chapter 4 jan12

7/29/2019 Chapter 4 Jan12

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CHAPTER 4

AQUEOUS REACTIONS

& SOLUTION

STOICHIOMETRY


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CONTENTS

4.1 General Properties of Aqueous Solutions

4.2 Precipitation Reactions

4.3 Acid-Base Reactions

4.4 Oxidation-Reduction Reactions

4.5 Concentrations of Solutions

4.6 Solution Stoichiometry and Chemical

Analysis


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Learning Outcomes

Able to identify types ofmetathesis reaction

Able to determine the precipitate forms in a

reactionAble to calculate oxidation numberof

element and identify the oxidizing agent in a

redox reactionAble to calculate unknown concentration in a

titration of acid-base solutions (reaction).


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4.1 General Properties ofAqueous Solutions

A solution is a homogeneous mixture.

Contains: a solvent (greater quantity) andsolute(s).

Solvent - substance in the mixture which actsas the dissolving medium.

Whatever else is dissolved in the solution iscalled the solute(s).


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Cont: 4.1 General Properties ofAqueous Solutions

E.g.

NaCl (s) + H2O (l) NaCl (aq)

solute solvent aqueous solution

Solution in which water is the dissolving medium

are called aqueous solutions.


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Cont: 4.1 General Properties ofAqueous Solutions

All aqueous solutions can be classified in terms ofits electrical conductivity.

Solution conducts electricity:

forms ions in solution

is an electrolyte. e.g. NaCl (an ionic compound)

# conducts electricity well: strong electrolyte

(NaCl)# conducts electricity poorly: weak electrolytes

(CH3COOH)


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Cont: 4.1 General Properties oAqueous Solutions

Solution does not conduct electricity

does not form ions in solution

is a non electrolyte e.g. sugar (amolecular compound)


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4.1.1 Ionic Compounds inWater

Ionic compounds dissolve in water,dissociate into anions and cations.

The solid no longer exists as a well ordered

arrangement, each ion is surrounded bywater molecules.

The ions are dispersed uniformly throughoutthe solution.

The relative concentrations of the ionsdepend on the chemical formula of thecompound. (Sec: 4.5.2)


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4.1.2 Molecular Compounds inWater

Molecular compounds (e.g. CH3OH, sugaretc.) dissolve and exist as dispersed

molecules throughout the solution. The structural integrity of the compound

is maintained.

Does not form ions

no ions to conductelectricity.

they are non-electrolytes


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Cont: 4.1.2 MolecularCompounds in Water

Some important exceptions:

NH3

dissolves in water to form NH4

+

and

OH-

HCl (g) in water ionizes to form H+ and Cl-


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4.1.3 Strong and WeakElectrolytes

Electrolyte: any substance whose aqueoussolution contains ions that conduct electricity.

Strong electrolytes - ionize 100 %in a solvent.

E.g. (i) Most ionic compounds (salt), NaOH(ii) Strong acids and bases: HCl, HBr, HClO4

A single arrow ( ): ionization of strongelectrolytes.

A single arrow indicates that ions have notendency to recombine.


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Cont: 4.1.3 Strong and WeakElectrolytes

Weak electrolytes - Incompletely ionization(small amount of ions).E.g. (i) Weak acid : HF, H2S, HC2H3O etc.

(ii) Weak base : NH3, amine A 1.0 M solution of acetic acid HC2H3O2:

Small fraction (about 1 %) is present as H+ and

C2H3O2-

ions. Most is present as HC2H3O2 molecules.

HC2H3O2(aq) H+(aq)+ C2H3O2

-(aq)


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Cont: 4.1.3 Strong and WeakElectrolytes

The double arrow () is significant:

HC2H3O2 molecules are ionizing to form H+

and C2H3O2- and at the same time H+ andC2H3O2

- ions are recombining.

Note: A double arrow ionization of weakelectrolytes.


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4.2 Precipitation Reactions

Precipitation: the reaction forming

insoluble product (called precipitate).

caused by the attraction between the oppositely

charged ions in the solid being too great for thewater molecules to pull them apart.


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4.2.1 Solubility Guidelines forIonic Compounds

Know the solubilities of differentcompounds

Solubility guidelines are used to predictwhether a precipitate will form whensolutions are mixed.


