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Anderson - MCHS Solution Stoichiometry AP Chem. – Unit 02

Solution Stoichiometry

Unit Overview

Unit 2: Chapter 4 - Reactions in Aqueous Solutions Chapter 13 – Properties of Solutions Unit 2: Solution Stoichiometry Chapter 4) Aqueous Reactions (2 wk)

Nature of solutions, Molarity, Solution Stoichiometry Types of reactions, precipitation reactions, acid-base reactions, oxidation – reduction reactions.

Sec 4.1 – 4.2 Quiz – Solutions & Metathesis reactions Sec 4.3 Quiz – Acid & Bases Sec 4.4 Quiz - Redox Sec 4.5 -4.6 Quiz – Molarity & Solution Stoichiometry Chapter 4 Test Chapter 13) Properties of Solutions (2 wk)

Properties of solutions, mole fraction, molality, molarity, conversion between concentration units, dilution problems. Effects on solubility of temperature and pressure, Henry’s Law, Raoult’s Law, Boiling point elevation, freezing point depression, osmotic pressure. Molar mass determination from colligative properties.

Sec 13.1 – 13.3 Quiz – Solution process, classification & factors affecting solubility Sec 13.4Quiz – Ways of expressing Concentration and Conversions Sec 13.5 – 13.6 Quiz – Colligative Properties Chapter 13 Test Related Labs:

Acid – Base Titrations: (AP requirements # 6 – Standardization of a Solution Using a Primary Standard) and #7 Determination of Concentration by Acid-BaseTitration , Including a Weak Acid or a Weak Base)( Vonderbrink)

Oxidation – Reduction Titrations: (AP Requirement #8 – Determination of Concentration by Oxidation – Reduction Titration)( Vonderbrink)

Molar Mass by Freezing Point Depression: (AP requirements #4 – Determination of Molar Mass By Freezing Point Depression)( Vonderbrink)

Determining the Stoichiometry of Chemical Reactions: (AP requirements # 9 – Determination of Mass & Mole Relationships in a Chemical Reaction) (Vonderbrink)

Names and Formulas of Ionic Compounds (Virtual Chem Lab 2-5) Summer assignment

Exam 2


Unit Objectives

Chapter 4 Aqueous Reactions and Solution Stoichiometry

AP04-1,3-01 Predict whether a substance is a nonelectrolyte, strong electrolyte, or

weak electrolyte from its chemical behavior.

AP04-1,3-02 Predict the ions formed by electrolytes when they dissociate or ionize.

AP04-1,3-03 Identify substances as acids, bases, or salts

AP04-2-01 Use solubility rules to predict whether a precipitate forms when electrolyte

solutions are mixed.

AP04-2-02 Predict the products of metathesis reactions (including both neutralization

and precipitation reactions) and write balanced chemical equations for

them.

AP04-2-03 Identify the spectator ions and write the net ionic equations for solution

reactions, starting with their molecular equations.

AP04-4-01 Determine whether a chemical reaction involves oxidation and reduction.

AP04-4-02 Assign oxidation numbers to atoms in molecules and ions.

AP04-4-03 Use the activity series to predict whether a reaction will occur when a

metal is added to an aqueous solution of either a metal salt or an acid, and

write the balanced molecular and net ionic equations for the reaction.

AP04-5-01 Calculate molarity, solution volume, or number of moles of solute given

any two of these quantities.

AP04-5-02 Calculate the volume of a more concentrated solution that muse be

diluted to obtain a given quantity of a more dilute solution.

AP04-6-01 Calculate the volume of a solution required to react with a volume of a

different solution using molarity and the stoichiometry of the reaction.

AP04-6-02 Calculate the amount of a substance required to react with a given volume

of a solution using molarity and the stoichiometry of the reaction.

AP04-6-03 Calculate the concentration or mass of solute in a sample from titration

data.

Chapter 13 Properties of Solutions

AP13-1,2,3-01 Describe the energy changes that occur in the solution process in terms of

the solute-solute, solvent-solvent, and solute-solvents attractive forces;

describe the role of disorder in the solution process.

AP13-1,2,3-02 Rationalize the solubilities of substances in various solvents in terms of the

molecular structure and intermolecular forces.

AP13-1,2,3-03 Describe the effects of pressure and temperature on solubilities.

AP13-4-01 Define mass percent, parts per million, mole fraction, molarity and molality


and calculate concentrations in any of these concentration units.

AP13-4-02 Convert concentration in one concentration unit into any other (given

density of the solution where necessary)

AP13-5-01 Determine the concentration and molar mass of a nonvolatile

nonelectrolyte form its effect on the colligative property of a solution.

AP13-5-02 Explain the difference between the magnitudes of changes in colligative

properties caused by electrolytes compared to those caused by

nonelectrolyte.

AP13-5-03 Describe the effects of solute concentration on the vapor pressure, boiling

point, freezing point, and osmotic pressure of a solution, and calculate any

of the properties given appropriate concentration data.

AP13-6-01 Describe how a colloid differs from a true solution.

•

Schedule & Assignments

Text Readings, Notes & Practice --- Option 1 Complete the readings before discussions in class. Carefully read through the sample

exercises and attempt the practice exercise in the text. Pay particular attention to vocabulary

& concept reviews. After class discussion – practice a min. of 5 EOC’s for each section.

(Correct answers are in the back of the book or on AP Chemistry website) Do what you need

to MASTER each concept. You should be able to summarize the main concepts in the section

both conceptually and mathematically.

