acid base titration

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Chapter14 p Fundamentals of Analytical Chemistry Eighth Edition

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  • Chapter14 p

    Fundamentals of

    Analytical Chemistry Eighth Edition

  • Chapter14 p

    Principles of Neutralization Titrations

    CHAPTER

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    Acids and bases are very important in the environment. Acid rain falling on the surface waters of lakes and ribers can cause these waters to become acidic.

    Lime stone (calcium carbonate), which reacts with CO2 and H2O to form bicarbonate. Bicarbonate in turn neurtalizes acids tomaintain the pH relatively constant.

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    Figure 9F-1Changes in p H of lakes between 1930 and 1975.

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    Figure 9F-2Effect of pH of lades on their fish population.

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    Figure 9F-3Effect of presence of limestone on pH of lakes in the United States. Shaded areas contain little limestone.

  • Figure 9F-4Sulfur dioxide emissions from selected plants in the United States Have dropped below the levels required by law.

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    Titrimetric Method

    Titrimetric methodsofanalysisarecapableofrapidandconvenientanalytedeterminationswithhighaccuracyandprecision.

    Titrimetric analysisisbasedonthecompletereactionbetweentheanalyte andareagent,thetitrant.

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    aA +tT products

    whereAandTrepresenttheanalyteandtitrant,respectively,andaandtarethestoichiometric coefficients.

    Titrationsareoftenclassifiedbythenatureofthistitrationreaction:acidbase,redox,precipitation andcomplexation reactions arethemostcommonreactiontypes.

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    Volumetric Titration

    Forvolumetrictitrations,theamount,nA,ofanalyte inthesamplecanbecalculatedusing

    nA =a/tCTVT

    whereCT istheconcentrationofthetitrant,andVT isthevolumeoftitrant neededtoreachtheendpoint.Againaandtarethestoichiometric coefficients.

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    Quantitativedeterminationoftheanalyteconcentrationrequiresthefollowing: Thereisastoichiometric reactionbetweenanalyte andtitrant.Thisreactionshouldbefastandcomplete,andthevaluesofaandtmustbeknown.

    Theconcentrationofthetitrant solution,CT,mustbeknownaccurately.Thetitrant solutionmustbestandardized eitherbypreparingitusingaprimarystandardor,morecommonly,titratingitagainstasolutionpreparedwithaprimarystandard.

    Theendpoint volumemustbemeasuredaccuratelyusinganappropriatechemicalindicatororinstrumentalmethod.Ifaninstrumentalmethodisusedtofollowtheprogressofthetitrationreaction,atitrationcurvemaybegenerated,whichallowsfortheanalysisofmixturesand/orthedetectionofinterferences.

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    Figure 13-1

    The titration process.

    Typical setup for carrying out a titration. The apparatus consists of a buret, a buretstand and clamp with a white porcelain baseto provide an appropriate background for viewing indicator changes, and a wide-mouth Erlenmeyer flask containing a precisely known volume of the solution to be titrated. The solution is normally delivered into the flask using a pipet, as shown in FIGURE 2-22.

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    Figure 13-1The titration process.

    Detail of the buret graduations. Normally, the bret is filled with titratntsolution to within 1 or 2 mL of the zero position at the top. The initial volume of the buret is read to the nearest 0.01 mL. The reference point on the meniscus and the proper position of the eye for reading are depicted in figure 2-21.

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    Figure 13-1The titration process.

    Before the titration begins. The solution to be titrated, an acid in this example, is placed in the flask and the indicator is added as shown in the photo. The indicator in this case is phenolphthalein, which turns pink in basic solution.

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    Figure 13-1The titration process.

    During titration. The titrant is added to the flask with swirling until the color of the indicator persists. In the initial region of the titration, titrant may be added rather rapidly, but as the end point is approached, increasingly smaller portions are added; at the end point, less than half a drop of titrantshould cause the indicator to change color.

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    Figure 13-1The titration process.

    Titration end point. The end point is achieved when the barely perceptible pink color of phenolphthalein persists. The flask on the left shows the titration less than half a drop prior to the end point; the middle flask shows the end point. The final reading of the buret is made at this point, and the volume of base delivered in the titration is calculated from the difference between the initial and final buret readings. The flask on the right shows what happens when a slight excess of base is added to the titration mixture. The solution turns a deep pink color, and the end pointhas been exceeded. In color plate 9, the color change at the end point is much easier to see than in this black-and-white version.

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    Figure 13-2Two types of titration curves.

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    Figure 13-3Titration curve for the titration of 50.00 mLof 0.1000 M AgNO3with 0.1000 M KSCN.

    Equivalence point -

    End point

    Equivalence point End point

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    ExampleApplication:AnalysisofChlorideinSurfaceWaters

    Chloride isfrequentlyamajoranioninsurfaceandgroundwater;certainlyisamajorconstituentofseawater.