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H


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4.2.2 Metathesis Reaction

Metathesis reactions involve swapping ions insolution:

AX + BY AY + BX

AgNO3(aq) + KCl(aq) AgCl(s)+ KNO3(aq)

Positive ions (cations) and negative ions (anions) -exchange partners.

Sometimes known as double-displacement reaction.


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Cont: 4.2.2 MetathesisReaction

Displacement reaction

Single displacement Double displacement

(Metathesis)

X + YZ XZ + Y WX + YZ WZ + YX

Eg: Ox-Red Rxn(Activity Series)(Sec. 4.4)

Eg: Ppt RxnAcid-Base Rxn


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Cont: 4.2.2 MetathesisReaction

E.g: Acid-base neutralization reaction

HCl(aq) + NaOH(aq) NaCl(aq) + H2

O(l)

H+ + OH- H2O ions are removed from thesolution

For a metathesis reaction to occur, ions mustbe removed from the solution (Sec. 4.2.3)


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Cont: 4.2.2 MetathesisReaction: Driving Forces

Three chemical processes can lead to theremoval of ions from solution (driving forcefor metathesis to occur):

(i) The formation of an insoluble product(called a precipitate).(4.2)

(ii) The formation of either a weak

electrolyte or a non electrolyte.(4.3.4)(iii) The formation of a gas that escapesfrom solution.(4.3.3)


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Cont: 4.2.2 MetathesisReaction (to predict products)

Metathesis reaction is used to predict whethertwo salts or two reactants can react to producethe products(s).

Some Rules:1. If one reactant is a salt containing the carbonate

ionor the hydrogen carbonate ionand the other

is a strong acid, the products will be a saltcontaining the anion of the acid, waterand CO2.

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


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Cont: 4.2.2 MetathesisReaction (To predict products)

2. If the reactants are an acidandabase, carryout a neutralization reaction.

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


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Cont: 4.2.2 MetathesisReaction (To predict products)

3. Insoluble metal oxidesreact with strong acidsto form a salt (containing the anion of the acid)

and water.

MgO(s) + 2HCl(aq) MgCl2(aq) + H2O(l)

4. If both reactants are weak acids/ weak bases,no reaction occurs.


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4.2.3 Ionic Equations

Chemical equations Molecular equation

Ionic equation: complete ionic equation: net ionic equation

indicate whether dissolved substances arepresent in solution predominantly as ions oras molecules.


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Cont: 4.2.3 Ionic Equations

Molecular equation - list all species (reactantsand products) in their molecular forms:

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

Complete ionic equation - lists allions in thereaction:

H+(aq)+ Cl-(aq)+ Na+(aq)+ OH-(aq) H2O (l)+ Na +(aq)+ Cl-(aq)


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Spectator ions : ions that appear inidentical form on both sides of the arrow

(they play no role in the reaction).H+(aq)+ Cl-(aq)+ Na+(aq)+ OH-(aq)

H2O (l)+ Na +(aq)+ Cl-(aq)

Na+(aq)and Cl-(aq)appear on both sides

- spectator ions.


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Omitting the spectator ions: Na+(aq)and Cl(aq)

Net ionic equation:

H+(aq)+ OH-(aq) H2O (l)

Include only the ions and moleculesdirectly involved in the chemical reaction.

Charge is conserved in the reaction (sumof the charges are the same on both sidesof the arrow).


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To write an ionic equation:

Only strong electrolytes are written in ionicform.

Non electrolytes, insoluble are written in

molecular form.

Note: take note on the difference between ionicequation and net ionic equation.


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Guidelines for writing a balanced ionicequation:

Write a balanced molecular equationfor the reaction

Rewrite the equation to show the ions

formed from dissociation (for strongelectrolytes).

Identify and cancel spectator ions.


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Example 1

Write the net ionic equation forzinc metal with

hydrochloric acid to form hydrogen gas and

aqueous solution of zinc chloride.


