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Chapter 4
Reactions in Aqueous Solution
▶ DEEP SEA VENTS are amazing
places. Superheated water (up to
400 oC) is released from cracks in
the bottom of the ocean. Rocks
dissolve and reform. The locally
high mineral content and sulfurcontaining substances in the water
provide an environment that favors
unusual organisms that are found
nowhere else in the world.
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Reactions in Aqueous Solution
What’s Ahead
4.1 GENERAL PROPERTIES OF AQUEOUS SOLUTIONS
We begin by examining whether substances dissolved in water exist as ions, molecules,
or a mixture of the two.
4.2 PRECIPITATION REACTIONS
We identify reactions in which soluble reactants yield an insoluble product.
4.3 ACIDS, BASES, AND NEUTRALIZATION REACTIONS
We explore reactions in which protons, H+ ions, are transferred from one reactant to
another.
4.4 OXIDATION-REDUCTION REACTIONS
We examine reactions in which electrons are transferred from one reactant to another.
4.5 CONCENTRATIONS OF SOLUTIONS
We learn how the amount of a compound dissolved in a given volume of a solution can
be expressed as a concentration. Concentration can be defined in a number of ways, the
most commonly used being moles of compound per liter of solution (molarity).
4.6 SOLUTION STOICHIOMETRY AND CHEMICAL ANALYSIS
We see how the concepts of stoichiometry and concentration can be used to calculate
amounts or concentrations of substances in solution through a process called titration.
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4.1 GENERAL PROPERTIES OF AQUEOUS SOLUTIONS
Solutions
• Solutions are defined as
homogeneous mixtures of
two or more pure substances.
• The solvent is present in
greatest abundance.
• All other substances are
solutes.
• When water is the solvent,
the solution is called an
aqueous solution.
Figure 4.3(b) An aqueous solution: Methanol
(CH3OH, solute) dissolved in water (solvent).
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4.1 GENERAL PROPERTIES OF AQUEOUS SOLUTIONS
Aqueous Solutions of Ionic Compounds
• When an ionic substance
dissolves in water, the
solvent pulls the
individual ions from the
crystal and solvates them.
• This process is called
dissociation.
Figure 4.3(a) An aqueous solution: Sodium
chloride (NaCl, solute) dissolved in water (solvent).
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4.1 GENERAL PROPERTIES OF AQUEOUS SOLUTIONS
Aqueous Solutions
• Substances can dissolve in water by different ways:
– Ionic compounds dissolve by dissociation, where water surrounds
the separated ions.
– Molecular compounds interact with water, but most do NOT dissociate.
• A few molecular substances have aqueous solutions that contain ions. Ex)
HCl(aq) ionizes; it dissociates into H+(aq) and Cl–(aq) ions.
– Some molecular substances react with water when they dissolve.
Figure 4.3 Dissolution in water. (a)
When an ionic compound, such as
sodium chloride, NaCl, dissolves in water,
H2O molecules separate, surround, and
uniformly disperse the ions into the liquid.
(b) Molecular substances that dissolve in
water, such as methanol, CH3OH, usually
do so without forming ions. We can think
of this as a simple mixing of two molecular
species. In both (a) and (b) the water
molecules have been moved apart so that
the solute particles can be seen clearly.
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4.1 GENERAL PROPERTIES OF AQUEOUS SOLUTIONS
Electrolytic Properties
• Pure water is a very poor conductor of electricity.
• The conductivity of bathwater originates from the substances
dissolved in the water.
• When ions are present in solution, the ions carry electrical
charge from one electrode to the other, completing the circuit.
Figure 4.2 Completion of an electrical circuit with an electrolyte turns on the light.
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4.1 GENERAL PROPERTIES OF AQUEOUS SOLUTIONS
Electrolytes and Nonelectrolytes
• An electrolyte is a substance that dissociates into
ions when dissolved in water.
– A strong electrolyte dissociates completely when dissolved
in water.
– A weak electrolyte only dissociates partially when dissolved
in water.
• A nonelectrolyte may dissolve in water, but it does
not dissociate into ions when it does so.