Assignments for Unit: Option 1 Chapter 4 Aqueous Reactions & Solution

Stoichiometry

Text Section EOCQ’s Date

completed Topic

4-1 General Properties of Aqueous Solutions Pg119 - 123 11 - 20

4-2 Precipitation Reactions Pg 124 – 128 *must know

table 4.1 pg 125 21 - 28

4-3 Acid – Base Reactions Pg 128 - 135 29 - 44

Chemistry Put to Work “Antacids” pg 135

4 -4 Oxidation Reduction reactions Pg 135 – 142 * must know

table 4.5 pg 141 45 - 58

A Closer Look The Aura of Gold pg143

Strategies in Chemistry Analyzing Chemical Reactions

4-5 Concentrations of Solutions Pg 142 - 149 59 - 78

4-6 SolutionStoichiometry and Chemical

Analysis Pg 149 - 154 79 - 90

Chapter 4 Summary & Key Terms Pg 155 - 156

Visualizing Concepts Pg 156 1-10

Additional Exercises Pg 161 92,94,97,100


,101,104

Integrative Exercises Pg 162 106,109,113,

115

Chapter 13 Properties of Solutions Text Section EOCQ’s

Topic

13-1 the Solution Process Pg 527 - 533

A Closer Look Hydrates pg 533 - 534 9-16

13-2 Saturated Solutions and solubility Pg 534 - 535 17-36

13-3 Factors Affecting Solubility Pg 535 - 541 89

Chemistry and Life Fat & Water soluble Vitamins

pg 538 17-36

Chemistry and Life Blood gases and deep sea

diving pg 540 33-36,93,94

13-4 Ways of Expressing Concentration Pg 542 - 546 37-60

13-5 Colligative Properties Pg 546 - 556 96

A Closer Look

Ideal solutions with two or

more volatile components

Pg 548

37-60

A Closer Look Colligative Properties of

Electrolyte Solutions Pg 554 37-60

13-6 Colloids Pg 556 - 561

Chemistry & Life

Sickle-Cell Anemia Pg 559

Chapter Review, Key Skills & Key Equations Pg 561 -563

Visualizing Concepts Pg 563 1-12

Additional Exercises Pg 569-570

81,83,85,87,

90,93,96,99,

102

Intergrative Exercises Pg570 - 571 104,108,111


Chapter 4: Reactions in Aqueous Solutions

Common Student Misconceptions

Chapter 4. Aqueous Reactions and Solution Stoichiometry

• Molarity is moles of solute per liter of solution, not per liter of solvent.

• Students sometimes use moles instead of molarity in MinitialVinitial = MfinalVfinal.

• Students often disregard rules for significant figures when calculating or using molarities.

• Students sometimes think that water is a good conductor. • Students sometimes have a problem with the arbitrary difference between strong and weak

electrolytes. • Students often think that nonelectrolytes produce no ions in aqueous solution at all. • Students sometimes cannot tell the difference between dissolution and dissociation.

• The symbols (equilibrium) and (resonance) are often confused.

• Students often do not see that the net ionic equation for the reaction between strong acids

and strong bases is always H+(aq) + OH

–(aq) H2O(l).

• Weaknesses in recollection of ionic nomenclature and the structure of common ions often

make it difficult for students to write molecular, complete ionic, and net ionic equations for

metathesis reactions.

• Students try to split polyatomic ions into smaller ions when they write net ionic equations.

• Students often think that a compound consisting of nonmetals only must be molecular

[counter-example: (NH4)2SO4, which is ionic!]

• Students do not realize that insoluble really means poorly soluble.

• Students do not appreciate the difference between equivalence point and end point.

• Students usually think that an oxidation necessarily involves a reaction with oxygen and/or

addition of an atom of oxygen to the formula.

• Students often think that all atoms of the same element must have the same oxidation

number and that this number is uniquely related to the atom’s location in the periodic table.

• It is important that students know that titrations can be conducted not only using acids and

bases, but also in precipitation and oxidation–reduction reactions.

4.1 General Properties of Aqueous Solutions (EOCQ’S11-20)

• A solution is a homogeneous mixture of two or more substances.

• A solution is made when one substance (the solute) is dissolved in another (the

solvent).

• The solute is the substance that is present in the smallest amount.

• Solutions in which water is the solvent are called aqueous solutions.


Electrolytic Properties

• All aqueous solutions can be classified in terms of whether or not they conduct

electricity.

• If a substance forms ions in solution, then the substance is an electrolyte and the

solution conducts electricity. An example is NaCl.

• If a substance does not form ions in solution, then the substance is a

nonelectrolyte and the solution does not conduct electricity. Examples are

sucrose and water.

Ionic Compounds in Water

• When an ionic compound dissolves in water, the ions are said to dissociate.

• This means that in solution, the solid no longer exists as a well-ordered

arrangement of ions in contact with one another.

• Instead, each ion is surrounded by several water molecules; it is called

an aqueous ion or a hydrated ion.

• This tends to stabilize the ions in solution and prevent cations and anions

from recombining.

• The positive ions have the oxygen atoms of water pointing towards the

ion; negative ions have the hydrogen atoms of water pointing towards the ion.

• The transport of ions through the solution causes electric current to flow

through the solution.

Molecular Compounds in Water

• When a molecular compound (e.g. CH3OH ) dissolves in water, only a very

limited number of ions are formed.

• Therefore, there is nothing in the solution to transport electric charge and the

solution does not conduct electricity.

• There are some important exceptions.

• For example, NH3(g) reacts with water to form NH4+(aq) and OH

– (aq).

• For example, HCl(g) in water ionizes to form H+(aq) and Cl

– (aq).


Strong and Weak Electrolytes

• Compounds whose aqueous solutions conduct electricity well are called strong

electrolytes.

• These substances exist in solution mostly as ions.

• Example: NaCl

NaCl(aq) Na+(aq) + Cl

–(aq)

• The single arrow indicates that the Na+ and Cl

– ions have no

tendency to recombine to form NaCl.

• In general, soluble ionic compounds are strong electrolytes.

• Other strong electrolytes include strong acids and soluble strong bases.

• Compounds whose aqueous solutions conduct electricity poorly are called weak

electrolytes

• These substances exist as a mixture of ions and un-ionized molecules in

solution.

• The predominant form of the solute is the un-ionized molecule.

• Example: acetic acid, HC2H3O2

HC2H3O2(aq) H+(aq) + C2H3O2

–(aq)

• The double arrow means that the reaction is significant in both

directions.

• It indicates that there is a balance between the forward and reverse

reactions.

• This balance produces a state of chemical equilibrium.

FORWARD REFERENCES:

• Double arrows () will be used in the chapter on chemical equilibria (Ch. 15) and

beyond.

• Strong vs. weak electrolytes will come up in chapters on acid-base and solubility

Equilibria. (Ch. 16 and 17) as well as in electrochemistry (Ch. 20).