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Example 1 (Answer)

Molecular Equation (complete):

Zn (s)+ 2HCl (aq) H2(g)+ ZnCl2(aq)

Ionic Equation (complete):

Zn (s)+ 2H+(aq)+ 2Cl-(aq)H2(g)+ Zn

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


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Example 1 (Answer)

Spectator ion is Cl-Zn (s)+ 2H+(aq)+ 2Cl-(aq)

H2(g)+ Zn2+

(aq)+ 2Cl-

(aq)

Net Ionic Equation:Zn (s)+ 2H+(aq) H2(g)+ Zn

2+(aq)

Reaction can occur when the net ionicequation can be written.


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Example 2

Predict whether any reaction occurs when

potassium chloride is mixed with sodium nitrate.


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Example 2 (Answer)

Molecular Equation (complete):

KCl(aq) + NaNO3(aq) KNO3(aq) + NaCl(aq)

Ionic Equation (complete):

K+(aq)+ Cl-(aq)+ Na+(aq)+ NO3-(aq)

K+(aq)+ NO3-(aq)+ Na+(aq)+ Cl-(aq) No net ionic equation.

No reaction occurs.


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Example 3

Write balanced molecular, ionic and net ionicequations for the precipitation reactions (if any)for the following compounds:

(a) Mg(NO3)2 and NaOH

(b) KCl and Na2SO4


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Example 3 (Answer)

(a) Write balanced molecular equation:

Mg(NO3)2(aq)+ 2NaOH (aq)

Mg(OH)2(s)+ 2NaNO3(aq)

Note: From guidelines table Mg(OH)2 willprecipitate.


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Example 3 (Answer)

Ionic equation:

Mg2+(aq)+ 2NO3-(aq)+ 2Na+(aq)+ 2OH-(aq)

Mg(OH)2(s)+ 2Na+

+ 2NO3-

(aq)

Net ionic equation:

Mg2+(aq)+ 2OH-(aq) Mg (OH)2

(s)

Note: Na+ and NO3- are spectator ions.


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Example 3 (Answer)

(b) Write balanced molecular equation:

2KCl (aq) + Na2SO

4(aq)

K2SO4(aq) + 2NaCl (aq)

Note: Both K2SO4 and NaCl are soluble and

dissociate in solution.


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Example 3 (Answer)

Ionic equation:

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

2K+(aq) + SO42-

(aq) + 2Na+(aq) + 2Cl-(aq)

No reaction; the solutes merely mix in the solution.


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4.3 Acid-Base Reactions

Acids

Definitions (based on Arrhenius)

Substance that is able to ionize to form ahydrogen ion (H+) in solution, e.g. HCl, HNO3etc.

Increase the concentration of H+(aq) ions.

H+ is a proton. Acids are called proton donors.


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Cont: 4.3 Acid-Base Reactions

Acids that ionize to form one H+ are calledmonoprotic acids. E.g: HCl and HNO3

Acids that ionize to form two H+ ions are called

diprotic acids. E.g:H2SO4 - yields two H+ ionsper molecule:

H2SO4(aq) H+(aq)+ HSO4-(aq)

HSO4- (aq) H+(aq)+ SO42-(aq) Aqueous solution of sulfuric acid contains a

mixture of H+(aq), HSO4-(aq)and SO4

2- (aq).


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Acids that ionize to form more than twoH+ ions are called polyprotic acids.

Bases: Soluble ionic compounds containing the

hydroxide ions, OH-, eg. NaOH, Ca(OH)2,

KOH etc. Increase the concentration of OH-(aq)

when added to water.


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React with the H+ ions to form water(protonacceptor)H+(aq)+ OH-(aq) H2O (l) (neutralization rxn)

Note: NH3 (ammonia) Compounds do not contain OH- ion BUT when

added to water, accepts H+

ion from the watermolecules increase the concentration of OH-ions in water.E.g. NH3(aq)+ H2O(l) NH4+ (aq)+ OH-(aq)


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4.3.1 Strong and Weak Acidsand Bases

Strong Acids and Bases

Strong electrolytes

Completely ionized in solution

Definition :

Acids - proton donor

Bases - proton acceptor


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H

nnnnnn


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Cont: 4.3.1 Strong and WeakAcids and Bases

How to classify non-electrolyte, weak electrolyte orstrong electrolyte:

Compounds Strong Weak Non

Ionic All None None

Molecular Strongacids

Weak acidsand bases

All othercompounds

Electrolyte


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Example 4

Classify each of the following aqueous solution asa non-electrolyte, weak electrolyte or strongelectrolyte.