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4.1 GENERAL PROPERTIES OF AQUEOUS SOLUTIONS
Electrolytes and Nonelectrolytes
• Soluble ionic compounds tend to be electrolytes.
• Molecular compounds tend to be nonelectrolytes,
except for acids and bases.
• Dissolution vs. ionization (CH3COOH vs. Ca(OH)2).
– CH3COOH: Highly soluble in water but a weak electrolyte.
– Ca(OH)2: Not very soluble in water but a strong electrolyte.
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4.1 GENERAL PROPERTIES OF AQUEOUS SOLUTIONS
Strong and Weak Electrolytes
• Strong electrolytes: Strong acids (HCl), strong bases (NaOH),
and water-soluble ionic compounds (NaCl, FeSO4, Al(NO3)3).
– Strong electrolytes exist in solution (nearly) completely as ions.
– Water-soluble ionic compounds: Combinations of metal and nonmetal elements (Ammonium containing compounds [NH4Br and
(NH4)2CO3] are exceptions to this rule of thumb.)
– A single arrow is used to represent the ionization of strong electrolytes.
• Weak electrolytes: Weak acids (CH3COOH) and weak bases.
– Weak electrolytes exist in solution mostly in the form of molecules with
only a small fraction (about 1%) in the form of ions.
– The aqueous solutions are in a state of chemical equilibrium.
– Half-arrows pointing in opposite directions are used to represent the
ionization of weak electrolytes.
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4.2 PRECIPITATION REACTIONS
Precipitation Reactions
• When two solutions containing soluble salts are
mixed, sometimes an insoluble salt will be produced.
A salt “falls” out of solution, like snow out of the sky.
This solid is called a precipitate.
Figure 4.4 A precipitation reaction.
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4.2 PRECIPITATION REACTIONS
Solubility Guidelines for Ionic Compounds
• Solubility: The amount of a substance dissolved in a
given quantity of solvent at a given temperature.
• Some ionic substances are insoluble in water because
the attraction between the opposite charges is too
great for the water molecules to separate the ions.
• No general rules to predict solubility of a substance.
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4.2 PRECIPITATION REACTIONS
Solubility of Ionic Compounds
• A list of solubility rules is used to decide what combination
of ions will dissolve.
– All common ionic compounds containing NO3– or CH3COO– are
soluble in water.
– All common ionic compounds containing the alkali metal ions or
NH4+ are soluble in water.
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4.2 PRECIPITATION REACTIONS
Exchange (Metathesis) Reactions
• Metathesis comes from a Greek word that means
“to transpose.”
• It appears as though the ions in the reactant
compounds exchange, or transpose, ions, as seen
in the equation below.
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4.2 PRECIPITATION REACTIONS
Completing and Balancing Metathesis Equations
• Steps to complete and balance the equation for a
metathesis reaction.
1. Use the chemical formulas of the reactants to determine
which ions are present.
2. Write formulas for the products: cation from one reactant,
anion from the other. Use charges to write proper subscripts.
3. Check your solubility rules. For a precipitation reaction to
occur, at least one product must be insoluble in water.
4. Balance the equation.
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4.2 PRECIPITATION REACTIONS
Molecular Equation
• Ways to Write Metathesis Reactions
– Molecular equation
– Complete ionic equation
– Net ionic equation
• The molecular equation lists the complete chemical
formulas of reactants and products without indicating the
ionic nature of the compounds.
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4.2 PRECIPITATION REACTIONS
Complete Ionic Equation
• In the complete ionic equation all strong electrolytes
(strong acids, strong bases, and soluble ionic salts) are
dissociated into their ions.
– This more accurately reflects the species that are found in the
reaction mixture.
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4.2 PRECIPITATION REACTIONS
Net Ionic Equation
• To form the net ionic equation, cross out anything that does
not change from the left side of the equation to the right.
• The ions crossed out are called spectator ions, K+ and
NO3–, in this example. They play no direct role in the reaction.
• The remaining ions are the reactants that form the product–
an insoluble salt in a precipitation reaction.
