Chapter 4.1-4.5 Types of Chemical Reactions and Solution Stoichiometry PHall AP info. NOT integrated...

Chapter 4.1-4.5

Types of Chemical

Reactions and Solution

Stoichiometry

PHall AP info. NOT integrated into this presentation

Chapter Four:

TYPES OF CHEMICAL REACTIONS AND

SOLUTION STOICHIOMETRY

Sugar & Potassium Chlorate video not sure where it goes...

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Chapter 4

Table of Contents

Copyright © Cengage Learning. All rights reserved 2

4.1 Water, the Common Solvent

4.2 The Nature of Aqueous Solutions: Strong and Weak Electrolytes

4.3 The Composition of Solutions

4.4 Types of Chemical Reactions

4.5 Precipitation Reactions

Section 4.1

Water, the Common Solvent

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• One of the most important substances on Earth.

• Can dissolve many different substances.

• A polar molecule because of its unequal charge distribution.

Copyright © Houghton Mifflin Company. All rights reserved. 4–

Dissolution of a Solid in a LiquidClick here to watch video.


SoluteSoluteA solute is the dissolved substance in a A solute is the dissolved substance in a solution.solution.

A solvent is the dissolving medium in a A solvent is the dissolving medium in a solution.solution.

SolventSolvent

Salt in salt water Sugar in soda drinks

Carbon dioxide in soda drinks

Water in salt water Water in soda

““Like Dissolves Like”Like Dissolves Like”

FatsFats BenzeneBenzene

SteroidsSteroids HexaneHexane

WaxesWaxes TolueneToluene

Polar and ionic solutes dissolve best in polar solvents

Nonpolar solutes dissolve best in nonpolar solvents

Inorganic SaltsInorganic Salts WaterWater

SugarsSugars Small alcoholsSmall alcohols

Acetic acidAcetic acid

Section 4.2

The Nature of Aqueous Solutions: Strong and Weak Electrolytes

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• Solute – substance being dissolved.• Solvent – liquid water.• Electrolyte – substance that when dissolved in

water produces a solution that can conduct electricity.

Nature of Aqueous Solutions

Section 4.2

The Nature of Aqueous Solutions: Strong and Weak Electrolytes

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• Strong Electrolytes – conduct current very efficiently (bulb shines brightly).

• Weak Electrolytes – conduct only a small current (bulb glows dimly).

• Nonelectrolytes – no current flows (bulb remains unlit).

Electrolytes


Electrolytesclick here to watch video.



Electrolyte BehaviorClick here to watch visualization.


An electrolyte is:An electrolyte is:

A substance whose aqueous solution conducts an electric current.

A nonelectrolyte is:A nonelectrolyte is:

A substance whose aqueous solution does not conduct an electric current.

Definition of Electrolytes and Definition of Electrolytes and NonelectrolytesNonelectrolytes

The ammeter measures the flow of electrons (current) through the circuit.

If the ammeter measures a current, and the bulb glows, then the solution conducts.

If the ammeter fails to measure a current, and the bulb does not glow, the solution is non-conducting.

Electrolytes vs. NonelectrolytesElectrolytes vs. Nonelectrolytes

1.Pure water2.Tap water3.Sugar solution4.Sodium chloride solution5.Hydrochloric acid solution6.Lactic acid solution7.Ethyl alcohol solution8.Pure sodium chloride

1.Pure water2.Tap water3.Sugar solution4.Sodium chloride solution5.Hydrochloric acid solution6.Lactic acid solution7.Ethyl alcohol solution8.Pure sodium chloride

Try to classify the following substances Try to classify the following substances as electrolytes or nonelectrolytes…as electrolytes or nonelectrolytes…

ELECTROLYTES: NONELECTROLYTES:

Tap water (weak)

NaCl solution

HCl solution

Lactate solution (weak)

Pure water

Sugar solution

Ethanol solution

Pure NaCl

Answers to ElectrolytesAnswers to Electrolytes

Ionic Compounds “Dissociate”Ionic Compounds “Dissociate”

NaCl(s)

AgNO3(s)

MgCl2(s)

Na2SO4(s)

AlCl3(s)

Na+(aq) + Cl-(aq)

Ag+(aq) + NO3-(aq)

Mg2+(aq) + 2 Cl-(aq)

2 Na+(aq) + SO42-(aq)

Al3+(aq) + 3 Cl-(aq)

The reason for this is the polar nature of the water molecule…

Positive ions associate with the negative end of the water dipole (oxygen).

Negative ions associate with the positive end of the water dipole (hydrogen).

Ions tend to stay in solution where they canconduct a current rather than re-forming a solid.

Covalent acids form ions in solution, with the help of the water molecules.

For instance, hydrogen chloride molecules,which are polar, give up their hydrogens towater, forming chloride ions (Cl-) and hydronium ions (H3O+).

Some covalent compounds IONIZE in solution

Other examples of strong acids include:

Sulfuric acid, H2SO4

Nitric acid, HNO3

Hydriodic acid, HI Perchloric acid, HClO4

Strong acids such as HCl are completelyionized in solution.

