mass & Molar Mass
• Formula Mass - Average mass of the
Molecules that compose a compound
(in a.m.u.’s)
• Molar Mass – Average mass of the
molecules that compose a mole
(6.02x1023) of the compound (in
grams/mole)
Same values different units!
Molar Mass Calculation
What is the molar mass of barium nitrate,
Ba(NO3)2?
• 261.35 g/mol
What is the mass in grams of 2.50 mol of
oxygen gas (O2)?
• 80.0g
Molar Mass as a Conversion
Factor…
Ibuprofen, C13H18O2, is the active ingredient
in many nonprescription pain relievers. Its
molar mass is 206.31 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?
1 mol C13H18O2
33 g C13H18O2 
 0.16 mol C13H18O2
206.31 g C13H18O2
6.022  1023 molecules
0.16mol C13H18O2 

mol
9.6  1022 molecules C13H18O2
13 mol C
12.01 g C
0.16 mol C13H18O2 

 25 g C
mol C13H18O2
mol C
Percentage
Chapter
7
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. mass of element in sample of compound
mass of sample of compound
 100 
% element in compound
• The mass percentage of an element in a
compound is the same regardless of the sample’s
size.
Percentage Composition of Iron
Oxides
Percentage Composition
Calculations
Chapter
7
Percentage
Composition,
continued
Find the percentage composition of
copper(I) sulfide, Cu2S.
63.55 g Cu
2 mol Cu 
 127.1 g Cu
mol Cu
1 mol S 
32.07 g S
 32.07 g S
mol S
Molar mass of Cu2S = 159.2 g
127.1 g Cu
 100  79.85% Cu
159.2 g Cu2S
32.07 g S
 100  20.15% S
159.2 g Cu2S
Chapter 7
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 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.
– example: the empirical formula of the gas
diborane is BH3, but the molecular formula is
B2H6.
Chapter 7
Calculation
of Empirical Formulas
• To determine a compound’s 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.
Calculation
of Empirical Formulas,
Chapter 7
continued
• Next, the mass composition of each element
is converted to a composition in moles by
dividing by the appropriate molar mass.
78.1 g B 
1 mol B
 7.22 mol B
10.81 g B
1 mol H
21.9 g H 
 21.7 mol H
1.01 g H
• These values give a mole ratio of 7.22
mol B to 21.7 mol H.
Chapter 7
Calculation
of Empirical Formulas,
continued
• To find the smallest whole number ratio, divide
each number of moles by the smallest number in
the existing ratio.
7.22 mol B 21.7 mol H
:
 1 mol B : 3.01 mol H
7.22
7.22
• Because of rounding or experimental error, a
compound’s 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.
Chapter
7
Calculation
of Empirical
Formulas, continued
Quantitative analysis shows that a
compound contains 32.38% sodium, 22.65%
sulfur, and 44.99% oxygen. Find the
empirical formula of this compound.
32.38 g Na 
1 mol Na
 1.408 mol Na
22.99 g Na
22.65 g S 
1 mol S
 0.7063 mol S
32.07 g S
1 mol O
44.99 g O 
 2.812 mol O
16.00 g O
1.408 mol Na 0.7063 mol S 2.812 mol O
:
:

0.7063
0.7063
0.7063
1.993 mol Na : 1 mol S : 3.981 mol O
Chapter 7 of Molecular Formulas
Calculation
• 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 compound’s
empirical formula and its molecular formula:
x(empirical formula) = molecular formula
Chapter
7
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 compound’s formula mass.
– Dividing the experimental formula mass by the
empirical formula mass gives the value of x.
• A compound’s molecular formula mass is numerically
equal to its molar mass, so a compound’s molecular
formula can also be found given the compound’s
empirical formula and its molar mass.
Calculation of Molecular Formulas
In Sample Problem M, 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 compound’s molecular
formula?
283.89 amu
x=
 2.0001
141.94 amu
The compound’s molecular formula is therefore P4O10.