Chapter 4 Notes- Aqueous Reactions and Solution Stoichiometry
4.1 General Properties of Aqueous Solutions [p.120]
1. General information
a. Aqueous Solution
b. Solution
c. Solute
d. Solvent
2. Electrolytic Properties
a. Ions in an aqueous solution allow for the conduction of electricity
b. Electrolyte- substance whose aqueous solution contains ions
c. Nonelectrolyte – a substance that does not form ions in an aqueous solution
3. Ionic Compounds in Water
a. Ionic compounds dissociate into ions when dissolving
b. Water is a polar molecule making it a very effective solvent.
i. Cations are attracted to the negative side of the molecule
ii. Anions are attracted to the positive side of the molecule
iii. Water molecules surround ions keeping the ions dissociated
c. Solvation- process that stabilizes the ions in solution
d. Solvent and ions are free to move which allows for the uniform dispersion of ions
e. Important to know the ions that form a compound so that the ions in solution can be
determined
4. Molecular Compounds in Water
a. Most molecular (covalent) compounds dissolve in water as intact molecules
i. Not electrolytes
ii. Example is Alcohol
b. Acids that are molecular compounds due ionize
5. Strong and Weak Electrolytes
a. Strong electrolytes
i. Conduct electricity
ii. Completely or almost completely dissociate
iii. Examples include NaCl and HCl
iv. Soluble ionic compounds
1. metal-nonmetal
2. contain ammonium
b. Weak Electrolytes
i. Poor conductors
ii. Dissolve as molecules with only a few ions possible
iii. Examples include acetic acid
c. STRENGTH OF ELECTROLTYE HAS NOTHING TO DO THE ABILITY OF AN
ELECTROLYTE TO DISSOLVE
d. Chemical Equilibrium is established when a weak electrolyte ionizes in solution
CH3COOH (aq) 
 CH3COO  (aq)  H  (aq)
e. Strong electrolytes completely ionize therefore no equilibrium is established.
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Chapter 4 Notes- Aqueous Reactions and Solution Stoichiometry
4.2 Precipitation Reactions [p.124]
1. General Definitions
a. Precipitation reaction
b. Precipitate
2. Solubility Guidelines for Ionic Compounds
a. Solubility – amount of substance that can be dissolved in a given quantity of solvent
b. Insoluble substances have a solubility of less than .001 mol/L
i.
Charges between particles in ionic compound are too strong for water to
separate
c. Must know solubility rule.
i.
Found on p.125
ii. Additional rules found on Regents Chemistry Reference Tables
iii. Examples: all ionic compounds containing nitrate, acetate, alkali metals, and
ammonium are soluble
iv.
Must use solubility rules to determine the precipitate in a double replacement
reaction
3. Exchange (Metathesis) Reactions
a. Also known as double replacement reactions
b. Reaction where cations and anions switch patterns
c. General formula
AB + CD  AD + CB
d. Examples: Acid- base reactions and precipitation reactions
e. Must be able write equation and balance
4. Ionic Equations
a. Molecular equations show the chemical formulas of the reactants and products.
b. Ionic equations indicates that soluble electrolytes as ions
c. Spectator ions appear on both sides of an equation
d. Net ionic equation only includes ions and molecules involved in a reaction
e. If all ions are spectators, then no reaction occurred.
f. Steps for writing a net ionic equation
i.
Write balanced molecular equation.
ii. Rewrite as complete ionic equation
4.3 Acid-Base Reactions [p. 128]
1. Acids
a. Substances that ionize in aqueous solution to form hydrogen ions [H+]- proton donors
b. Proton in water- H+(aq)
c. Ionization of acid is different for different acids
i. Monoprotic acids- release one H+
1. HCl
2. HNO3
ii. Diprotic acids release 2 H+
1. Example Sulfuric Acid (H2SO4)
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Chapter 4 Notes- Aqueous Reactions and Solution Stoichiometry
Step1 : H 2 SO2 (aq) 
 H  (aq)  HSO4 (aq)
Step2 : HSO4 (aq) 
 H  (aq)  SO42 (aq)
***Note step 1 completely ionizes, but not step 2. Therefore, H+, SO42-,
HSO4- are present in solution
2. Acetic acid has 4 hydrogen in only hydrogen ion attached to –COOH
will ionize
2. Bases
a.
b.
c.
d.
