Download Chapter 4 - Bakersfield College

Chapter 4
Molecular View of Reactions in
Aqueous Solutions
Part I
Chemistry: The Molecular Nature of
Matter, 7E
Jespersen/Hyslop
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Chapter in Context
 Describe solutions qualitatively and quantitatively
 Distinguish electrolytes from non-electrolytes
 Write balanced molecular, ionic, and net ionic
equations
 Identify acids and bases and learn names and
formulas
 Use metathesis reactions to plan chemical
syntheses
 Define and use molarity in calculations
 Understand titrations and chemical analysis
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2
Importance of Water
 One of the most common compounds on earth
 Dissolves many different substances
 Responsible, in part, for evolution of life
 60% of the human body is water
Distinct Properties
 Dissolves ionic compounds
 Acid-base reactions occur in water
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Reactions in Solution
 For a reaction to occur
 Reactants needs to come into physical contact
 Happens best in gas or liquid phase
 Movement occurs
Solution
 Homogeneous mixture
 Two or more components mix freely
 Molecules or ions completely intermingled
 Contains at least two substances
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Definitions
Solvent
 Medium that dissolves solutes
 Component present in largest amount
 Can be gas, liquid, or solid
 Aqueous solution—water is solvent
Solute
 Substance dissolved in solvent




