Unit 2: Types of
Chemical Equations
Fall 2013
Ms. Nickerson
Unit 1 Outline
Page 1
 Taking measurements in Lab
 Uncertainty in Measurement
 The Mole & Molar Mass
 Percent Composition
 Empirical and Molecular Formulas
 Chemical Formulas & Nomenclature
 Polyatomic ions
Unit 1 Outline
Page 2
 Writing chemical equations and drawn representations
 Balancing chemical equations
 Stoichiometry: Applying mole concept to chemical
equations
 Limiting Reactants
 Theoretical yield & Percent (%) yield
 Significant Figures
Unit 1 Outline
Page 3
 Classification of matter
 Law of definite proportions
 Chemical and physical changes
 Separation of mixtures
Ionic Compounds in Water
 Electrolyte – A substance that splits apart into ions
when dissolved in water (aqueous)
 Soluble ionic compounds
 Nonelectrolyte – A substance that does not split into
ions when dissolved in water (aq)
 Insoluble ionic compounds & soluble molecules (covalent
bonding)
Ionic Compounds in Water
 Solvation – The process of separating ions in water
 The water molecules surround the ions to stabilize the ions
(prevents them from recombining)
Electrolytes
 Strong electrolytes – exist completely (or almost
completely) as ions in water
 Weak electrolytes – exist mostly in the form of neutral
molecules (a small fraction may form ions)
Precipitation Reactions
 Precipitation Reactions – reactions that result in the
formation of an insoluble product called a precipitate
 One of the products comes out of solution as a solid
 Use solubility rules to determine which products will form as
precipitates
Double Replacement Reations
 Two compounds react to form two new compounds. No
changes in oxidation numbers occur.
 The ion pairs switch, and the products each contain a
cation (+) and an anion (-).
 All double replacement reactions must have a "driving
force" that removes a pair of ions from solution.
Precipitation Reactions
 A precipitate is an insoluble substance formed by the reaction
of two aqueous substances. Two ions bond together so strongly
that water can not pull them apart. You must know your
solubility rules to write these net ionic equations
 Ex. Solutions of silver nitrate and lithium bromide are mixed.
Ag+(aq) + Br-(aq)  AgBr(s)
Precipitation Reactions
 Formation of a gas: Gases may form directly in a double
replacement reaction or can form from the decomposition of a
product such as H2CO3 or H2SO3.
 Common gases: CO2, SO2, SO3, H2S, NO2, NH3, O2, H2
 Ex. Excess hydrochloric acid solution is added to a solution of
potassium sulfite.
H+(aq) + SO32-(aq)  H20(l) + SO2(g)
Precipitation Reactions
 Formation of a gas: Gases may form directly in a double
replacement reaction or can form from the decomposition of a
product such as H2CO3 or H2SO3.
 Common gases: CO2, SO2, SO3, H2S, NO2, NH3, O2, H2
 Ex. A solution of sodium hydroxide is added to a solution of
ammonium chloride.
OH-(aq) + NH4+(aq) NH3(g) + H2O(l)
Precipitation Reactions
 Formation of a molecular substance: When a molecular
substance such as water or acetic acid is formed, ions are
removed from solution and the reaction "works".
 Ex. Dilute solutions of lithium hydroxide and hydrobromic acid
are mixed.
OH-(aq) + H+(aq)  H2O(l)
(HBr, HCI, and HI are strong acids)
Precipitation Reactions
 Formation of a molecular substance: When a molecular
substance such as water or acetic acid is formed, ions are
removed from solution and the reaction "works".
 Ex. Gaseous hydrofluoric acid reacts with solid silicon dioxide.
HF(g) + SiO2(s)  SiF4(l) + H2O(l)
PREDICTING THE PRODUCTS OF DOUBLE REPLACEMENT
REACTIONS
 Follow the steps to predict the products and balance the final equation:
 Predict the identity of the precipitate that forms when
aqueous solutions of BaCl2 and K2SO4 are mixed.
 Write the balanced chemical equation for the reaction.
 What compound precipitates when aqueous solutions of
Fe2(SO4)3 and LiOH are mixed? Write a balanced equation for
the reaction.
NET IONIC EQUATIONS WITH DOUBLE REPLACEMENT
REACTIONS
Write the net ionic equation for the reaction between aqueous
solutions of Fe2(SO4)3 and LiOH.