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4.2 PRECIPITATION REACTIONS
Procedure for Writing Net Ionic Equations
1. Write a balanced molecular equation for the reaction.
AgNO3(aq) + KCl(aq) → AgCl(s) + KNO3(aq)
2. Dissociate all strong electrolytes.
Ag+(aq) + NO3−(aq) + K+(aq) + Cl−(aq) →
AgCl(s) + K+(aq) + NO3−(aq)
3. Identify and cancel spectator ions.
Ag+(aq) + NO3−(aq) + K+(aq) + Cl−(aq) →
AgCl(s) + K+(aq) + NO3−(aq)
4. Write the net ionic equation with the species that remain.
Ag+(aq) + Cl−(aq) → AgCl(s)
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4.3 ACIDS, BASES, AND NEUTRALIZATION REACTIONS
Acids and Bases
• Many acids and bases are industrial and household
substances.
• Hydrochloric acid: An important industrial chemical and
the main constituent of gastric juice in our stomach.
• Acids and bases are also common electrolytes.
Figure 4.5 Vinegar and lemon juice are common household acids. Ammonia
and baking soda (sodium bicarbonate) are common household bases.
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4.3 ACIDS, BASES, AND NEUTRALIZATION REACTIONS
Acids
• Arrhenius defined acids as substances that increase
the concentration of H+ when dissolved in water.
• Brønsted and Lowry defined them as proton donors.
– Monoprotic acids: HCl and HNO3,
one H+ per molecule of acid.
– Diprotic acids: H2SO4, two H+ per
molecule of acid.
– CH3COOH (acetic acid): The primary
component in vinegar.
Figure 4.6 Molecular models
of four common acids.
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4.3 ACIDS, BASES, AND NEUTRALIZATION REACTIONS
Bases
• Arrhenius defined bases as substances that increase the
concentration of OH− when dissolved in water.
• Brønsted and Lowry defined them as proton (H+) acceptors.
– NaOH, KOH and Ca(OH)2 are common bases.
– NH3 is also a common base although it does not have OH−.
Figure 4.7 Hydrogen ion transfer. An H2O molecule acts as a proton
donor (acid), and NH3 acts as a proton acceptor (base). Only a fraction of
the NH3 molecules react with H2O. Consequently, NH3 is a weak electrolyte.
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4.3 ACIDS, BASES, AND NEUTRALIZATION REACTIONS
Strong and Weak Acids and Bases
• Strong acids and bases are completely ionized in solution.
– There are only seven strong acids.
– The strong bases are the soluble metal salts of hydroxide ion: Alkali
metals, calcium, strontium, and barium hydroxides.
• Weak acids only partially dissociate. Ex) CH3COOH and HF
• Weak bases only partially react to produce hydroxide anions.
– Ex) NH3
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4.3 ACIDS, BASES, AND NEUTRALIZATION REACTIONS
Identifying Strong and Weak Electrolytes
• All soluble ionic compounds are strong electrolytes.
• All strong acids are strong electrolytes, and all weak acids
are weak electrolytes.
• Weak base, NH3 is a weak electrolyte.
• Any molecular substance that we encounter in this chapter
that is not an acid or NH3 is probably a nonelectrolyte.
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4.3 ACIDS, BASES, AND NEUTRALIZATION REACTIONS
Different Properties of Acids and Bases
• Sour (acids) and bitter (bases) taste.
• Different changes of the colors of certain dyes.
Figure 4.8 Litmus paper. Litmus paper is
coated with dyes that change color in
response to exposure to either acids or bases.
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4.3 ACIDS, BASES, AND NEUTRALIZATION REACTIONS
Acid-Base Reactions
• In an acid–base reaction, the acid (H2O above) donates
a proton (H+) to the base (NH3 above).
• Reactions between an acid and a base are called
neutralization reactions.
• A neutralization reaction between an acid and a metal
hydroxide produces water and a salt.
– The term salt has come to mean any ionic compound whose
cation comes from a base (for example, from NaOH) and
whose anion comes from an acid (for example, from HCl).