Solution is named by solute
Can be gas—CO2 in soda
Liquid—ethylene glycol in antifreeze
Solid—sugar in syrup
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Iodine Molecules in Ethanol
Crystal of solute
placed in solvent
Solute molecules dispersed
throughout solvent
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Solutions
 May be characterized using
Concentration
 Solute-to-solvent ratio
g solute
g solvent
or
g solute
g solution
 Percent concentration
g solute
percentage concentration =
100 g solution
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Relative Concentration
Dilute solution
 Small solute to solvent ratio
e.g., Eye drops
Concentrated
solution
 Large solute to solvent
ratio
e.g., Pickle brine
 Dilute solution contains less solute per unit
volume than more concentrated solution
 ‘Dilute’ and ‘concentrated’ are relative terms
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Concentration
Solubility
g solute needed to make saturated solution
Solubility =
100 g solvent
 Temperature dependent
Saturated solution
 Solution in which no more solute can be dissolved
at a given temperature
Unsaturated solution
 Solution containing less solute than maximum
amount
 Able to dissolve more solute
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Solubilities of Some Common
Substances
Substance
Solubility
Formula (g/100 g water)
Sodium chloride
NaCl
Sodium
hydroxide
NaOH
Calcium
carbonate
CaCO3
35.7 at 0 C
o
39.1 at 100 C
o
42 at 0 C
o
347 at 100 C
o
0.0015 at 25 C
o
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Concentrations
Supersaturated Solutions
 Contains more solute than required for saturation
at a given temperature
 Formed by careful cooling of saturated solutions
 Unstable
 Crystallize out when add seed crystal – results in
formation of solid or precipitate (ppt.)
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Precipitates
Precipitate
 Solid product formed when reaction carried out in
solutions and one product has low solubility
 Insoluble product
 Separates out of solution
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Precipitates
Precipitation Reaction
Reaction that produces precipitate
Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq)
Solid
precipitate
Pb2+ NO3–
K+ I–
PbI2(s)
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Electrolytes in Aqueous Solution
 Ionic compounds conduct electricity
 Molecular compounds don’t conduct electricity
Why?
Bright
light
No
light
Ions
present
Molecular
CuSO4 and water
Sugar and water
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Ionic Compounds (Salts) in Water
 Water molecules arrange themselves around ions
and remove them from lattice.
Dissociation
 Salts break apart into
ions when entering solution
Separated ions
 Hydrated
 Conduct electricity
 Note: Polyatomic ions
remain intact
 e.g., KIO3  K+ + IO3–
NaCl(s)  Na+(aq) + Cl–(aq)
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Molecular Compounds In Water
 When molecules dissolve in water
 Solute particles are surrounded by water
 Molecules do not dissociate
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Electrical Conductivity
Electrolyte
 Solutes that yield electrically conducting solutions
 Separate into ions when enter into solution
Strong electrolyte
 Electrolyte that dissociates 100% in water
 Yields aqueous solution that conducts electricity
 Good electrical conduction
 Ionic compounds, e.g., NaCl, KNO3
 Strong acids and bases, e.g., HClO4, HCl
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Electrical Conductivity
Non-electrolyte
 Aqueous solution that doesn’t conduct electricity
 Molecules remain intact in solution
e.g., Sugar, alcohol
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Electrical Conductivity
Weak electrolyte
 When dissolved in water only a small
percentage of molecules ionize
 Common examples are weak acids and bases
 Solutions weakly conduct electricity
 e.g., Acetic acid (CH3COOH), ammonia (NH3)
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Strong vs. Weak Electrolyte
HCl(aq)
CH3COOH(aq)
NH3(aq)
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Dissociation Reactions
 Ionic compounds dissolve to form hydrated ions
 Hydrated = surrounded by water molecules
 In chemical equations, hydrated ions are
indicated by
 Symbol (aq) after each ions
 Ions are written separately
KBr(s)  K+(aq) + Br–(aq)
Mg(HCO3)2(s)  Mg2+(aq) + 2HCO3–(aq)
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Learning Check
Write the equations that illustrate the dissociation
of the following salts:
Na3PO4(aq)
→ 3Na+(aq) + PO43–(aq)
Al2(SO4)3(aq) → 2Al3+(aq) + 3SO42–(aq)
CaCl2(aq)
2+(aq) + 2Cl–(aq)
Ca
→
Ca(MnO4)2(aq) → Ca2+(aq) + 2MnO4–(aq)
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Equations of Ionic Reactions
 Consider the reaction of Pb(NO3)2 with KI
Pb2+
NO3–
K+
I–
PbI2(s)
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Equations of Ionic Reactions