    Althoughchlorideinfreshwaterisusuallyofgeologicalorigin,runofffromroadssaltedduringthewintermaysignificantlyincreasethechloridecontentofsurroundingstreams,riversandlakes.

    Ahighchlorideconcentrationmayimpartanoticeablysaltytastetopotablewater,andcanalsodamagemetallicpipesandgrowingplants.

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    ExampleApplication:AnalysisofChlorideinSurfaceWaters(contd) Argentometric titration ofwatersamplesisastandardmethodforchloridedetermination;

    Concentrationsinthelowppm rangemaybedetectedusingpotentiometric titration.

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    Figure 13-4Titration curve for A, 50.00mL of 0.0500 M NaClwith 0.1000 M AgNO3, and B, 50.00mL of 0.00500 M NaClwith 0.0100 M AgNO3.

    Equivalence point -

    End point

    Equivalence point End point

  • Chapter14 p356

    Figure 13-5Effect of reaction completeness on precipitation titration curves. For each curve, 50.00m of a 0.0500 M solution of the anon was titrated with 0.1000 M AgNO3. Note that smaller values of Ksp give much sharper breaks at the end point.

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    EndpointDetection Avarietyofchemicalindicatorsareusedtoindicatetheendpointofargentometric titrations,includingtheFajans,Volhard,andMohr methods

    Asilverwireorringisasufficientindicatorelectrodeforpotentiometric titrationsusingAgNO3,whileafluorideISEissuitableforpotentiometric endpointdetectionforfluorideanalysisusingLa3+orPb2+ titrant solutions.

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    adsorption

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    without

    ()

    ()

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    Thiocyanate.Iron(III) Complex

    a. Thiocyanate reacts with iron(III) to produce a deep, red color:

    Fe3+ + SCN1- FeSCN2+ (red)

    The red color can be used to detect the presence of iron (III)In titrations

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    Testing for iron(III) ions with thiocyanate ionsThis provides an extremely sensitive test for iron(III) ions in solution. If you add thiocyanate ions, SCN-, (from, say, sodium or potassium or ammonium thiocyanate solution) to a solution containing iron(III) ions, you get an intense blood red solution containing the ion [Fe(SCN)(H2O)5]2+.

  • Chapter14 pIndicator fluorescein

    adsorbed

    adsorption

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    adsorbed

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    adsorbed

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    adsorption

    adsorbed

    adsorbedadsorbed

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    adsorption

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    2-Furoic acid

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    Thestandardreagentsusedinacid/basetitrationsarealwaysstrongacidsorstrongbase,mostlycommonlyHCl,HClO4,H2SO4,NaOH,andKOH.

    Weakacidsandbasesareneverusedasstandardreagentsbecausetheyreactincompletelywithanalytes

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    Acid/Base Indicators

    HIn + H2O In- + H3O+acid color base color

    In + H2O InH+ + OH-base color acid color

    Ka = [H3O+][In-]/[HIn]

    [H3O+] = Ka ([HIn]/ [In-])

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    Figure 14-2Indicator color as a function of pH (pKa=5.0).

    AcompletecolorchangewhenthepHofthesolutioninwhichitisdossolved changesfrom4to6.

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    Figure 14-1Color change and molecular model for phenolphthalein.

  • Chapter14 pColorless Pink Magenta

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  • Chapter14 pYellow Blue

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    2.28

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    pH 2.28

    pH 7.00

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    Beyond the equivalence point, we first calculate pOH and then pH.

    pH =pKw pOH = 14.00 pOH

    Kw=[H3O+][OH-]- log Kw = -log [H3O+][OH-] = -log[H3O+]-[OH-]

    pKw = pH + pOH-log 10-14 = pH + pOH =14

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    Figure 14-3Titration curves for HCl with NaOH. Curve A:50.00mL of 0.0500 M HCl with 0.1000 M NaOH. Curve B: 50.00 mL of 0.000500 M HCl with 0.001000 M NaOH.

  • Chapter14 p377

    Figure 14-4Titration curves for NaOH with HCl. Curve A:50.00 mLof 0.0500 M NaOHwith 0.1000 M HCl. Curve B: 50.00 mL of 0.00500 M NaOHwith 0.0100 M HCl.

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    HA+H2O H3O+ + A-

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    A-+H2O HA + OH-

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    Henderson-Hasselbalch equation is used to calculate the pH of buffer solutions.

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    HA+H2O H3O+ + A-

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    A-+H2O HA + OH-

    10-14/6.31x10-5

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    Kb

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    Figure 14-5Curve for the titration of acetic acid with sodium hydroxide. Curve A: 0.1000 M acid with 0.1000 M base. Curve B: 0.001000 M acid with 0.001000 M base.

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    Figure 14-6The effect of acid strength (dissociation constant) on titration curves. Each curve represents the titration of 50.00 mL of 0.1000 M acid with 0.1000 M base.

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    Figure 14-7The effect of base strength (Kb) on titration curves. Each curve represents the titration of 50.00mL of 0.1000 M base with 0.1000 M HCl.