(a) HBrO (f) Sucrose

(b) HF (g) O2

(c) HNO3(d) KOH

(e) CoSO4


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Example 4 (Answer)

(a) HBrO - weak acid - weak electrolyte

(b) HF - weak acid - weak electrolyte

(c) HNO3 - strong acid - strong electrolyte

(d) KOH - strong base - strong electrolyte

(e) CoSO4 - salt - strong electrolyte

(f) Sucrose - molecular - non electrolyte

(g) O2 - molecular - non electrolyte


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Example 5

Classify the followings as a strong electrolyte,weak electrolyte or non electrolyte.

HBr; H2S; NH3; Ba(OH)2; KCl; C6H6; I2


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Example 5 (Answer)

If a substance is a salt (metallic cation and anon-metallic anion) strong electrolyte.

Ba (OH)2 strong electrolyte HBr is a strong acid strong electrolyte H2S is a weak acid - weak electrolyte NH3 is a weak base - weak electrolyte

C6H6 is a hydrocarbon- non electrolyte I2 is a homonuclear diatomic element - not salts

and not listed as acids or bases nonelectrolyte


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4.3.2 Neutralization Reactionsand Salts

Neutralization Reactions and Salts

A neutralization reaction occurs when an acid

and a base react to form a salt.E.g. HCl (aq)+ NaOH (aq) H2O (l)+ NaCl (aq)

A saltis an ionic compound whose cation comes

from a base and anionfrom an acid.

Net ionic equation: H+(aq)+ OH-(aq) H2O (l)


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4.3.3 Acid-Base Reaction withGas Formation

One of the driving force for metathesisreaction to occur.

One of the products is a gas.

Carbonates and hydrogen carbonates(bicarbonates) react with acids to form CO2 gas.

Example :

Sodium hydrogen carbonate (NaHCO3) reactswith HCl


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Cont: 4.3.3 Acid-Base Reactionwith Gas Formation

Molecular equation:HCl (aq)+ NaHCO3(aq)

NaCl (aq)+ H2O (l)+ CO2(g)

Ionic equation:H+(aq)+ Cl-(aq)+ Na+(aq)+ HCO3 -(aq)

Na+(aq)+ Cl-(aq)+ H2O (l)+ CO2(g)

Net ionic equation:H+(aq)+ HCO3-(aq) H2O (l)+ CO2(g)


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H

nnnnnn


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H

nnnnnn

4.3.4 Acid-Base Reactions with


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4.3.4 Acid Base Reactions withWeak Electrolyte or Nonelectrolyte

Formation

Common weak electrolyte : H2O

Example : Neutralization reaction:

H++ OH- H2O

Example: Mg(OH)2 - white suspension dissolveswhen reacts with HCl (aq)

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

Cont: 4.3.4 Acid-Base Reactions


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Con t: 4.3.4 Acid Base Reactionswith Weak Electrolyte orNonelectrolyte Formation

Ionic equation:

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

Mg2+(aq)+ 2Cl-(aq)+ 2H2O (l)

Net ionic equation:

Mg(OH)2(s)+ 2H+(aq) Mg2+(aq)+ 2H2O (l)

Note: H2O is the driving force


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Example 6

Write balanced net ionic equation for thereaction (if any) that may occur and indicate thedriving forces when the following pair is mixed.

Cr(C2H3O2)2(aq)and HNO3(aq)

Answer

Molecular equation:Cr(C2H3O2)2(aq)+ 2HNO3(aq)

Cr(NO3)2(aq)+ 2HC2H3O2(aq)


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Example 6 (Answer)

Ionic equation:Cr2+ (aq)+ 2C2H3O2

-(aq)+ 2H+(aq)+ 2NO3-(aq)

Cr2+ (aq)+ 2HC2H3O2(aq)+ 2NO3- (aq)

Net ionic equation:2C2H3O2-(aq)+ 2H+(aq)2HC2H3O2(aq)

C2H3O2-(aq)+ H+(aq)HC2H3O2(aq)

Note: HC2H3O2 is a driving force


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4.4 Oxidation-ReductionReactions (Redox)

Corrosion - conversion of a metal into a metalcompound.