 When two soluble ionic solutions are mixed,
sometimes an insoluble solid forms.
 Three types of equations used to describe
1. Molecular equation
 Substances listed as complete formulas
2. Ionic equation
 All soluble substances broken into ions
3. Net ionic equation
 Only lists substances that actually take part in
reaction
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Equations of Ionic Reactions
1. Molecular Equation
 Complete formulas for all reactants and products
 Formulas written with ions together
 Does not indicate presence of ions (no charges)
 Gives identities of all compounds
 Good for planning experiments
e.g.,
Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq)
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Equations of Ionic Reactions
2. Ionic Equation
 Emphasizes the reaction between ions
 All strong electrolytes dissociate into ions
 Used to visualize what is actually occurring in
solution
 Insoluble solids written together as they don’t
dissociate to any appreciable extent
e.g.,
Pb2+(aq) + 2NO3–(aq) + 2K+(aq) + 2I–(aq) 
PbI2(s) + 2K+(aq) + 2NO3–(aq)
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Equations of Ionic Reactions
Spectator Ions
 Ions that don’t take part in reaction
 They hang around and watch
 K+ and NO3– in our example
3. Net Ionic Equation
 Eliminate all spectator ions
 Emphasizes the actual reaction
 Focus on chemical change that occurs
e.g., Pb2+(aq) + 2I–(aq)  PbI2(s)
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Equations of Ionic Reactions
Criteria for ionic and net ion equations
 Material balance
 The same number of each kind of atom must be
present on both sides of the arrow.
 Electrical balance
 The net electrical charge on the reactants must
equal the net electrical charge on the products
 Charge does not necessarily have to be zero
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Net Ionic Equations
 Many ways to make PbI2
1. Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq)
2. Pb(C2H3O2)2(aq) + 2NH4I(aq) 
PbI2(s) + 2NH4C2H3O2(aq)
 Different starting reagents
 Same net ionic equation
 Pb2+(aq) + 2I–(aq)  PbI2(s)
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Converting Molecular Equations to
Ionic Equations
Strong electrolytes exist as dissociated ions in
solution
Strategy
1. Identify strong electrolytes
2. Use subscript coefficients to determine total
number of each type of ion
3. Separate ions in all strong electrolytes
4. Show states as recorded in molecular equations
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Learning Check: Convert Molecular
to Total Ionic Equations:
Write the correct ionic equation for each:
Pb(NO3)2(aq) + 2NH4IO3(aq) →
Pb(IO3)2(s) + 2NH4NO3(aq)
Pb2+(aq) + 2NO3–(aq) + 2NH4+(aq) + 2IO3–(aq) →
Pb(IO3)2(s) + 2NH4+(aq) + 2NO3–(aq)
2NaCl (aq) + Hg2(NO3)2 (aq) → 2NaNO3 (aq) + Hg2Cl2 (s)
2Na+(aq) + 2Cl–(aq) + Hg22+(aq) + 2NO3–(aq) →
2Na+(aq) + 2NO3–(aq) + Hg2Cl2(s)
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Converting Ionic Equations to Net
Ionic Equations
Strategy
1. Identify spectator ions
2. Cancel from both sides
3. Rewrite equation using only substances that
actually react.
4. Show states as recorded in molecular and ionic
equations
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Learning Check: Convert Ionic
Equation to Net Ionic Equation
Write the correct net ionic equation for each.
Pb2+(aq) + 2NO3–(aq) + 2K+(aq) + 2IO3–(aq) →
Pb(IO3)2(s)+ 2K+(aq) + 2NO3–(aq)
Pb2+(aq) + 2IO3–(aq) → Pb(IO3)2(s)
2Na+(aq) + 2Cl–(aq) + Hg22+(aq) + 2NO3–(aq) →
2Na+(aq)+ 2NO3–(aq) + Hg2Cl2(s)
2Cl–(aq) + Hg22+(aq) → Hg2Cl2(s)
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Criteria for Balancing Ionic and
Net Ionic Equations
1. Material Balance
 There must be the same number of atoms of
each kind on both sides of the arrow
2. Electrical Balance
 The net electrical charge on the left must
equal the net electrical charge on the right
 Charge does not have to be zero
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Learning Check: Balancing Equations
for Mass & Charge
Balance molecular equation for mass
2Na3PO4(aq) + 3Pb(NO3)2(aq) 
6NaNO3(aq) + Pb3(PO4)2(s)
 Can keep polyatomic ions together when counting
Balance total ionic equation for charge
6Na+(aq) + 2PO43–(aq) + 3Pb2+(aq) + 6NO3–(aq) 
6Na+(aq) + 6NO3–(aq) + Pb3(PO4)2(s)
 Charge must add up to zero on both sides.
Net ionic equation balanced for mass and charge
3Pb2+(aq) + 2PO43–(aq)  Pb3(PO4)2(s)
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Acids and Bases
 Common laboratory reagents
 Also found in food and household products
 vinegar, citrus juice, and cola contain acids
 drain cleaners and ammonia contain bases
 Acids
 Tart, sour taste
 Bases
 Bitter taste and slippery feel
 Caution: Never taste, feel, or smell laboratory chemicals
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Arrhenius Acid
 Substance that reacts with water to produce the
hydronium ion, H3O+
HCl(g) + H2O