• Equilibria involving insoluble or poorly soluble compounds and their ions will be

discussed in more detail in Ch. 17 (section 4).

• Dissolving of substances in solvent and properties of solution will be discussed in

Ch. 13.

• Interactions between ions and molecules of a solvent (ion–dipole interactions) will

be further discussed in Ch. 11 and Ch. 13.


4.2 Precipitation Reactions (EOCQ’S 21-28)

• Reactions that result in the formation of an insoluble product are known as

precipitation reactions.

• A precipitate is an insoluble solid formed by a reaction in solution.

• Example: Pb(NO3)2(aq) + 2KI(aq) PbI2(s) + 2KNO3(aq)

Solubility Guidelines for Ionic Compounds

• The solubility of a substance at a particular temperature is the amount of that

substance that can be dissolved in a given quantity of solvent at that temperature.

• A substance with a solubility of less than 0.01 mol/L is regarded as being

insoluble.

• Experimental observations have led to empirical guidelines for predicting

solubility.

• Solubility guidelines for common ionic compounds in water:

• Compounds containing alkali metal ions or ammonium ions are soluble.

• Compounds containing NO3– or C2H3O2

– are soluble.

• Compounds containing Cl–, Br

– or I

– are soluble.

• Exceptions are the compounds of Ag+, Hg2

2+, and Pb

2+.

• Compounds containing SO42–

are soluble.

• Exceptions are the compounds of Sr2+

, Ba2+

, Hg22+

, and Pb2+

.

• Compounds containing S2–

are insoluble.

• Exceptions are the compounds of NH4+, the alkali metal cations,

and Ca2+

, Sr2+

, and Ba2+

.

• Compounds of CO32–

or PO43–

are insoluble.

• Exceptions are the compounds of NH4+ and the alkali metal

cations.

• Compounds of OH– are insoluble.

• Exceptions are the compounds of NH4+, the alkali metal cations,

and Ca2+

, Sr2+

, and Ba2+

.


Exchange (Metathesis) Reactions

• Exchange reactions, or metathesis reactions, involve swapping ions in solution:

AX + BY AY + BX.

• Many precipitation and acid–base reactions exhibit this pattern.

Ionic Equations

• Consider 2KI(aq) + Pb(NO3)2(aq) PbI2(s) + 2KNO3(aq).

• Both KI(aq) + Pb(NO3)2(aq) are colorless solutions. When mixed, they form a

bright yellow precipitate of PbI2 and a solution of KNO3.

• The final product of the reaction contains solid PbI2, aqueous K+, and aqueous

NO3– ions.

• Sometimes we want to highlight the reaction between ions.

• The molecular equation lists all species in their complete chemical forms:

Pb(NO3)2(aq) + 2KI(aq) PbI2(s) + 2KNO3(aq)

• The complete ionic equation lists all strong soluble electrolytes in the reaction as

ions:

Pb2+

(aq) + 2NO3–(aq) + 2K

+(aq) + 2I

–(aq)

PbI2(s) + 2K+(aq) + 2NO3

–(aq)

• Only strong electrolytes dissolved in aqueous solution are written in

ionic form.

• Weak electrolytes and nonelectrolytes are written in their complete

chemical form.

• The net ionic equation lists only those ions which are not common on both sides

of the reaction:

Pb2+

(aq) + 2I–(aq) PbI2(s)

• Note that spectator ions, ions that are present in the solution but play no

direct role in the reaction, are omitted in the net ionic equation. FORWARD REFERENCES:

• Net ionic equations will be frequently used in chapters dealing with acid–base

reactions (Ch. 16 and Ch. 17) as well as in electrochemistry (Ch. 20, Appendix E)

• Equilibria involving insoluble or poorly soluble compounds and their ions will be

discussed in more detail in Ch. 17 (section 17.4).


4.3 Acid–Base Reactions (EOCQ’S 29-44)

Acids

• Acids are substances that are able to ionize in aqueous solution to form H+.

• Ionization occurs when a neutral substance forms ions in solution.

An example is HC2H3O2 (acetic acid).

• Since H+ is a naked proton, we refer to acids as proton donors and bases as proton acceptors.

• Common acids are HCl, HNO3, vinegar, and vitamin C.

• Acids that ionize to form one H+ ion are called monoprotic acids.

• Acids that ionize to form two H+ ions are called diprotic acids.

Bases

• Bases are substances that accept or react with the H+ ions formed by acids.

• Hydroxide ions, OH–, react with the H

+ ions to form water:

H+(aq) + OH

–(aq) H2O(l)

• Common bases are NH3 (ammonia), Draino, and milk of magnesia.

• Compounds that do not contain OH– ions can also be bases.

• Proton transfer to NH3 (a weak base) from water (a weak acid) is an example of an acid–

base reaction.

• Since there is a mixture of NH3, H2O, NH4+, and OH

– in solution, we write

NH3(aq) + H2O(l) NH4+(aq) + OH

–(aq)

Strong and Weak Acids and Bases

• Strong acids and strong bases are strong electrolytes.

• They are completely ionized in solution.

• Strong bases include: Group 1A metal hydroxides, Ca(OH)2, Ba(OH)2, and Sr(OH)2.

• Strong acids include: HCl, HBr, HI, HClO3, HClO4, H2SO4, and HNO3.

• We write the ionization of HCl as:

HCl H+ + Cl

–

• Weak acids and weak bases are weak electrolytes.

• They are partially ionized in aqueous solution.

• HF(aq) is a weak acid; most acids are weak acids.

• We write the ionization of HF as:

HF + F

–

Identifying Strong and Weak Electrolytes

• Compounds can be classified as strong electrolytes, weak electrolytes, or nonelectrolytes by

looking at their solubility.

• Strong electrolytes:

• Soluble ionic compounds are strong electrolytes.

• Molecular compounds that are strong acids are strong electrolytes.


• Weak electrolytes:

• Weak acids and bases are weak electrolytes.

• Nonelectrolytes:

• All other compounds, including water.

Neutralization Reactions and Salts

• A neutralization reaction occurs when an acid and a base react:

• HCl(aq) + NaOH(aq) H2O(l) + NaCl(aq)

• (acid) + (base) (water) + (salt)

• In general an acid and a base react to form a salt.