When a metal undergoes corrosion, it loses

electrons and form cations.

E.g. Ca (s) + 2H+ (aq)Ca2+ (aq) + H2(g)


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Cont: 4.4 Oxidation-ReductionReactions

When atom, ion or molecule becomes morepositively charged (has lost electrons) - it hasbeen oxidized.

Loss of electrons by a substance is calledoxidation.

OIL RIG (Oxidation Is Losing, ReductionIs Gaining- in

term of electrons)


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Example: Metal react with O2 in air to formmetal oxides. The metal loses electrons to

oxygen, forming metal ion and oxide ion.2Ca (s)+ O2(g)2CaO (s)

Oxygen is transformed from neutral O2 to theO 2- ion.


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When an atom, ion or molecule become morenegatively charged, it is reduced.

The gain of electrons by a substance is called

reduction. One reactant loses electrons, another

reactant must gain electrons. This is calledoxidation-reduction or redox reactions.

e-Substance

oxidised

(loses electron)

Substance

reduced

(gains electron)


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4.4.1 Oxidation Number

Oxidation number of an atom in a substance is theactual charge of the atom if it is a monatomicion; otherwise, it is the hypothetical charge

assigned to the atom using a set of rules.

1. Atom in elemental form, the oxidation number iszero. E.g. S, Ar.

2. Monatomic ion, the oxidation number equals thecharge on the ion. E.g. K+ ox. no. = +1, O2- ox. no.= -2.


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Cont: 4.4.1 Oxidation Number

3. Group 1A elements always have ox. no. = +1.Similarly, Group 2A = +2, Group 3A = +3.

4. The ox. no. of oxygen is usually -2 in both ionicand molecular compounds (except peroxides O22-with ox. no. =-1).

5. The ox. no. of hydrogen is +1 when bonded to

nonmetals and -1 when bonded to metals.

6. The ox. no. of fluoride ion is 1 for allcompounds.


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Cont: 4.4.1 Oxidation Number

7. The ox. no. of halides (Group 7A other thanfluoride) is 1 in binary compounds BUT positivevalues in oxyanions.

8. The sum of ox. no. of all atoms in a neutralcompound is zero.

9. The sum of ox. no. in a polyatomic ion equals the

charge of the ion.10. Oxidation process will lead to increase in ox. no.,

while reduction process decrease in ox. no.


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4.4.2 Oxidation of Metals byAcids and Salts

Acids

Metal react with acids to form salts andhydrogen gas.

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

Complete ionic equation:Mg (s)+ 2H+(aq)+ 2Cl-(aq)

Mg2+ (aq)+ 2Cl- ( aq)+ H2(g)


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Cont: 4.4.2 Oxidation of Metalsby Acids and Salts

Net ionic equation:

Mg (s)+ 2H+(aq) Mg2+ (aq)+ H2

(g)

Metal is oxidized by acid.

H+ ions are reduced to H2.

id i f l


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Salts

Metal can be oxidized by aqueous solution ofvarious salts.

A + BX AX + B

Displacement reaction occur if A is more easilyoxidized than B.

Example:

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

C 2 O id i f l


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Net ionic equation:

Fe (s) + Ni2+ (aq) Fe2+ (aq) + Ni (s)

Oxidation of iron is accompanied by thereduction of Ni2+ to Ni.


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4.4.3 The Activity Series

Activity series is a list of metals arrangedin order of decreasing ease of oxidation.

(i) Metals at the top of the table are mosteasily oxidized - react to form compounds.

(ii) Alkali metals and alkaline earth metals

are at the top - active metals.(iii) The transition elements from group 1B

and 8B are at the bottom - noble metals.


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Cont: 4.4.3 The Activity Series

How to use The Activity Series ??

(i) Any metal on the list can be oxidized by theions of elements below it.

E.g. Cu (s)+ 2Ag+(aq) Cu 2+(aq)+ H2(g)

(ii) Metals above hydrogen are able to reactwith acids to form H2.


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H

nnnnnn


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Example 7

Use the activity series to predict which of thefollowing reactions will occur.

(a) Hg (l) + MnSO4(aq) HgSO4(s) + Mn (s)

(b) 2Ag (s) + H2SO4(aq) Ag2SO4(aq) + H2(g)

(c) Ca (s) + 2H2O (l) Ca(OH)2(aq) + H2(g)


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Example 7 (Answer)

M + HX MX + H

M must be higher in the activity series than H.