Acid + H2O

Anion + H3O+
HA + H2O

A– + H3O+
Cl–(aq) + H3O+(aq)
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Arrhenius Acid
Another example
HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2−(aq)
Ionization reaction definition
 Ions form where none have been before
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Arrhenius Base
 Substance that produces OH–
 Ionic substances containing OH– or O2-
 Molecular substances
Ionic compound containing OH–
a. Metal hydroxides
 Dissociate into metal and hydroxide ions
NaOH(s)  Na+(aq) + OH–(aq)
Mg(OH)2(s)  Mg2+(aq) + 2OH–(aq)
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Strong Acids
HClO4(aq)
HClO3(aq)
HCl(aq)
HBr(aq)
HI(aq)
HNO3(aq)
H2SO4(aq)
perchloric acid
chloric acid
hydrochloric acid
hydrobromic acid
hydroiodic acid
nitric acid
sulfuric acid
 Dissociate completely when dissolved in water
e.g., HBr(g) + H2O  H3O+(aq) + Br–(aq)
 Good electrical conduction (i.e., strong electrolytes)
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Strong Bases
 Bases that dissociate completely in water
 Soluble metal hydroxides
 KOH(aq)  K+(aq) + OH–(aq)
 Good electrical conductors (i.e., strong electrolytes)
 Behave as aqueous ionic compounds
 Common strong bases are:
 Group 1A metal hydroxides
 LiOH, NaOH, KOH, RbOH, CsOH
 Group 2A metal hydroxides
 Ca(OH)2, Sr(OH)2, Ba(OH)2
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Weak Acids
 Any acid other than seven strong acids
 Are also weak electrolytes, i.e, ionize < 100%
Organic acids
HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2–(aq)
Acetic Acid
Molecule,
HC2H3O2
Only this H comes off as H+
Acetate ion, C2H3O2–
e.g.,
HCO2H(aq) + H2O  H3O+(aq) + HCO2–(aq)
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Why is Acetic Acid Weak?
CH3COOH(aq) + H2O  CH3COO-(aq) + H3O+(aq)
CH3COO–(aq) + H3O+ (aq)  CH3COOH (aq) + H2O
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Dynamic Equilibrium
 Two opposing reactions occurring at same rate
 Also called chemical equilibrium
Equilibrium
 Concentrations of substances present in solution do
not change with time
Dynamic
 Both opposing reactions occur continuously
 Represented by double arrow
HC2H3O2(aq) + H2O
H3O+(aq) + C2H3O2–(aq)
Forward reaction – forms ions
Reverse reaction – forms molecules
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Arrhenius Bases
2. Molecular Bases
 Undergo ionization (hydrolysis) reaction to form
hydroxide ions
Base + H2O  BaseH+(aq) + OH–(aq)
B + H2O  BH+(aq) + OH–(aq)
NH3(aq) + H2O  NH4+(aq) + OH–(aq)
NH3
H2O

NH4+
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OH–
45
Weak Bases
 Molecular bases
 Do not dissociate
 Accept H+ from water inefficiently
 Are weak electrolytes
e.g.,
NH3(aq) + H2O
NH4+(aq) + OH–(aq)
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Equilibrium for Weak Base
Forward reaction
Reverse reaction
Net is dynamic equilibrium
NH3(aq) + H2O
NH4+(aq) + OH–(aq)
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General Ionization Equations
 Strong acid in water
HX(aq) + H2O  H3O+(aq) + X–(aq)
 Strong base, M(OH)n
M(OH)n  Mn+(aq) + nOH–(aq)
 Weak acid in water
HA(aq) + H2O
H3O+(aq) + A–(aq)
 Weak base in water
B(aq) + H2O
HB+(aq) + OH–(aq)
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Learning Check
 Write the ionization equation for each of the
following with water:
1. Weak base methylamine, CH3NH2
CH3NH2(aq) + H2O
CH3NH3+(aq) + OH–(aq)
2. Weak acid nitrous acid, HNO2
HNO2(aq) + H2O
H3O+(aq) + NO2–(aq)
3. Strong acid chloric acid, HClO3
HClO3(aq) + H2O  H3O+(aq) + ClO3–(aq)
4. Strong base strontium hydroxide, Sr(OH)2
Sr(OH)2(aq)  Sr2+(aq) + 2OH–(aq)
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Brief summary
 Strong acids and bases are strong electrolytes
 Weak acids and bases are weak electrolytes
 Strong electrolyte
 Weak electrolyte
 Completely ionizes
 Small % ionizes
 Forward reaction
dominates
 Reverse rxn dominates
 Mostly products
 Mostly reactants
 Strong acids & bases  Weak acids and bases
 Little reverse reaction  Lots of reverse reaction
 Write eqn. as 
 Write eqn. as
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Polyprotic Acids
Monoprotic Acids
 Furnish only one H+
HNO3(aq) + H2O  H3O+(aq) + NO3–(aq)
HC2H3O2(aq) + H2O  H3O+(aq) + C2H3O2–(aq)
Diprotic acids — furnish two H+
H2SO3(aq) + H2O  H3O+(aq) + HSO3–(aq)
HSO3–(aq) + H2O  H3O+(aq) + SO32–(aq)
Polyprotic acids
 Furnish more than one H+
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Polyprotic Acids
Polyprotic acids
 Triprotic acids — furnish three H+
– H+
– H+
– H+
H3PO4  H2PO4–  HPO42–  PO43–
 Stepwise equations
H3PO4(aq) + H2O  H3O+(aq) + H2PO4–(aq)
H2PO4–(aq) + H2O  H3O+(aq) + HPO42–(aq)
HPO42–(aq) + H2O  H3O+(aq) + PO43–(aq)
Net:
H3PO4(aq) + 3H2O  3H3O+(aq) + PO43–(aq)
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Learning Check
 Write the stepwise ionization reactions for citric
acid, H3C6H5O7, in water.
H3C6H5O7(aq) + H2O  H3O+(aq) + H2C6H5O7–(aq)
H2C6H5O7–(aq) + H2O  H3O+(aq) + HC6H5O72-(aq)
HC6H5O72-(aq) + H2O  H3O+(aq) + C6H5O73-(aq)
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Acidic Anhydrides
Nonmetal Oxides
 Act as Acids
 React with water to form molecular acids that
contain hydrogen
SO3(g) + H2O  H2SO4(aq)
sulfuric acid
N2O5(g) + H2O  2HNO3(aq)
nitric acid
CO2(g) + H2O  H2CO3(aq)
carbonic acid
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Ionic Oxides
b. Basic Anhydrides
 Soluble metal oxides
 Undergo ionization (hydrolysis) reaction to
form hydroxide ions
 Oxide reacts with water to form metal hydroxide
CaO(s) + H2O  Ca(OH)2(aq)
O2– + H2O