Substances that accept hydrogen ions
Produce OH- when they disociate
Common bases- NaOH, KOH, Ca(OH)2
Bases that do not contain hydroxide ions include ammonia, NH3, (weak electrolyte)


NH 3 (aq)  H 2 O(l ) 
 NH 4 (aq)  OH  (aq)
3. Strong and Weak Acids and Bases
a. Strong electrolytes [strong acids and bases] completely ionize in solution
i. Strong acids HCl, HBr, HI, chloric acid, perchloric acid, nitric acid, sulfuric acid
ii. Strong Bases: Group 1A metal hydroxides and heavy group 2A metal
hydroxides
iii. More reactive when reactivity is dependent on concentration only
b. Weak Electrolytes [weak acids and bases] partly ionize
i. Excluding the metal hydroxides in the strong base category, metal hydroxides
are insoluble
4. Identifying Strong and Weak Electrolytes
a. To determine the strong v. weak electrolytes follow the following steps
Compound
Ionic
All strong
electrolytes
Molecular
Strong electrolytesStrong Acids
5. Neutralization Reactions and Salts
a. Acid properties
i. Sour taste
ii. Litmus turns red
b. Basic properties
i. Bitter taste
ii. Litmus turns blue
c. Definition- Acid –base reaction
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Weak ElectrolytesWeak acids and weak
bases
NonelectrolytesAll other compounds
Chapter 4 Notes- Aqueous Reactions and Solution Stoichiometry
i. Acid + Base  water + salt
ii. Salt = ionic compound formed by cation of base and anion of base
iii. Net ionic equation are the reaction between HCl and NaOH is
H+(aq) + OH-(aq) H2O(l)
d. Exchange of ions makes acid-base reactions a metathesis reaction
6. Acid-Base Reactions with Gas Formation
a. Bases not containing OH- react with a hydrogen ion to gaseous molecular compounds
with low solubility’s
i. Bases containing the sulfide ion and the carbonate ion combine with acid to
gases
ii. Bases containing carbonate ions and bicarbonate ions react with acid to form
CO2
iii. Carbonic acid is unstable so it will decompose into water and carbon dioxide
4.4 Oxidation- Reduction Reactions [p.135]
1. Oxidation and Reduction
a. Electrons are transferred between reactants
b. Also called redox reaction
c. Corrosion is an example of redox
d. Oxidation= loss of electrons
e. Reduction = gain of electrons
f. Use LEO says GER to remember definitions
2. Assigning Oxidation Numbers
a. Loner elements have an oxidation number = 0 includes diatomic molecules
b. Monoatomic ions – oxidation number = charge of the ion
c. Nonmetals typically have a negative charge
i. Oxygen ion usually -2 except when part of a peroxide
ii. Hydrogen is typically a +1 except when it forms a metal hydride
iii. Fluorine is -1 always
d. Neutral compounds- sum of oxidation numbers = 0
e. Ions – sum of oxidation states = charge of ion
3. Oxidation of Metals by Acids and Salts
a. Types of reactions that are redox
i. Combustion
ii. Single replacement ( frequently involve a metal reacting with an acid or a metal
salt)
4. The Activity Series
a. Ease of oxidation varies based on the properties of a metal
b. Alkali metals and alkaline earth metals high on activity series ( active metals)
c. Transition metals from 8B and 1B are very low on the activity series called noble
metals
4.5 Concentration of Solutions [p.142]
1. Concentration = amount of solute dissolved in a given quantity of solvent
2. Molarity
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Chapter 4 Notes- Aqueous Reactions and Solution Stoichiometry
a. Symbol M
b. Moles of solute per liter of solution
3. Expressing the Concentration of an Electrolyte
a. Moles of ions per liter of solution
b. Can be completed for anion or cation
4. Interconverting Molarity, Moles, and Volume
a. Use dimensional analysis
b. Requires knowledge of 2 of the 3 variables needed to find molarity
5. Dilution
a. Used to lower the concentration of a stock solution
b. M concVconc  M dilVdil
“conc” means concentrate
“dil” means dilution
4.6 Solution Stoichiometry and Chemical Analysis [p.149]
1. Titrations
a. Used to determine the concentration of a solute in solution
b. Standard solutions = solution with a known concentration
c. Neutralization reaction = acid-base reaction
d. Equivalence point occurs when the moles of H+ = the moles OHe. Indicators are dyes that help identify when the equivalence point has been reached.
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