 Molecular:
 Complete Ionic:
 Net Ionic:
NET IONIC EQUATIONS WITH DOUBLE REPLACEMENT
REACTIONS
 Write the net ionic equation for the reaction between
HCl(aq) and Pb(NO3)2.
 Molecular:
 Complete Ionic:
 Net Ionic:
NET IONIC EQUATIONS WITH DOUBLE REPLACEMENT
REACTIONS
 Sodium carbonate solution is combined with calcium
hydroxide solution. What is the net ionic equation for this
reaction?
 Molecular:
 Complete Ionic:
 Net Ionic:
 Will a precipitate form when solutions of Ba(NO3)2 and KOH
are mixed?
 Will a precipitate form when solutions of LiBr and NH4OH
are mixed?
SOLUTION STOICHIOMETRY
 Looking for the molarity of an ion WITHIN a
compound? Multiply the ions by their
subscripts.
EX:
1.0 M CaCl2 = 1.0 M Ca2+ ions AND
2.0 M Cl1- ions
SOLUTION STOICHIOMETRY
DIMENSIONAL ANALYSIS SETUP
 What is the concentration of lithium ions are in
3.5 M Li2SO4?
 How many sulfate ions are in 0.12 M Li2SO4?
Use the molarity of ions in your solution stoichiometry
problems:
 How many moles of Na+ ions are needed to
produce 5.0 grams of Pb(NO3)2?
 Molecular: Na2SO4(aq) + Pb(NO3)2(aq)  2NaNO3(aq) + PbSO4(s)
 Complete Ionic:
 Net Ionic:
Use the molarity of ions in your solution stoichiometry
problems:
 Set up your stoichiometry problem as usual, but
watch for where to take the problem (moles Na+):
 Net Ionic: Pb2+(aq) + SO42-(aq) PbSO4(s)
 Consider the reaction:
Li2S(aq) + Co(NO3)2(aq)  2LiNO3(aq) + CoS(s)
 What volume of 0.150 M Li2S is required to completely react with 125
mL of 0.250 M Co(NO3)2?
 a) A 55.0 mL sample of 0.102 M potassium sulfate solution is mixed
with how many mL of 0.114 M lead (II) acetate solution in order to
fully react all the potassium sulfate?
Molecular Equation:
 b) How much precipitate product will be produced if “55.0 mL sample
of 0.102 M potassium sulfate” is your given value?
 Molecular Equation:
Acids & Bases
 Acid--any compound that, on reaction with water,
produces an ion called the hydronium ion, H3O+
[or H+], and an anion (Arrhenius definition)
 Base--any compound that provides a hydroxide,
OH−, and a cation in water (Arrhenius definition)
 Ammonia, NH3 is an exception, so Brønsted-Lowry defined
it as a proton acceptor!!
Acid-Base Reactions
 Neutralization—when moles acid = moles base,
each is neutralized [pH is not necessarily 7.0].
 Neutralization means your limiting reagent has been used
up.
 Usually involves an aqueous solution of acid and an
aqueous solution of base
 Titration – the process of neutralizing an acid or
base. You add one to the other (acid/base) until
you reach neutralization
Acid-Base Titrations
 Volumetric analysis—a technique for
determining the amount of a certain
substance by doing a titration
 Titrant—the substance delivered from a
buret so that its volume is accurately known
 Analyte—the substance being analyzed; its
mass or volume must also be accurately
known
 More terms to know:  Equivalence point – moles of OH− equals (is
equivalent to)
# moles of H3O+
 Indicator – undergoes a color change near the
equivalence point.
 Standardization – a procedure for establishing
the exact concentration of a reagent
 The products formed are a salt [ask yourself if it is soluble]
and water:
Neutralization:
Acid
+
HCl(aq) +
Base 
KOH(aq) 
Water
+
HOH(l) +
Salt
KBr(aq)
H+(aq) + Cl-(aq) + K+(aq) + OH-(aq)  HOH(l) + K+(aq) + Br-(aq)
H+(aq) + OH-(aq) 
HOH(l)
 What volume of a 0.100 M HC1 solution is needed to
neutralize 25.0 mL of 0.350 M NaOH ?
1. The titration of a 10.00 mL sample of an HCl solution of
unknown concentration requires 12.54 mL of a 0.100 M NaOH
solution to reach the equivalence point. What is the
concentration of the unknown HCl solution in M?