• A salt is any ionic compound whose cation comes from a base and anion from an acid.

• The other product, H2O, is a common nonelectrolyte.

• A typical example of a neutralization reaction is the reaction between an acid and a metal

hydroxide:

• Mg(OH)2 (milk of magnesia) is a suspension.

• As HCl is added, the magnesium hydroxide dissolves, and a clear solution

containing Mg2+

and Cl– ions is formed.

• Molecular equation:

Mg(OH)2(s) + 2HCl(aq) MgCl2(aq) + 2H2O(l)

• Net ionic equation:

Mg(OH)2(s) + 2H+(aq) Mg

2+(aq) + 2H2O(l)

• Note that the magnesium hydroxide is an insoluble solid; it appears in

the net ionic equation.

Acid–Base Reactions with Gas Formation

• There are many bases besides OH– that react with H

+ to form molecular compounds.

• Reaction of sulfides with acid gives rise to H2S(g).

• Sodium sulfide (Na2S) reacts with HCl to form H2S(g):


Na2S(aq) + 2HCl(aq) H2S(g) + 2NaCl(aq)


2H+(aq) + S

2–(aq) H2S(g)

• Carbonates and hydrogen carbonates (or bicarbonates) will form CO2(g) when

treated with an acid.

• Sodium bicarbonate (NaHCO3; baking soda) reacts with HCl to form bubbles

of CO2(g):


NaHCO3(s) + HCl(aq) NaCl(aq) + H2CO3(aq) H2O(l) + CO2(g) + NaCl(aq)


H+(aq) + HCO3

–(aq) H2O(l) + CO2(g)

• Ammonium salts will form NH3(g) when treated with a hydroxide.


• Ammonium nitrate (NH4NO3) reacts with NaOH(aq) to form a gas with a

characteristic ammonia, NH3(g), smell.


NH4NO3 (aq) + NaOH(aq) NH4OH(aq) + NaNO3(aq) H2O(l) +

NH3(g) + NaNO3(aq)


NH4+(aq) + OH

–(aq) H2O(l) + NH3(g)

FORWARD REFERENCES:

• Strong acids and bases will be revisited in Ch. 16.

• Strong acids and bases will be used as titrants in acid–base titrations (Ch. 17)

• Equilibria involving weak acids and bases will be further discussed in Ch. 16 and 17.

• Environmental impact of weak acid equilibria will be discussed on Ch. 18.

4.4 Oxidation–Reduction Reactions (EOCQ’S 45-58)

Oxidation and Reduction

• Oxidation-reduction, or redox, reactions involve the transfer of electrons

between reactants.

• When a substances loses electrons, it undergoes oxidation:

Ca(s) + 2H+(aq) Ca

2+(aq) + H2(g)

• The neutral Ca has lost two electrons to 2H+ to become Ca

2+.

• We say Ca has been oxidized to Ca2+

.

• When a substance gains electrons, it undergoes reduction:

2Ca(s) + O2(g) 2CaO(s).

• In this reaction the neutral O2 has gained electrons from the Ca to

become O2–

in CaO.

• We say O2 has been reduced to O2–

.

• In all redox reactions, one species is reduced at the same time as another is

oxidized.


Oxidation Numbers

• Electrons are not explicitly shown in chemical equations.

• Oxidation numbers (or oxidation states) help up keep track of electrons during

chemical reactions.

• Oxidation numbers are assigned to atoms using specific rules.

• For an atom in its elemental form, the oxidation number is always zero.

• For any monatomic ion, the oxidation number equals the charge on the

ion; positive for metals and negative for nonmetals.

• The oxidation number of oxygen is usually –2.

• The major exception is in peroxides (containing the O22–

ion).

• The oxidation number of hydrogen is +1 when bonded to nonmetals and

–1 when bonded to metals.

• The oxidation number of fluorine is –1 in all compounds. The other

halogens have an oxidation number of –1 in most binary compounds.

• The sum of the oxidation numbers of all atoms in a neutral compound is

zero.

• The sum of the oxidation numbers in a polyatomic ion equals the charge

of the ion.

• The oxidation of an element is evidenced by an increase in its oxidation number;

reduction is accompanied by a decrease in an oxidation number.

Oxidation of Metals by Acids and Salts

• The reaction of a metal with either an acid or a metal salt is called a displacement

reaction.

• The general pattern is:

A + BX AX + B

• Example: It is common for metals to produce hydrogen gas when they

react with acids. Consider the reaction between Mg and HCl:

Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)

• In the process the metal is oxidized and the H+ is reduced.


• Example: It is possible for metals to be oxidized in the presence of a

salt:

Fe(s) + Ni(NO3)2(aq) Fe(NO3)2(aq) + Ni(s)

• The net ionic equation shows the redox chemistry well:

Fe(s) + Ni2+

(aq) Fe2+

(aq) + Ni(s)

• In this reaction iron has been oxidized to Fe2+

, while the Ni2+

has

been reduced to Ni.

• Always keep in mind that whenever one substance is oxidized, some other

substance must be reduced.

The Activity Series

• We can list metals in order of decreasing ease of oxidation.

• This list is an activity series.

• The metals at the top of the activity series are called active metals.

• The metals at the bottom of the activity series are called noble metals.

• A metal in the activity series can only be oxidized by a metal ion below it.

• If we place Cu into a solution of Ag+ ions, then Cu

2+ ions can be formed because

Cu is above Ag in the activity series:

Cu(s) + 2AgNO3(aq) Cu(NO3)2(aq) + 2Ag(s)

or

Cu(s) + 2Ag+(aq) Cu

2+(aq) + 2Ag(s)

FORWARD REFERENCES:

• Oxidation numbers will be frequently used in electrochemistry (Ch. 20, Appendix E,

etc.)

• Balancing of redox reactions will be covered in Ch. 20


4.5 Concentrations of Solutions (EOCQ’S 59-78)

• The term concentration is used to indicate the amount of solute dissolved in a

given quantity of solvent or solution.

Molarity

• Solutions can be prepared with different concentrations by adding different

amounts of solute to solvent.