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4.3 ACIDS, BASES, AND NEUTRALIZATION REACTIONS
Neutralization Reactions
• Neutralization reaction between hydrochloric acid and a
solution of sodium hydroxide.
– Molecular equation:
– Complete ionic equation:
– Net ionic equation:
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4.3 ACIDS, BASES, AND NEUTRALIZATION REACTIONS
Neutralization Reactions and Salts
• Neutralization reaction between hydrochloric acid and the
water-insoluble base Mg(OH)2.
Figure 4.9 Neutralization reaction between
Mg(OH)2(s) and hydrochloric acid. Milk of
magnesia is a suspension of water-insoluble
magnesium hydroxide, Mg(OH)2(s), in water.
When sufficient hydrochloric acid, HCl(aq), is
added a reaction ensues that leads to an
aqueous solution containing Mg2+(aq) and
Cl–(aq) ions.
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4.3 ACIDS, BASES, AND NEUTRALIZATION REACTIONS
Neutralization Reactions with Gas Formation
• These metathesis reactions do not give the expected product
(H2CO3). It decomposes to give a gaseous product (CO2):
CaCO3(s) + 2 HCl(aq) → CaCl2(aq) + CO2(g) + H2O(l)
NaHCO3(aq) + HBr(aq) → NaBr(aq) + CO2(g) + H2O(l)
– When a CO32– (carbonate) or HCO3– (bicarbonate) reacts with an
acid, the products are a salt, CO2, and H2O.
• Reaction between sodium bicarbonate and hydrochloric acid:
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4.3 ACIDS, BASES, AND NEUTRALIZATION REACTIONS
Neutralization Reactions with Gas Formation
• Similarly, when a sulfite reacts with an acid, the products are
a salt, sulfur dioxide, and water:
SrSO3(s) + 2HI(aq) → SrI2(aq) + SO2(g) + H2O(l)
• The reaction below gives the predicted product. Just as in the
previous examples, a gas (H2S with a characteristic “rotten
egg” smell) is formed as a product:
Na2S(aq) + H2SO4(aq) → Na2SO4(aq) + H2S(g)
– Reaction of hydrochloric acid with sodium sulfide:
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4.3 ACIDS, BASES, AND NEUTRALIZATION REACTIONS
Antacids
•
•
•
We can address the problem of excess stomach acid in two ways: (1) removing the excess
acid or (2) decreasing the production of acid. Substances that remove excess acid are
called antacids, whereas those that decrease acid production are called acid inhibitors.
FIGURE 4.10 shows several common over-the-counter antacids, which usually contain
hydroxide, carbonate, or bicarbonate ions (TABLE 4.4).
Antiulcer drugs, such as Tagamet® and Zantac®, are acid inhibitors. They act on acidproducing cells in the lining of the stomach. Formulations that control acid in this way are
now available as over-the-counter drugs.
Figure 4.10 Antacids. These products all serve
as acid-neutralizing agents in the stomach.
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4.4 OXIDATION-REDUCTION REACTIONS
Oxidation and Reduction
• Oxidation: Loss of electrons by a substance.
– An oxidation occurs when an atom or ion loses electrons.
• Reduction: Gain of electrons by a substance.
– A reduction occurs when an atom or ion gains electrons.
• One cannot occur without the other: Redox reaction
– Reaction in which electrons are transferred from one
reactant to another.
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4.4 OXIDATION-REDUCTION REACTIONS
Oxidation and Reduction
• Corrosion: An example of redox reactions.
– The conversion of a metal into a metal compound by a reaction
between the metal and some substance.
– When a metal corrodes, each metal atom loses electrons and so forms
a cation, which can combine with an anion to form an ionic compound.
– Example: The green coating on the Statue of Liberty (Cu2+ combined
with carbonate and hydroxide anions), rust (Fe3+ combined with oxide
and hydroxide anions), and silver tarnish (Ag+ combined with sulfide
anions).
Figure 4.11 Familiar corrosion products. (a) A green coating forms when copper is
oxidized. (b) Rust forms when iron corrodes. (c) A black tarnish forms as silver corrodes.
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4.4 OXIDATION-REDUCTION REACTIONS
Oxidation and Reduction
• The oxidations of alkali and alkaline earth metals.