2. The titration of a 20.0 mL sample of an H2SO4 solution of unknown
concentration requires 22.87 mL of a 0.158 M KOH solution to
reach the equivalence point. What is the concentration of the
unknown H2SO4 solution?
Acid-Base Net Ionic Equations (NIE)
 H2SO4(aq) + Ca(OH)2(aq) 
 HClO4(aq) + KOH(aq) 
 HCl(aq) + Ba(OH)2(aq) 
 Write the balanced molecular and net ionic equations for the
reaction between hydrobromic acid (HBr) and potassium
hydroxide.
Electrolytes in Net Ionic Equations
 A salt is any ionic compound (cation, anion combination)
 Some salts are very soluble (strong electrolytes)
 Some salts are slightly soluble (weak electrolytes)
 Acids and bases are electrolytes because they ionize in
water (dissociate into ions)
 Strong acids & strong bases are strong electrolytes
 Weak acids & weak bases are weak electrolytes
Electrolytes in Net Ionic Equations
SIX Strong Acids
HCl – hydrochloric acid
HNO3 – nitric acid
HBr – hydrobromic acid
HClO4 – perchloric acid
HI – hydroiodic acid
H2SO4 – sulfuric acid
SEVEN Strong Bases
NaOH – sodium hydroxide
KOH – potassium hydroxide
Ca(OH)2 – calcium hydroxide
RbOH – rubidium hydroxide
Sr(OH)2 – strontium hydroxide
CsOH – cesium hydroxide
Ba(OH)2 – barium hydroxide
Strong Electrolytes
 Strong acids & strong bases
 Soluble salt (ionic compounds)
 Dissociates
100%
into its ions
 Therefore in a NIE, the compound can be written:
M+(aq) + X-(aq)
Weak Electrolytes
 Weak acids & strong bases
 Insoluble salt (ionic compounds)
 Does not
its ions
fully
dissociate into
 Therefore in a NIE, the compound can be written:
MX(aq)
*When the reaction occurs, the small amount that is dissociated
can react. We’ll talk about this later in the course.
Does it
start with
H?
Does it
contain
hydroxide
(OH-)?
Is it on the
strong
acid list?
STRONG
WEAK
(aq, split
ions)
(aq, keep
together)
Is it on the
strong
base list?
Is it soluble
according to
solubility
rules?
STRONG
STRONG
WEAK
(aq, split
ions)
(aq, split
ions)
(s, keep
together)
Write each compound as it would be written in a net ionic equation:
 HNO2: __________________
 HNO3: ___________________
 KOH: ___________________
 Fe(NO3)3: __________________
 KNO2: ___________________
 LiOH: ____________________
 HCl: ____________________
 HClO2: ____________________
 NH4Cl: __________________
 CoCl2: ____________________
 HF: ____________________
 NH4OH: ___________________
 H2CO3: __________________
 RbOH: ___________________
 Fe(OH)3: _________________
 H2SO3: ____________________
Write the NIE when each of the following pairs of aq soln’s are added together:
 A. HNO2(aq) + KOH(aq) 
 B. HCl(aq) + NH3(aq) 
Write the NIE when each of the following pairs of aq soln’s are added together:
 C. Na2CO3(aq) + HF(aq) 
 D. Fe(OH)3(s) + HNO3(aq) 
Write the NIE when each of the following pairs of aq soln’s are added together:
 E. HNO3(aq) + KOH(aq) 
*Reaction E is between a strong acid and a strong base.
 The net ionic equation between a strong acid and strong base is
always:
H+(aq) + OH-(aq)  H2O(l)
 Summarize a procedure for reading volume involving a
meniscus:
The Mole
 Chemists use moles to describe amounts of a substance.
1 mole = 6.02 x 1023 particles
1 mole = atomic mass in grams
Molar Mass
ATOMIC MASS CALCULATIONS
 1) The atomic masses of the elements are
based on Carbon-12 = 12 amu
 2) The atomic mass on the periodic table is
the weighted average of all the naturally
occurring isotopes of an element
“isotopes” have the
same #of ____, but
different # of _____
“weighted”
means it takes
into account the
% of each
isotope
Molar Mass
ATOMIC MASS CALCULATIONS
 1) The atomic masses of the elements are
based on Carbon-12 = 12 amu
 2) The atomic mass on the periodic table is
the weighted average of all the naturally
occurring isotopes of an element
 3) A mass spectrometer is an instrument that
can be used to separate the isotopes of an
element and find their natural abundances.