• The amount (moles) of solute per liter of solution is the molarity or molar

concentration (symbol M) of the solution:

• By knowing the molarity of a quantity of liters of solution, we can easily calculate

the number of moles (and, by using molar mass, the mass) of solute.

• Consider weighed copper sulfate, CuSO4 (39.9 g, 0.250 mol) placed in a 250. mL

volumetric flask. A little water is added and the flask swirled to ensure the copper

sulfate dissolves. When all the copper sulfate has dissolved, the flask is filled to

the mark with water.

• The molarity of the solution is

0.250 mol CuSO4 / 0.250 L solution = 1.00 M.

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.

• Example: for a 1.0 M solution of NaCl:

• The solution is 1.0 M in Na+ ions and 1.0 M in Cl

– ions.

• Example: for a 1.0 M solution of Na2SO4:

• The solution is 2.0 M in Na+ ions and 1.0 M in SO4

2– ions.

Interconverting Molarity, Moles, and Volume

• The definition of molarity contains three quantities: molarity, moles of solute, and

liters of solution.

• If we know any two of these, we can calculate the third.

• Dimensional analysis can be helpful in these calculations.

solution of liters

solute moles=Molarity


Dilution

• A solution in concentrated form (stock solution) is mixed with solvent to obtain a

solution of lower solute concentration.

• This process is called dilution.

• An alternate way of making a solution is to take a solution of known molarity and

dilute it with more solvent.

• Since the number of moles of solute remains the same in the concentrated and

diluted forms of the solution, we can show:

MconcVconc = MdilVdil

• An alternate form of this equation is:

MinitialVinitial = MfinalVfinal

FORWARD REFERENCES:

• Molarity will be used throughout the course as the most common form of

concentration.

• The concept of molarity is not limited to solutions; one can calculate molarity for

gases and use them in Kc expressions in Ch. 15 and beyond.

• Molarity can be converted into other concentrations (molality, normality, ppm, etc.)

as shown in Ch. 13; molarity will be used to calculate osmotic pressure (section

13.5).

• In some later chapters (14 and beyond) molarity will be also symbolized as [solute].

• Dilutions will come up in select acid-base equilibrium problems in Ch. 16 and in

titrations (Ch. 17).

4.6 Solution Stoichiometry and Chemical Analysis (EOCQ’S 79-90)

• In approaching stoichiometry problems:

• recognize that there are two different types of units:

• laboratory units (the macroscopic units that we measure in lab)

and

• chemical units (the microscopic units that relate to moles).

• Always convert the laboratory units into chemical units first.

• Convert grams to moles using molar mass.

• Convert volume or molarity into moles using M = mol/L.

• Use the stoichiometric coefficients to move between reactants and

products.

• This step requires the balanced chemical equation.

• Convert the laboratory units back into the required units.

• Convert moles to grams using molar mass.

• Convert moles to molarity or volume using M = mol/L.


Titrations

• A common way to determine the concentration of a solution is via titration.

• We determine the concentration of one substance by allowing it to undergo a

specific chemical reaction, of known stoichiometry, with another substance whose

concentration is known (standard solution).

• Monoprotic acids and bases react with each other in a stoichiometric ratio of 1:1.

• Example: Suppose we know the molarity of an NaOH solution and we want to

find the molarity of an HCl solution.

• What do we know?

• molarity of NaOH, volume of HCl

• What do we want?

• molarity of HCl

• What do we do?

• Take a known volume of the HCl solution (i.e., 20.00 mL) and measure

the number of mL of 0.100 M NaOH solution required to react completely

with the HCl solution.

• The point at which stoichiometrically equivalent quantities of NaOH and

HCl are brought together is known as the equivalence point of the

titration.

• The equivalence point is a theoretical concept that can be calculated “on

paper” only.

• In a titration we often use an acid–base indicator to allow us to determine

when the equivalence point of the titration has been reached.

• Acid–base indicators change color at the end point of the titration.

• The indicator is chosen so that the end point

corresponds to the equivalence point of the titration; the end point is

determined experimentally.

• What do we get?

• We get the volume of NaOH. Since we already have the molarity

of the NaOH, we can calculate moles of NaOH.

• What is the next step?

• We also know HCl + NaOH NaCl + H2O; note the 1:1

stoichiometric ratio between HCl and NaOH.

• Therefore, we know moles of HCl.

• Can we finish?


• Knowing mol (HCl) and volume of HCl, we can calculate the

molarity.

FORWARD REFERENCES:

• Acid–base titrations will be discussed in detail in Ch. 17.

ADDITIONAL EXERCISES: PAGE 161 Q’S 92,94,97,100,101,104

INTEGRATIVE EXERCISES: PAGE 162 Q’S 106, 109, 113, 115

Chapter 13: Properties of Solutions

Common Student Misconceptions

• Students often confuse dilute and concentrated; weak and strong are often confused.

• Students often do not appreciate the driving forces behind the formation of a solution.

• Students often confuse dissolution with melting.

• Even insoluble compounds dissolve to some extent in water.

• Students do not realize that crystallization is the reverse of dissolution.

• Many student think that solutions can only be made either by mixing two liquids or

dissolving a solid in a liquid.

• Students often confuse solution with solvation.

• Increased temperature often but not always improves solubility.

• Students often think that every mixture is a solution.

• Students often do not appreciate how unusual water is.

• Students often forget that calculations of molality require the mass of solvent, not

solution.

• Students often do not realize that colloids, like solutions, can occur in all three states of

matter.

13.1 The Solution Process

(EOCQ’S 9-16)

• A solution is a homogeneous mixture of solute and solvent. • Solutions may be gases, liquids, or solids. • Each substance present is a component of the solution.

o The solvent is the component present in the largest amount.

o The other components are the solutes.


The Effect of Intermolecular Forces

• Intermolecular forces become rearranged in the process of making solutions

with condensed phases. • Consider NaCl (solute) dissolving in water (solvent):

o Water molecules orient themselves on the NaCl crystals.

o H-bonds between the water molecules have to be broken.

o NaCl dissociates into Na+ and Cl–.

o Ion–dipole forces form between the Na+ and the negative end of the water

dipole.

Similar ion–dipole interactions form between the Cl– and the positive

end of the water dipole.

o Such an interaction between solvent and solute is called solvation.