Many of these weaker acids are “organic” acids that contain a “carboxyl” group.

The carboxyl group does not easily give up itshydrogen.

Weak acids such as lactic acid usually ionize less than 5% of the time.

Other organic acids and their sources include:

o Citric acid – citrus fruito Malic acid – appleso Butyric acid – rancid buttero Amino acids – proteino Nucleic acids – DNA and RNAo Ascorbic acid – Vitamin C

This is an enormous group of compounds; these are only a few examples.

Because of the carboxyl group, organic acids are sometimes called “carboxylic acids”.

SSoolluuttiioonnss

solutionsxt

Solution ConcentrationSolution Concentration

Calculations of Solution Concentration:Calculations of Solution Concentration:MolarityMolarity

MolarityMolarity is the ratio of moles of solute to liters of solution

Section 4.3

The Composition of Solutions

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• We must know: The nature of the reaction. The amounts of chemicals present in

the solutions.

Chemical Reactions of Solutions

Section 4.3


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• Molarity (M) = moles of solute per volume of solution in liters:

Molarity

Section 4.3


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Exercise

A 500.0-g sample of potassium phosphate is dissolved in enough water to make 1.50 L of solution. What is the molarity of the solution? (calculate #moles per L to get molarity)

1.57 M

Preparation of Molar SolutionsPreparation of Molar Solutions

Problem: How many grams of sodium chloride are needed to prepare 1.50 liters of 0.500 M NaCl solution?

Step #1: Ask “How Much?” (What volume to prepare?)

1.500 L

Step #2: Ask “How Strong?” (What molarity?)

0.500 mol

1 L

Step #3: Ask “What does it weigh?” (Molar mass is?)

58.44 g

1 mol= 43.8 g

Section 4.3


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• For a 0.25 M CaCl2 solution:

CaCl2 → Ca2+ + 2Cl–

Ca2+: 1 × 0.25 M = 0.25 M Ca2+

Cl–: 2 × 0.25 M = 0.50 M Cl–.

Concentration of Ions

Note that there were 2 times as many ions of Cl- formed, so the molarity was doubled due to

having twice as many moles like in the titration lab with citric acid which had 3 moles

hydrogen ions produced for each mole of the base.

Section 4.3


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Concept Check

Which of the following solutions containsthe greatest number of ions?

a) 400.0 mL of 0.10 M NaCl.

b) 300.0 mL of 0.10 M CaCl2.

– 200.0 mL of 0.10 M FeCl3.

– 800.0 mL of 0.10 M sucrose.

Section 4.3


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• Where are we going? To find the solution that contains the greatest

number of moles of ions.

• How do we get there? Draw molecular level pictures showing each

solution. Think about relative numbers of ions. How many moles of each ion are in each

solution?

Let’s Think About It

Section 4.3


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• The solution with the greatest number of ions is not necessarily the one in which: the volume of the solution is the

largest. the formula unit has the greatest

number of ions.

Notice

Section 4.3


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• The process of adding water to a concentrated or stock solution to achieve the molarity desired for a particular solution.

• Dilution with water does not alter the numbers of moles of solute present.

• Moles of solute before dilution = moles of solute after dilution

M1V1 = M2V2

Dilution

Section 4.3


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Concept Check

A 0.50 M solution of sodium chloride in an open beaker sits on a lab bench. Which of the following would decrease the concentration of the salt solution?

a) Add water to the solution.

b) Pour some of the solution down the sink drain.

c) Add more sodium chloride to the solution.

d) Let the solution sit out in the open air for a couple of days.

e) At least two of the above would decrease the concentration of the salt solution.

Serial DilutionSerial DilutionProblem: What volume of stock (11.6 M) hydrochloric acid is needed to prepare 250. mL of 3.0 M HCl solution?

MstockVstock = MdiluteVdilute

(11.6 M)(x Liters) = (3.0 M)(0.250 Liters)

x Liters = (3.0 M)(0.250 Liters) 11.6 M

= 0.065 L


DilutionClick here to watch video


Section 4.3


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Exercise

What is the minimum volume of a 2.00 M NaOH solution needed to make 150.0 mL of a 0.800 M NaOH solution?

60.0 mL

Section 4.4

Types of Chemical Reactions

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• Precipitation Reactions• Acid–Base Reactions• Oxidation–Reduction Reactions

Precipitation ReactionsPrecipitation Reactions

Graphic: Wikimedia Commons User Tubifex

Section 4.5

Precipitation Reactions

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• A double displacement reaction in which a solid forms and separates from the solution. When ionic compounds dissolve in

water, the resulting solution contains the separated ions.

Precipitate – the solid that forms.

Precipitation Reaction

Section 4.5


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• Ba2+(aq) + CrO42–(aq) → BaCrO4(s)

The Reaction of K2CrO4(aq) and Ba(NO3)2(aq)


Precipitation of Silver Chloride

Click here to watch visualization


Double Replacement Reactions

The ions of two compounds exchange places in an aqueous solution to form two new compounds.