Converting with Moles
Mole-to-ply (multiply)
x
n
z
y
Molar Mass
 MOLAR MASS CALCULATIONS:
Molar mass of X = mass (g) of 1 mole
of X
• Look @ periodic table for masses
of each element.
• Multiply the mass by how many of
that element.
Molar Mass
 MOLAR MASS CALCULATIONS:
Examples:
C = 12.011 g/mol
CO2 = 1(12.011) + 2(16.00)
= 44.01 g/mol
Molar mass of X = mass (g) of 1 mole
of X
• Look @ periodic table for masses
of each element.
• Multiply the mass by how many of
that element.
Molar mass example 1
Find the molar mass of Ca5(PO4)3F, a compound which is
commonly used as a source of phosphate in fertilizers.
Molar mass example 2
Hydrates are ionic compounds, which have a set amount
of water contained in their crystal structure. Find the
molar mass of copper (II) sulfate pentahydrate.
Percent Composition
PERCENT COMPOSITION BY MASS:
Because compounds must have a
constant composition, it is possible to
find the percent by mass of each
element in the compound:
Total of
ALL that
element.
mass of element
% composition =
´100%
mass of compound
% Composition example 1
% composition =
mass of element
´100%
mass of compound
Ammonium is used extensively in the production of
fertilizer, where the ammonia is converted in ammonium
nitrate. Find the percent nitrogen in ammonium nitrate.
% Composition example 1
% composition =
mass of element
´100%
mass of compound
Calculate the percent water in strontium hydroxide octahydrate
Empirical and Molecular Formula
 An empirical formula is the simplest ratio of each
element in a compound (use GCF from math)
 A molecular formula gives the actual number of each
element in a compound.
EX: The empirical formula for hydrogen peroxide is HO,
but its molecular formula is H2O2.
EMPIRICAL FORMULA DETERMINATION
When in
doubt,
convert to
moles!
1) Convert grams to moles
2) Simplify moles by dividing all moles by smallest number
of calculated moles
3) If not a whole number ratio, multiple by the
appropriate value to create whole numbers
4) Insert ratio as subscripts
Assume % are grams, and the sample is
100 grams total (100%)
Empirical Formula example 1
 Methane is a hydrocarbon which is the main component
of natural gas. It is composed of 75% carbon and 25%
hydrogen. What is the empirical formula of methane?
Empirical Formula example 2
 Aluminum is a metal with a very high strength to
weight. It forms an oxide which is 52.9 % aluminum.
What is the empirical formula of the oxide?
Empirical Formula example 3
 When a hydrate of sodium carbonate is heated it loses
54.3% of its mass. What is the empirical formula of the
hydrate?
MOLECULAR FORMULA DETERMINATION: The molecular
formula is always an exact whole number multiple of the
empirical formula.
molecular weight
Multiplier # =
empirical weight
 EX: The empirical formula for hydrogen peroxide is HO
(weight = 17.01), but the molecular weight is 34.01.
34.01
=1.999 » 2
17.01
H2O2
Molecular Formula example 1
 A hydrocarbon has an empirical formula of CH2 and a
molecular mass of 56. What is the molecular formula of
this hydrocarbon?
Molecular Formula example 2
 An unknown hydrocarbon is found to be 84.2% carbon by
mass. It has been determined that the molar mass of
the hydrocarbon is 114 g/mole. Find the molecular
formula of this hydrocarbon.
EMPIRICAL FORMULA BY ANALYSIS: Involves the
combustion of organic compounds
1) The mass of carbon is calculated from the CO2
produced
2) The mass of hydrogen is calculated from the H2O
produced
3) Oxygen or any other element is calculated by
subtraction of Carbon & Hydrogen form the total sample.
Analysis example 1
 A 27.0 gram sample of an unknown hydrocarbon was burned in
excess oxygen to form 88.0 grams of CO2 and 27 grams of H2O.
The hydrocarbon was found to have a molecular mass of 54
g/mole. Find the molecular formula of the hydrocarbon.
Analysis example 2
 A 2.78 mg sample of an organic compound contain carbon,
hydrogen, and oxygen was burned completely. Combustion of
the compound produced 6.32 mg of carbon dioxide and 2.58 mg
of water. What is the empirical formula of this compound?