– Quick reactions upon exposure to air.
• Cf) The reaction between iron and oxygen tends to be relatively slow.
– CaO formation: Ca is oxidized to Ca2+ and neutral O2 is transformed to
ions.
Figure 4.12 Oxidation of calcium metal by molecular oxygen. The oxidation involves
transfer of electrons from the calcium metal to the O2, leading to formation of CaO.
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4.4 OXIDATION-REDUCTION REACTIONS
Oxidation Numbers
• To determine if an oxidation–reduction reaction has
occurred, we assign an oxidation number to each
element in a neutral compound or charged entity.
• For neutral molecules and polyatomic ions, the oxidation
number of a given atom is a hypothetical charge.
– This charge is assigned by artificially dividing up the electrons
among the atoms in the molecule or ion.
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4.4 OXIDATION-REDUCTION REACTIONS
Rules for Assigning Oxidation Numbers
• Rules for assigning oxidation numbers:
– Elements in their elemental form have an oxidation number of 0.
– The oxidation number of a monatomic ion is the same as its charge.
– Nonmetals tend to have negative oxidation numbers, although some
are positive in certain compounds or ions.
• Oxygen has an oxidation number of −2, except in the peroxide ion, in
which it has an oxidation number of −1.
• Hydrogen is −1 when bonded to a metal, +1 when bonded to a nonmetal.
• Fluorine always has an oxidation number of −1.
• The other halogens have an oxidation number of −1 when they are
negative; they can have positive oxidation numbers, however, most
notably in oxyanions.
– The sum of the oxidation numbers in a neutral compound is 0.
– The sum of the oxidation numbers in a polyatomic ion is the
charge on the ion.
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4.4 OXIDATION-REDUCTION REACTIONS
Oxidation of Metals by Acids and Salts
• The general pattern of the reaction between a metal and
either an acid or a metal salt:
– Displacement reactions: The ion in solution is displaced (replaced)
through oxidation of an element.
– In displacement reactions, ions oxidize an element, and then the ions
are reduced.
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4.4 OXIDATION-REDUCTION REACTIONS
Oxidation of Metals by Acids and Salts
• Oxidation of metals by acids.
– Reaction of magnesium metal with hydrochloric acid.
• Molecular equation:
• Net ionic equation:
Figure 4.13 Reaction of
magnesium metal with
hydrochloric acid. The
metal is readily oxidized by
the acid, producing hydrogen
gas, H2(g), and MgCl2(aq).
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4.4 OXIDATION-REDUCTION REACTIONS
Oxidation of Metals by Acids and Salts
• Oxidation of metals by salts.
– Oxidation of iron metal by aqueous solutions of Ni2+ such as
Ni(NO3)2(aq):
• Molecular equation:
• Net ionic equation:
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4.4 OXIDATION-REDUCTION REACTIONS
The Activity Series
Figure 4.14 Reaction of copper metal with silver ion. When copper metal is placed in a solution
of silver nitrate, a redox reaction forms silver metal and a blue solution of copper(II) nitrate.
• In this reaction, silver ions oxidize copper metal:
Cu(s) + 2Ag+(aq) → Cu2+(aq) + 2Ag(s)
• The reverse reaction does not occur. Why not?
x Cu(s) + 2Ag+(aq)
Cu2+(aq) + 2Ag(s) →
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4.4 OXIDATION-REDUCTION REACTIONS
The Activity Series
• A list of metals arranged in order of decreasing ease of
oxidation.
– Elements higher on the activity series are more reactive; They
are more likely to exist as ions.
– Zn is oxidized by aqueous solution of Cu2+, but Ag is not.
– The elements
above hydrogen
will react with
acids to produce
hydrogen gas;
The metal is
Any metal on the
oxidized to a
list can be oxidized
by the ions of
cation.
elements below it.
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4.5 CONCENTRATIONS OF SOLUTIONS
Molarity
• The quantity of solute in a solution can matter to a
chemist.
• Concentration is used to designate the amount
of solute dissolved in a given quantity of solvent
or quantity of solution.