AX + BY →AY + BX

One of the compounds formed is usually a precipitate (an insoluble solid), an insoluble gas that bubbles out of solution, or a molecular compound, usually water.

Double replacement forming a precipitate…

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

Pb2+(aq) + 2 NO3-(aq) + 2 K+(aq) +2 I-(aq) PbI2(s) + 2K+(aq) + 2 NO3

-(aq)

PbPb2+2+(aq) + 2 I(aq) + 2 I--(aq) (aq) PbIPbI22(s)(s)

Double replacement (ionic) equation

Complete ionic equation shows compounds as aqueous ions

Net ionic equation eliminates the spectator ions

Lead(II) nitrate + potassium iodide →lead(II) iodide + potassium nitrate

Solubility Rules – Mostly SolubleSolubility Rules – Mostly Soluble

IonIon Solubility Solubility ExceptionsExceptions

NONO33-- SolubleSoluble NoneNone

ClOClO44-- SolubleSoluble NoneNone

NaNa++ SolubleSoluble NoneNone

KK++ SolubleSoluble NoneNone

NHNH44++ SolubleSoluble NoneNone

ClCl--, I, I-- SolubleSoluble PbPb2+2+, Ag, Ag++, Hg, Hg222+2+

SOSO442-2- SolubleSoluble CaCa2+2+, Ba, Ba2+2+, Sr, Sr2+2+, Pb, Pb2+2+, Ag, Ag++, Hg, Hg2+2+

Solubility Rules – Mostly InsolubleSolubility Rules – Mostly Insoluble

IonIon Solubility Solubility ExceptionsExceptions

COCO332-2- InsolubleInsoluble Group IA and NHGroup IA and NH44

++

POPO443-3- InsolubleInsoluble Group IA and NHGroup IA and NH44

++

OHOH-- InsolubleInsoluble Group IA and CaGroup IA and Ca2+2+, Ba, Ba2+2+, Sr, Sr2+2+

SS2-2- InsolubleInsoluble Groups IA, IIA, and NHGroups IA, IIA, and NH44++


Table 4.1 Simple Rules for the Solubility of Salts in Water


Solubility RulesClick here to watch video.


Solubility Solubility Chart:Chart:

Common Common saltssalts

at 25at 25CC

S = SolubleI = InsolubleP = Partially SolubleX = Other

Section 4.5


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• Soluble – solid dissolves in solution; (aq) is used in reaction.

• Insoluble – solid does not dissolve in solution; (s) is used in reaction.

• Insoluble and slightly soluble are often used interchangeably.

Precipitates

Section 4.5


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1. Most nitrate (NO3−) salts are soluble.

2. Most alkali metal (group 1A) salts and NH4+ are soluble.

3. Most Cl−, Br−, and I− salts are soluble (except Ag+, Pb2+, Hg2

2+).

4. Most sulfate salts are soluble (except BaSO4, PbSO4, Hg2SO4, CaSO4).

5. Most OH− are only slightly soluble (NaOH, KOH are soluble, Ba(OH)2, Ca(OH)2 are marginally soluble).

6. Most S2−, CO32−, CrO4

2−, PO43−- salts are only slightly

soluble, except for those containing the cations in Rule 2.

Simple Rules for Solubility

Solubility TrendsSolubility Trends The solubility of MOST solids increases The solubility of MOST solids increases

with temperature.with temperature. The rate at which solids dissolve increases The rate at which solids dissolve increases

with increasing surface area of the solid.with increasing surface area of the solid. The solubility of gases decreases with The solubility of gases decreases with

increases in temperature.increases in temperature. The solubility of gases increases with the The solubility of gases increases with the

pressure above the solution.pressure above the solution.

Therefore…Therefore…Solids tend to dissolve best when:

o Heatedo Stirredo Ground into small particles

Gases tend to dissolve best when:

o The solution is coldo Pressure is high

Solubility ChartSolubility Chart

Saturation of SolutionsSaturation of Solutions A solution that contains the maximum amount A solution that contains the maximum amount

of solute that may be dissolved under existing of solute that may be dissolved under existing conditions is conditions is saturated.saturated.

A solution that contains less solute than a A solution that contains less solute than a saturated solution under existing conditions is saturated solution under existing conditions is unsaturated.unsaturated.

A solution that contains more dissolved solute A solution that contains more dissolved solute than a saturated solution under the same than a saturated solution under the same conditions is conditions is supersaturated.supersaturated.

Section 4.5


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Concept Check

Which of the following ions form compounds with Pb2+ that are generally soluble in water?

a) S2–

b) Cl–

c) NO3–

d) SO42–

e) Na+

57

Section 4.5


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END OF SLIDES FOR SECTION 4.1-4.5

• END OF SLIDES FOR SECTION 4.1-4.5

Chapter 4.1-4.5 Types of Chemical Reactions and Solution Stoichiometry PHall AP info. NOT integrated...

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