Analysis example 3
 A 0.116 g sample of a compound of C, H, & N was combusted,
producing 0.164 g of CO2 and 0.168 g of H2O. Find the empirical
formula of the compound.
Formulas of Molecular Compounds
 Covalent bonds form between nonmetals, and involve
sharing of valence electrons.
 Lewis structures show the elemental symbol surrounded
by valence electrons.
 H, C, N, O, F, P, S, Cl, Br, and I are all nonmetals.
 This type of atom tends to gain electrons (perhaps to
form anions).
Formulas of Ionic Compounds
 When cations and anions combine, ionic compounds are
formed.
 Ionic compounds are ELECTRICALLY NEUTRAL. This
means there is no overall charge. All charges cancel to
equal zero.
Ionic Compound Naming
 The first element/word in an ionic compound is a cation
(always a metal, except ammonium).
 The second element/word in an ionic compound is an
anion (nonmetal or polyatomic ion).
 Roman numerals tell the reader the CHARGE of the ion.
Chemical Formulas & Nomenclature
 What is the rule for the total overall charge for a
compound?
Steps to writing an ionic chemical
formula:
(I) Write the symbols of the two elements.
(II) Write the valence of each as superscripts.
(III) Drop the positive and negative signs.
(IV) Crisscross the superscripts so they become subscripts.
(V) Reduce when possible.
 Molecular Compounds: atoms of two or more different
elements which are bound together by strong covalent
bonds
 Exist as separate molecules which behave as one unit
 Molecular formulas give the actual number of each
atom within the compound
 Ionic Compounds: Ions of two or more different
element held together by a strong electrostatic (+ to -)
attraction
 Exist as a crystal lattice structure of many ions (does
not exist as separate molecules)
 Empirical formula gives the simplest ratio of ion
present with the crystal lattice structure
 Monatomic ions: ions composed of one atom
 Name usually ends in –ide
 Polyatomic ions: ions composed of more than one
atom covalently bonded together with a net charge
 Behave as one unit
 Names usually end in –ate or ite
Polyatomic Ions (with a 1- charge)
Hydroxide
OH1-
Hypochlorite
ClO1-
Nitrate
NO31-
Cyanide
CN1-
Nitrite
NO2
Permanganate
KMnO41-
1-
Acetate
C3H2O31-
Perchlorate
ClO41-
Chlorate
ClO3
Chlorite
ClO21-
1-
Bicarbonate
(hydrogen
carbonate)
Bisulfate
HCO31-
(hydrogen sulfate)
HSO41-
Thiocyanate
SCN1-
Polyatomic Ions (other charges)
Ammonium
Phosphate
NH41+
PO43-
Sulfate
SO42-
Sulfite
SO32-
Carbonate
CO32-
Oxalate
C2O42-
Chromate
CrO42-
Dichromate
Cr2O72-
Naming Molecular Compounds
(Covalent)
Molecular Compound:
 Use Greek prefixes to denote the number of each atom
present
 Mono is never used when naming the first element
Prefix
MonoDiTriTetraPenta-
#
1
2
3
4
5
Prefix
HexaHeptaOctaNonaDeca-
#
6
7
8
9
10
Naming Acids
Acids: Molecular compounds which start with H+
Acids with Anions that do not contain oxygen:
 Name the acid beginning with the prefix hydro
 Change anion name to end in –ic acid
 Add the word acid
Acids with Anions that do contain oxygen(oxyacids):
 Do not use the prefix hydro-
 Change polyatomic ion’s ending -ate  -ic acid
-ite  -ous acid
Nomenclature Practice
Common
Name
Systematic Name
1) Table Salt
Sodium Chloride
2) Baking Soda
Sodium hydrogen carbonate
3) Muratic acid
Hydrochloric acid
4) Laughing gas
Dinitrogen monoxide
5) Rust
Iron (III) oxide
Chemical Formula
Nomenclature Practice
Common
Name
Systematic Name
6)
Vinegar
acetic acid
7)
phosphine
Phosphorous trihydride
8)
Lime
Calcium hydroxide
9)
Lye
10) Bleach
Chemical Formula
NaOH
NaClO
Nomenclature Practice
Common
Name
Systematic Name
Chemical
Formula
11) Limestone
CaCO3
12) Galena
PbS
13) Quartz
SiO2
14) Seltzer water
H2CO3
11) Limestone
CaCO3