• Molarity is one way to measure the concentration
of a solution as the number of moles of solute in a
liter of solution:
Molarity (M) =
moles of solute
volume of solution in liters
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4.5 CONCENTRATIONS OF SOLUTIONS
Molarity
Figure 4.15 Preparing 0.250 L of a 1.00 M solution of CuSO4.
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4.5 CONCENTRATIONS OF SOLUTIONS
Expressing the Concentration of an Electrolyte
• When an ionic compound dissolves, the relative
concentrations of the ions in the solution depend on
the chemical formula of the compound.
– Ex1) NaCl: A 1.0 M solution of NaCl is 1.0 M in Na+ ions
and 1.0 M in Cl– ions.
– Ex2) Na2SO4: A 1.0 M solution of Na2SO4 is 2.0 M in Na+
ions and 1.0 M in SO42– ions.
• The concentration of an electrolyte solution can be
specified either in terms of the compound used to
make the solution (1.0 M Na2SO4) or in terms of the
ions in the solution (2.0 M Na+ and 1.0 M SO42–).
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4.5 CONCENTRATIONS OF SOLUTIONS
Interconverting Molarity, Moles, and Volume
Molarity (M) =
moles of solute
volume of solution in liters
• If we know any two of the three quantities in the
above equation, we can calculate the third.
• Molarity, therefore, is a conversion factor between
volume of solution and moles of solute:
– Ex 1) The number of moles of solute in 2.0 L of 0.200 M
HNO3 solution:
– Ex 2) The volume of 0.30 M HNO3 solution required to
supply 2.0 mol of HNO3 (Use the reciprocal of molarity in
the conversion: Liters = moles × 1/M):
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4.5 CONCENTRATIONS OF SOLUTIONS
Dilution
• Dilution is a process to obtain solutions of lower
concentrations by adding water.
Figure 4.16 Preparing 250 mL of 0.100 M CuSO4 by dilution of 1.00 M CuSO4.
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4.5 CONCENTRATIONS OF SOLUTIONS
Dilution
• How can we prepare 250.0 mL of 0.100 M cupric sulfate
(CuSO4) solution from a 1.00 M CuSO4 stock solution?
– The main point: When solvent is added to a solution, the
number of moles of solute remains unchanged.
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4.5 CONCENTRATIONS OF SOLUTIONS
Dilution
• Using the above equation, we can simply calculate the
volume of 1.00 M CuSO4 needed to prepare 250.0 mL of
0.100 M CuSO4 solution.
• If the volume of a diluted solution is known, the molarity
of the solution also can be determined from the equation.
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4.6 SOLUTION STOICHIOMETRY AND CHEMICAL ANALYSIS
Using Molarities in Stoichiometric Calculations
Figure 4.17 Procedure for solving
stoichiometry problems involving
reactions between a pure substance
A and a solution containing a known
concentration of substance B.
• Let’s consider a question: How many grams of Ca(OH)2
are needed to neutralize 25.0 mL of 0.100 M HNO3?
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4.6 SOLUTION STOICHIOMETRY AND CHEMICAL ANALYSIS
Titrations
• Titration: An analytical technique in which one can
calculate the concentration of a solute in a solution.
– Standard solution: A reagent solution of known concentration.
– Equivalence point: The point at which stoichiometrically
equivalent quantities are brought together.
– Acid-base indicator: A dye that changes color on passing the
equivalence point.
Figure 4.18 Procedure for titrating an acid against a standard solution of NaOH. The acid–base
indicator, phenolphthalein, is colorless in acidic solution but takes on a pink color in basic solution.
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4.6 SOLUTION STOICHIOMETRY AND CHEMICAL ANALYSIS
Titrations
• Knowing the volumes of both solutions and the concentration
of the standard solution we can calculate the concentration
of the unknown solution.
Figure 4.19 Procedure for determining the concentration of a solution from titration
with a standard solution.
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Chapter 4. Homework
Exercises 4.13
4.20
4.23
4.26
4.34
4.40
4.52
4.56
4.61
4.72
4.81
4.88
4.105
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