2OH–
 Then metal hydroxide dissociates in water
Ca(OH)2(aq)  Ca2+(aq) + 2OH–(aq)
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Acid—Base Nomenclature
 System for naming acids and bases
Acids
 Binary acid system e.g., HCl(aq), H2S(aq)
 Oxoacid system e.g., H2SO4, HClO2
 Acid salt system e.g., NaHSO4, NaHCO3
Bases
 Metal hydroxide/oxide system e.g., NaOH, CaO
 Molecular base system e.g., NH3, (CH3)3N
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Naming Acids
A. Binary Acids — hydrogen + nonmetal




Take molecular name
Drop –gen from H name
Merge hydro– with nonmetal name
Replace –ide with –ic acid
Name of Molecular
Compound
Name of Binary Acid in
water
HCl(g) hydrogen chloride HCl(aq) hydrochloric acid
H2S(g) hydrogen sulfide
H2S(aq) hydrosulfuric
acid
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Naming Acids
B. Oxo Acids
 Acids with hydrogen, oxygen and another nonmetal
element
 A table of polyatomic ions can be found in the book
 To name:
 Based on parent oxoanion name
 Take parent ion name
 Anion ends in –ate change to –ic (more O's)
 Anion ends in –ite change to –ous (less O's)
 End name with acid to indicate H+
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Oxoacids (Aqueous)
Named according to the anion suffix
 Anion ends in -ite, acid name is -ous acid
 Anion ends in -ate, acid name is -ic acid
Name of Parent
Oxoanion
nitrate
NO3
Name of Oxoacid
HNO3
nitric acid
SO42
sulfate
H2SO4
sulfuric acid
ClO2
chlorite
HClO2
chlorous acid
PO32
phosphite
H2PO3
phosphorous acid
Jespersen/Hyslop, Chemistry7E, Copyright © 2015 John Wiley & Sons, Inc. All Rights Reserved
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Learning Check: Name Each
Aqueous Acid
 HNO2
 nitrous acid
 HCN
 hydrocyanic acid
 HClO4
 perchloric acid
 HF
 hydrofluoric acid
 H2CO3
 carbonic acid
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60
C. Naming Bases
Oxides & Hydroxides
 Ionic compounds
 Named like ionic compounds
 Ca(OH)2
 Li2O
calcium hydroxide
lithium oxide
Molecular Bases
 Named like molecules




NH3
CH3NH2
(CH3)2NH
(CH3)3N
ammonia
methylamine
dimethylamine
trimethylamine
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61