If water is the solvent, the interaction is called hydration.

Energy Changes and Solution Formation

• There are three steps involving energy in the formation of a solution:

• Separation of solute molecules (∆H1),

• Separation of solvent molecules (∆H2), and

• Formation of solute–solvent interactions (∆H3).

• We define the enthalpy change in the solution process as:

∆Hsoln = ∆H1 + ∆H2 + ∆H3

• ∆Hsoln can either be positive or negative depending on the intermolecular

forces.

• To determine whether ∆Hsoln is positive or negative, we consider the

strengths of all solute–solute, solvent–solvent, and solute–solvent interactions:

• Breaking attractive intermolecular forces is always endothermic.

• ∆H1 and ∆H2 are both positive.

•Forming attractive intermolecular forces is always exothermic.

•∆H3 is always negative.

• It is possible to have either ∆H3 > (∆H1 + ∆H2) or ∆H3 < (∆H1 + ∆H2).

• Examples:

o MgSO4 added to water has ∆Hsoln = –91.2 kJ/mol.

o NH4NO3 added to water has ∆Hsoln = + 26.4 kJ/mol.

o MgSO4 is often used in instant heat packs and NH4NO3 is often used in

instant cold packs.


• How can we predict if a solution will form?

• In general, solutions form if the ∆Hsoln is negative.

• If ∆Hsoln is too endothermic a solution will not form.

• “Rule of thumb”: polar solvents dissolve polar solutes.

•Nonpolar solvents dissolve nonpolar solutes.

• Consider the process of mixing NaCl in gasoline.

•Only weak interactions are possible because gasoline is nonpolar.

•These interactions do not compensate for the separation of ions from one

another.

Result: NaCl doesn't dissolve to any great extent in gasoline.

Consider the process of mixing water in octane (C8H18).

•Water has strong H-bonds.

•The energy required to break these H-bonds is not compensated for by

interactions between water and octane.

Result: water and octane do not mix.

Solution Formation, Spontaneity, and Disorder

• A spontaneous process occurs without outside intervention.

• When the energy of the system decreases (e.g., dropping a book and allowing it to

fall to a lower potential energy), the process is spontaneous.

• Spontaneous processes tend to be exothermic.

• However, some spontaneous processes do not involve the movement of the

system to a lower energy state (e.g., an endothermic reaction).

• Some endothermic processes occur spontaneously.

• The amount of randomness or disorder of the system is given by its entropy.

• In most cases, solution formation is favored by the increase in entropy that

accompanies mixing.

• Example: A mixture of CCl4 and C6H14 is less ordered than the two

separate liquids.

• Therefore, they spontaneously mix even though ∆Hsoln is very close to

zero.

• A solution will form unless the solute–solute or solvent–solvent interactions

are too strong relative to solute–solvent interactions.


Solution Formation and Chemical Reactions

• Some solutions form by physical processes and some by chemical processes.

• Consider:

Ni(s) + 2HCl(aq) NiCl2(aq) + H2(g)

• Note that the chemical form of the substance being dissolved has changed

during this process (Ni NiCl2)

• When all the water is removed from the solution, no Ni is found, only

NiCl2•6H2O remains.

• Therefore, the dissolution of Ni in HCl is a chemical process.

• By contrast: NaCl(s) + H2O (l) Na+(aq) + Cl–(aq).

• When the water is removed from the solution, NaCl is found.

• Therefore, NaCl dissolution is a physical process.

13.2 Saturated Solutions and Solubility (EOCQ’S 17-36)

• As a solid dissolves, a solution forms:

• Solute + solvent solution

• The opposite process is crystallization.

• Solution solute + solvent

• If crystallization and dissolution are in equilibrium with undissolved solute, the

solution is saturated.

• There will be no further increase in the amount of dissolved solute.

• Solubility is the amount of solute required to form a saturated solution.

• A solution with a concentration of dissolved solute that is less than the

solubility is said to be unsaturated.

• A solution is said to be supersaturated if more solute is dissolved than in a

saturated solution.

13.3 Factors Affecting Solubility (EOCQ’S 89) • The tendency of a substance to dissolve in another depends on:

• the nature of the solute.

• the nature of the solvent.

• the temperature.

• the pressure (for gases).

Solute–Solvent Interactions

• Intermolecular forces are an important factor in determining solubility of a solute

in a solvent.

• The stronger the attraction between solute and solvent molecules, the

greater the solubility.


• For example, polar liquids tend to dissolve in polar solvents.

• Favorable dipole–dipole interactions exist (solute–solute, solvent–

solvent, and solute–solvent).

• Pairs of liquids that mix in any proportions are said to be miscible.

• Example: Ethanol and water are miscible liquids.

• In contrast, immiscible liquids do not mix significantly.

• Example: Gasoline and water are immiscible.

• Consider the solubility of alcohols in water.

• Water and ethanol are miscible because the broken hydrogen bonds in both

pure liquids are re-established in the mixture.

• However, not all alcohols are miscible with water.

• Why?

• The number of carbon atoms in a chain affects solubility.

• The greater the number of carbons in the chain, the more the molecule

behaves like a hydrocarbon.

• Thus, the more C atoms in the alcohol, the lower its solubility

in water.

• Increasing the number of –OH groups within a molecule increases its

solubility in water.

• The greater the number of –OH groups along the chain, the more

solute-water H-bonding is possible.

• Generalization: “like dissolves like”.

• Substances with similar intermolecular attractive forces tend to be soluble in

one another.

• The more polar bonds in the molecule, the better it dissolves in a

polar solvent.

• The less polar the molecule the less likely it is to dissolve in a polar

solvent and the more likely it is to dissolve in a nonpolar solvent.

• Network solids do not dissolve because the strong intermolecular forces in the

solid are not reestablished in any solution.

Pressure Effects

• The solubility of a gas in a liquid is a function of the pressure of the gas over the

solution.

• Solubilities of solids and liquids are not greatly affected by pressure.

• With higher gas pressure, more molecules of gas are close to the surface of the

solution and the probability of a gas molecule striking the surface and entering the

solution is increased.

• Therefore, the higher the pressure, the greater the solubility.

• The lower the pressure, the smaller the number molecules of gas close to the

surface of the solution resulting in a lower solubility.