(a) Hg (l) + MnSO4(aq) no reaction

Hg lies below Mn in the activity series, thus the

reaction does not occur.


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Example 7 (Answer)

(b) 2Ag (s) + H2SO4(aq) no reaction

Ag lies below hydrogen, thus the reaction does not

occur.

(c) Ca (s) + 2H2O (l) Ca(OH)2(aq) + H2(g)

Ca lies above hydrogen, thus the reaction occurs.


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4.5 Concentrations of Solutions

Concentration - to designate the amount ofsolute dissolved in a given quantity of solvent or

solution.

The greater the amount of solute dissolved in acertain amount of solvent, the more

concentrated the solution.


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4.5.1 Molarity

Quantitative measurement of concentrationrequires accurate determination of the amounts

of solvent and solute present in a solution.

Molarity (M) is used to measure the amount ofsolute in a solution.


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Cont: 4.5.1 Molarity

M = number of moles of solute, n

volume of the solution in liters, V

Molarity can be used as a conversion factor to

change between volume of solution and numberof moles of substance.


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Example 8

What is the molarity of an ethanol (C2H6O)

solution containing 10.0 g of ethanol in waterwith a total volume of 100 mL.


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Example 8 (Answer)

MW of C2H6O = 46.07 g

10.0 g C2H6O 1 mol C2H6O = 0.217 mol C2H6O

46.07 g C2H6O

M = 0.217 mol C2H6O = 2.17 M

0.100 L


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Example 9

How many grams of HCl are contained in 500 mL

of a 0.250 M HCl solution?

1 mol HCl = 36.45 g HCl


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Example 9 (Answer)

HCl is the solute, M = moles solute

volume in liters

mol solute = M volume in liters

= 0.250 M x 0.500 L

= 0.125 mol

mass (g) = 0.125 mol molar mass HCl

= 0.125 mol 36.45 g/mol

= 4.56 g HCl

4 5 2 Expressing the


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4.5.2 Expressing theConcentration of Electrolyte

When an ionic compound dissolves, the relativeconcentrations of the ions depend on thechemical formula of the compound.

Examples:

(i) 1.0 M solution of sodium sulfate, Na2SO4 :

Na2SO4 2Na+

+ SO42-

2.0 M in Na+ ions, 1.0 M in SO42- ions

Cont: 4 5 2 Expressing the


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Con t: 4.5.2 Expressing theConcentration of Electrolyte

(ii) 0.025 M aqueous solution of calcium nitrate, Ca(NO3)2:

Ca (NO3)2 Ca2+ + 2NO3-

0.025 M in Ca2+ ions

2 0.025 M in NO3-

ions


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4.5.3 Dilution

Solutions of known concentration may bediluted with the solvent to produce a morediluted (less concentrated) solution.

The number of moles of solute - unchange. Moles of solute before dilution = moles of

solute after dilution.

moles solute = molarity volume

Minitial Vinitial = Mfinal Vfinal


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Example 10

What is the molarity of a solution of NaOH formed

by diluting 125 mL of a 3.0 M NaOH solution to 500

mL?

Solution

ffii VMVM

MmL

mLM

V

VMM

f

iif 75.0

500

1250.3

4 6 Solution Stoichiometry and


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4.6 Solution Stoichiometry andChemical Analysis

4.6.1 Titration

A titration is an experiment in which theunknown molarity of a substance ismeasured by using the known molarity ofanother substance.


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Cont: 4.6.1 Titration

Example: To determine the unknown concentrationof HCl solution .

Standard solution: NaOH solution, 0.100 M

Prepare 0.100 M NaOH

(i) Take a specific volume of HCl solution (example: 20ml)

(ii) Slowly add the standard 0.100 M NaOH until theneutralization reaction between HCl and NaOH iscompleted.


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H

nnnnnn


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The point at which stoichiometrically equivalentquantities are brought together is known as theequivalence point of the titration.

In acid-base titration, dyes known as acid-baseindicators are used to determine the end point.

e.g. Phenolphthalein is colorless in acidic

solution, red in basic solution.The color change from colorless to red endpoint of the titration.