• The solubility of a gas is directly proportional to the partial pressure of the

gas above the solution.

• This statement is called Henry's law.

• Henry's law may be expressed mathematically as:

Sg=kPg

• Where Sg is the solubility of gas, Pg the partial pressure, k = Henry’s

law constant.

• Note that the Henry's law constant differs for each solute–solvent pair

and differs with temperature.

• An application of Henry's law is the preparation of carbonated soda.

• Carbonated beverages are bottled under PCO2 > 1 atm.

• As the bottle is opened, PCO2 decreases and the solubility of CO2

decreases.

• Therefore, bubbles of CO2 escape from solution.

Temperature Effects

• Experience tells us that sugar dissolves better in warm water than

in cold water.

• As temperature increases, solubility of solids generally increases.

• Sometimes solubility decreases as temperature increases

[e.g., Ce2(SO4)3].

• Experience tells us that carbonated beverages go flat as they get warm.

• Gases are less soluble at higher temperatures.

• An environmental application of this is thermal pollution.

• Thermal pollution: if lakes get too warm, CO2 and O2 become less soluble

and are not available for plants or animals.

• Fish suffocate.

13.4 Ways of Expressing Concentration

(EOCQ’S 37-60) • All methods involve quantifying the amount of solute per amount

of solvent (or solution).

• Concentration may be expressed qualitatively or quantitatively.

• The terms dilute and concentrated are qualitative ways to

describe concentration.

• A dilute solution has a relatively small concentration of solute.

• A concentrated solution has a relatively high concentration

of solute.


• Quantitative expressions of concentration require specific information regarding

such quantities as masses, moles, or liters of the solute, solvent, or solution.

• The most commonly used expressions for concentration are:

• mass percentage.

• mole fraction.

• molarity.

• molality.

Mass Percentage, ppm, and ppb

• Mass percentage is one of the simplest ways to express concentration.

• By definition:

• Similarly, parts per million (ppm) can be expressed as the number of mg of

solute per kilogram of solution.

• By definition:

• The density of a very dilute aqueous solution is similar to

that of pure water (1g/ml)

• If the density of the solution is 1g/ml, then 1 ppm = 1 mg solute per liter of

solution.

• We can extend this again!

• Parts per billion (ppb) can be expressed as the number of µg of solute per

kilogram of solution.

• By definition:

• If the density of the solution is 1g/ml, then 1 ppb = 1 µg solute per liter of

solution.

Mole Fraction, Molarity, and Molality

• Common expressions of concentration are based on the number of moles of one or

more components.

• Recall that mass can be converted to moles using the molar mass.

•Recall:

• Note that mole fraction has no units.

• Note that mole fractions range from 0 to 1.

100solution of mass total

solnin component of masscomponent of % Mass


solnin component of masscomponent of (ppm)million per Parts


solnin component of masscomponent of (ppb)billion per Parts

componentsallofmolestotal

componentofmoles,componentoffractionMole X


• Recall:

• Note that molarity will change with a change in temperature (as the

solution volume increases or decreases).

• We can define molality (m), yet another concentration unit:

• Molality does not vary with temperature.

• Note that converting between molarity (M) and molality (m) requires

density.

• The molarity and molality of dilute solutions are often very similar.

13.5 Colligative Properties (EOCQ’S 96) • Colligative properties depend on number of solute particles.

• There are four colligative properties to consider:

• vapor pressure lowering (Raoult's Law).

• boiling point elevation.

• freezing point depression.

• osmotic pressure.

Lowering the Vapor Pressure

• Consider a volatile liquid in a closed container.

• After a period of time equilibrium will be established between the liquid and

its vapor.

• The partial pressure exerted by the vapor is the vapor pressure.

• Nonvolatile solutes (with no measurable vapor pressure) reduce the ability of the

surface solvent molecules to escape the liquid.

• Therefore, vapor pressure is lowered.

• The amount of vapor pressure lowering depends on the amount of

solute.

• Raoult’s law quantifies the extent to which a nonvolatile solute lowers the vapor

pressure of the solvent.

• If PA is the vapor pressure with solute, P˚A is the vapor pressure of the

pure solvent, and is the mole fraction of A, then

• An ideal solution is one that obeys Raoult’s law.

• Real solutions show approximately ideal behavior when:

• the solute concentration is low.

• the solute and solvent have similarly sized molecules.

solution of liters

solute of moles Molarity, M

solvent of kilograms

solute of moles Molality, m

oAAA PP


• the solute and solvent have similar types of

intermolecular attractions.

• Raoult’s law breaks down when the solvent–solvent and solute–solute

intermolecular forces are much greater or weaker than solute–solvent

intermolecular forces.

Boiling-Point Elevation

• A nonvolatile solute lowers the vapor pressure of a solution.

• At the normal boiling point of the pure liquid, the solution has a

has a vapor pressure less than 1 atm.

• Therefore, a higher temperature is required to reach a vapor pressure of 1

atm for the solution (∆Tb).

• The molal boiling-point-elevation constant, Kb, expresses how much ∆Tb

changes with molality, m:

∆Tb = Kbm

• The nature of the solute (electrolyte vs. nonelectrolyte) will impact the colligative

molality of the solute.

Freezing-Point Depression

• When a solution freezes, crystals of almost pure solvent are formed first.

• Solute molecules are usually not soluble in the solid phase

of the solvent.

• Therefore, the triple point occurs at a lower temperature because of the

lower vapor pressure for the solution.

• The melting-point (freezing-point) curve is a vertical line from the triple point.

• Therefore, the solution freezes at a lower temperature (∆Tf) than

the pure solvent.

• The decrease in freezing point (∆Tf) is directly proportional

to molality.

• Kf is the molal freezing-point-depression constant.

∆Tf = Kfm

• Values of Kf and Kb for most common solvents can be found in Table 13.4.

Osmosis

• Semipermeable membranes permit passage of some components of a solution.

• Often they permit passage of water but not larger molecules or ions.

• Examples of semipermeable membranes are cell membranes and

cellophane.

• Osmosis is the net movement of a solvent from an area of low solute

concentration to an area of high solute concentration.


• Consider a U-shaped tube with a two liquids separated by a semipermeable

membrane.

• One arm of the tube contains pure solvent.

• The other arm contains a solution.

• There is movement of solvent in both directions across

a semipermeable membrane.

• As solvent moves across the membrane, the fluid levels in the arms

become uneven.

• The vapor pressure of solvent is higher in the arm with

pure solvent.

• Eventually the pressure difference due to the difference in height of liquid in

the arms stops osmosis.

• Osmotic pressure, , is the pressure required to prevent osmosis.

• Osmotic pressure obeys a law similar in form to the ideal-gas law.

• For n moles, V= volume, M= molarity, R= the ideal gas constant, and

an absolute temperature, T, the osmotic pressure is:

V = nRT

• Two solutions are said to be isotonic if they have the same

osmotic pressure.

• Hypotonic solutions have a lower , relative to a

more concentrated solution.

• Hypertonic solutions have a higher , relative to a more

dilute solution.

• We can illustrate this with a biological system: red blood cells.

• Red blood cells are surrounded by semipermeable membranes.

• If red blood cells are placed in a hypertonic solution (relative to

intracellular solution), there is a lower solute concentration in the cell than

the surrounding tissue.

• Osmosis occurs and water passes through the membrane

out of the cell.

• The cell shrivels up.

• This process is called crenation.

• If red blood cells are placed in a hypotonic solution, there is a higher

solute concentration in the cell than outside the cell.

• Osmosis occurs and water moves into the cell.

• The cell bursts (hemolysis).

• To prevent crenation or hemolysis, IV (intravenous) solutions must be

isotonic relative to the intracellular fluids of cells.

MRTRTV

n


• Everyday examples of osmosis include:

• If a cucumber is placed in NaCl solution, it will lose water

to shrivel up and become a pickle.

• A limp carrot placed in water becomes firm because water

enters via osmosis.

• Eating large quantities of salty food causes retention of water

and swelling of tissues (edema).

• Water moves into plants, to a great extent, through osmosis.

• Salt may be added to meat (or sugar added to fruit) as

a preservative.

• Salt prevents bacterial infection: A bacterium placed on the salt will

lose water through osmosis and die.

• Active transport is the movement of nutrients and waste material through a

biological membrane against a concentration gradient.

• Movement is from an area of low concentration to an area

of high concentration.

• Active transport is not spontaneous.

• Energy must be expended by the cell to accomplish this.

Determination of Molar Mass

• Any of the four colligative properties may be used to determine molar mass.

13.6 Colloids (EOCQ’S )

• Colloids or colloidal dispersions are suspensions in which the suspended

particles are larger than molecules but too small to separate out of the suspension

due to gravity.

• Particle size: 5 to 1000 nm.

• There are several types of colloids:

• aerosol: gas + liquid or solid (e.g., fog and smoke),

• foam: liquid + gas (e.g., whipped cream),

• emulsion: liquid + liquid (e.g., milk),

• sol: liquid + solid (e.g., paint),

• solid foam: solid + gas (e.g., marshmallow),

• solid emulsion: solid + liquid (e.g., butter),

• solid sol: solid + solid (e.g., ruby glass).

• The Tyndall effect is the ability of colloidal particles to scatter light.

• The path of a beam of light projected through a colloidal suspension can be

seen through the suspension.


Hydrophilic and Hydrophobic Colloids

• Focus on colloids in water.

• Water-loving colloids are hydrophilic.

• Water-hating colloids are hydrophobic.

• In the human body, large biological molecules such as proteins are kept in

suspension by association with surrounding water molecules.

• These macromolecules fold up so that hydrophobic groups are away from

the water (inside the folded molecule).

• Hydrophilic groups are on the surface of these molecules and interact with

solvent (water) molecules.

• Typical hydrophilic groups are polar

(containing C–O, O–H, N–H bonds) or charged.

• Hydrophobic colloids need to be stabilized in water.

• One way to stabilize hydrophobic colloids is to adsorb ions on

their surface.

• Adsorption: when something sticks to a surface we say that it

is adsorbed.

• If ions are adsorbed onto the surface of a colloid, the colloid appears

hydrophilic and is stabilized in water.

• Consider a small drop of oil in water.

• Add a small amount of sodium stearate.

• Sodium stearate has a long hydrophobic hydrocarbon tail and

a small hydrophilic head.

• The hydrophobic tail can be absorbed into the oil drop, leaving the

hydrophilic head on the surface.

• The hydrophilic heads then interact with the water and the

oil drop is stabilized in water.

• A soap acts in a similar fashion.

• Soaps are molecules with long hydrophobic tails and hydrophilic

heads that remove dirt by stabilizing the colloid in water.

• Most dirt stains on people and clothing are oil based.

• Biological application of this principle:

• The gallbladder excretes a fluid called bile.

• Bile contains substances (bile salts) that form an emulsion with

fats in our small intestine.

• Emulsifying agents help form an emulsion.

• Emulsification of dietary fats and fat-soluble vitamins is important in

their absorption and digestion by the body.


Removal of Colloid Particles

• We often need to separate colloidal particles from the dispersing medium.

• This may be problematic.

• Colloid particles are too small to be separated by physical means (e.g.,

filtration).

• However, colloid particles often may be coagulated (enlarged) until they

can be removed by filtration.

• There are various methods of coagulation.

• Colloid particles move more rapidly when the colloidal dispersion is heated,

increasing the number of collisions. The particles stick to each other when

they collide.

• Adding an electrolyte neutralizes the surface charges on the

colloid particles.

• A biological application of another approach to separating colloidal particles from

the suspending medium is dialysis.

• In dialysis a semipermeable membrane is used to separate ions from

colloidal particles.

• In kidney dialysis, the blood is allowed to pass through a semipermeable

membrane immersed in a washing solution.

• The washing solution is isotonic in ions that must be retained.

• The washing solution does not have the waste products that are found

in the blood.

• Wastes therefore dialyze out of the blood (move from the blood into

the washing solution).

• The "good" ions remain in the blood.

ADDITIONAL EXERCISES: PAGE569 Q’S 81,83,85,87,90,93,96,9

INTEGRATIVE EXERCISES: PAGE 570 Q’S 104, 108, 111

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