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The equivalence point of the titration is thepoint where the stoichiometrically correctnumber of the moles of each reactant is

present. The end point of the titration is the point where

the indicator changes.

Note: The equivalence point and end point are notthe same, but coincide closely.


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Example 11

What is the molarity of a solution of H2SO4 if 20.00

mL of a 0.100 M NaOH solution is required to reactcompletely with 25.00 mL of the H2SO4 solution.


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Example 11 (Answer)

The chemical reaction describing the reaction.

H2SO4(aq)+ 2NaOH(aq) Na2SO4(aq)+ 2H20(l)

Note: H2SO4 has two ionizable H+ ions that

react with OH- ions from NaOH.


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Example 11 (Answer)

1

2

4242

SOHSOH

NaOHNaOH

VM

VM

242

42

SOH

NaOHNaOHSOH

V

VMM

MmL

mLM04.0

200.25

00.20100.0


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A 10g solid sample containing Zn(OH)2 isadded to 0.400L of 0.550M solution of HBr.

The solution that remains is still acidic. It is

then titrated with 0.5M NaOH solution which

requires 165 mL to reach the equivalence

point. What is the actual mass of Zn(OH)2 in

the sample? What is its mass percentage?

Example 12


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Example 12 (Answer)

Acid + Base water (neutralization rxn)H+ + OH- salt + H20

From Equation:Mol H+ (from HBr) = mol OH- (from Zn(OH)2 +

mol OH- (from NaOH)

H+ = 0.550M x 4.0L = 0.220 moles

OH- (from NaOH) = 0.5M x (165/1000)L = 0.0825 moles

OH- (from Zn(OH)2) = ? (x moles)


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(Answer)

X moles = 0.22 moles 0.0825 moles = 0.1375moles OH- from Zn(OH)2

Zn(OH)2 Zn2+ + 2OH-

If you we have 0.1375 moles of OH- = 0.1375/2moles Zn(OH)2 = 0.06875 moles

So Zn(OH)2 = 0.06875 mol x 99.41 g/mole = 6.8344 gMass percentage = 6.8344/10 x 100 = 68.34%

4.6.2 Stoichiometry for


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4.6.2 Stoichiometry forReactions in Solution

Step 1

Identify the species present and determine what

reaction occurs.

Step 2

Write the balanced net ionic equation for thereaction.

Cont: 4.6.2 Stoichiometry for


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Con t: 4.6.2 Stoichiometry forReactions in Solution

Step 3

Calculate the moles of reactants.

Step 4 Determine which reactant is limiting.

Step 5 Calculate the moles of product(s), convert to grams

or other units as required.

l 3


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Example 13

Calculate the mass of PbSO4, when 1.25 L of

0.0500 M Pb(NO3)2 and 2.00 L of 0.0250M Na2SO4

are mixed.

Molecular Equation:

Pb(NO3)2(aq) + Na2SO4(aq) PbSO4(s) +

2NaNO3(aq)

E l 13 (A )


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Example 13 (Answer)

Ionic Equation:

2Na+(aq) + SO42-(aq) + Pb2+(aq) + 2NO3

-(aq)

PbSO4(s) + 2Na+(aq) + 2NO

3

-(aq)

Net ionic equation:

SO4

2-(aq) + Pb2+(aq) PbSO4

(s)

E l 13 (A )


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Example 13 (Answer)

Moles of reactants

moles of Pb2+

=1.25 L 0.0500 mol = 0.0625 mol Pb2+

L

moles of SO42-

=2.00 L 0.0250 mol = 0.0500 mol SO42-

L

E l 13 (A )


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Example 13 (Answer)

Limiting Reactants

Pb2+ and SO42- react in a 1:1 ratio , the amount of

SO42- will be the limiting factor.

Grams of product

0.0500 mol PbSO4 303.3 g PbSO4 =15.2 gPbSO41 mol PbSO4

E i 4 1


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Exercise 4.1

What mass of NaOH is required to precipitate all

the Fe2+

ions from 50.0 mL of 0.200 M Fe(NO3)2solution?

Answer: 0.800 g NaOH


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END of CHAPTER 4


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END of CHAPTER 4

chapter 4 jan12

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