Titrimetry(anEm`pnmQwQy)

Quantitative method

The amount of substance(titrand) is calculated from the measured amount (usually volume)of a reagent solution(titrant or standard solution) of known concentration.

The process is called ‘Titration’

For a titrimetric determination1. The reaction must be fast.2. The reaction must be stoichiometric.3. It must go to completion with no side reactions.4. Should alter a physical or chemical property at the

completion of the reaction.5. An indicator should be available to determine this

change.

Classification of titrationsFrom type of reaction

Acid- base titrations Precipitation titrations Redox titrations Complexometric titrations

From type of titrant usedo Acidimetryo Alkalimetryo Argentimetryo Permanganometryo Iodometry and iodimetry

Classification of titrationsFrom the method of end point detection

Visual titrations Electrometric (Potentiometric titrations)

(Amperometric titrations) (Coulometric titrations)

(Conductometric titrations) Photometric (Colorimetric titrations….) Gravimetric titrations

Standard solution Primary standard (pY`}mQk pY`m`NQky) A compound that can be weighed and diluted to get an

exact concentration. These must be of known purity, preferably 100%. reaction with the reagent to be analyzed should be

stoicheometric, complete, fast and selective. easy to handle (weighing & dissolving). high molecular weight (to minimize weighing errors). readily available, inexpensive and easy to dry. stable in conditions used (in air & in solution). should not absorb water (hygroscopic) or CO2. soluble in the medium used (long term).

Examples of primary standards Bases (For standardizing acids) Sodium carbonate (Na2CO3)-commonly used but low MW. 4-aminopyridine – high purity and stability. Sodium tetraborate (Na2B4O7)

Acids (For standardizing bases) Benzoic acid Potassium hydrageniodate {KH(IO3)2} Potassium hydrogen phthalate (KHP)- most commonly

used.high MW(204.2 g/mol). High purity, thermally stable and reacts fast with NaOH and KOH.

2-Furonic acid- stronger acid than KHP.

Other common primary standards Ag, Ag(NO3), NaCl, KCl, KBr

K2Cr2O7, KBrO3, KIO3, Na2C2O4, As2O3,

Metals (Zn, Cu, Fe, Mg, Ni, Mn,…)

EDTA

**(NaOH,KOH, HCl, HNO3, H2SO4, H3PO4, KMnO4, Na2S2O3 are not primary standards.)

Secondary standards (q~vQwQk pY`m`NQky) A solution with an approximate concentration is

prepared and the exact concentration is established using a primary standard.(“standardized”)

The second material is then considered a secondary standard.

Equivalence point (smkw` l]&y) (theoretical end point)

The point where enough titrant (stoichiometric amount) is added to completely react with the titrand(analyte).

Endpoint (an~wl]&y) The amount of titrant required for the detection

of the equivalence point.

In a titration we ideally want the equivalence point and the end point to be the same.

But this seldom happens due to methods used to observe the end points.

The deviation from the true value is known as the “titration error”.

Acid-Base titrationsam|l-x’;~m anEm`pnAcid- base theories1. Arrhenius theory – acids are substances that

yield H+ in aqueous solutions. bases are substances that yield OH-.

2. Lewis theory – acid – compounds that can accept a lone pair of

electrons. base – compounds that can donate a lone pair of

electrons.

3. Bronsted-Lowry theory – acid – compounds/ions that can give up a proton

(proton donor) base – compounds/ions that can take in a proton (proton acceptor)

HA H+ + A- Acid Conjugated base

**(stronger the acid weaker the conjugated base)

BOH + H+ B+ + H2O Base Conjugated acid

Strong acids and bases Acids and bases that fully dissociate in aqueous

solution. eg.

• HCl, HNO3, H2SO4, HClO4 ……..• NaOH, KOH, NaNH2 …….

**in water all are of equal strength but strength depends on the medium used. eg. HCl is a weak acid in a acetic acid medium.

Weak acids and bases Partial dissociation in the medium.Acids HAc H+ + Ac-

Acid dissociation constant = Ka = [H+][Ac-] [HAc]Bases BOH B+ + OH-

Base dissociation constant = Kb = [B+][OH-] [BOH]

Autoprotolysis of water H2O + H2O H3O+ + OH-

Kw = [H3O+][OH-] = ion product of water [H2O]2 = water dissociation constant

Kw = [H+][OH-] = 1.00 x 10-14 at 250C.The value of Kw depends on temperature.

T (0C) Kw (x 10-14)

010254060

0.120.291.012.929.61

The concept of pH pH = -log[H+]

Kw = [H+][OH-]Hence pKw = pH + pOH = 14 at 250C.

at 250C pH of pure water is 7.0 and neutral. if pH > 7.0 solution is basic if pH < 7.0 solution is acidic.

Acidity constant of a baseFor a base BOH B+ + OH-

Kb = [B+][OH-] [BOH]

For the conjugated acid B+ + H2O BOH + H+

Ka = [BOH][H+] [B+][H2O] Ka.Kb = Kw

Hydrolysis of salts (lvN jlvQc’|@j~qny)

4 classeso Derived from strong acid + strong base – eg. NaCl neutralo weak acid + strong base – eg. NaAc basico strong acid + weak base – eg. NH4Cl acidico weak acid + weak base – eg. NH4Ac ??

1. Salt of strong acid and strong base eg. NaCl + H2O Na+ + Cl-

strong electrolyte

**pH of the solution is 7.00 (at 250C.)

Salt of weak acid and strong base

eg. NaAc + H2O Na+ + Ac-

A- + H2O HA + OH-

1-x x x

Kh = hydrolysis constant = [HA][OH-] = Kw/Ka

[A-][H2O]But [HA] = [OH-]

Kh =[OH-]2/[A-] since Kh is very small, [A-]= c

c = concentration of the salt in the solutionKh

= Kw= [OH-]2

Ka c

or [OH-] = c.Kw/Ka = Kw/[H+]

Hence [H+] =Kw/ c.Kw/Ka = Kw.Ka /c

pH = ½pKw +½pKa -½pc(or +½log c)

eg. For an 0.05M aqueous solution of sodium benzoate at 250C.

pKw=14.0pKa(benzoic acid)=4.2 (Ka is 6.37 x 10-5)pc =1.30 (for 0.05M)

pH of the solution =7.0 +2.1 – 0.65 = 8.45

Salt of strong acid and weak base

M+ + H2O MOH + H+

1-x x x

eg. NH4Cl + H2O NH4+ + Cl-

Kh = [MOH][H+] = Kw = [H+]2

[M+][H2O] Kb [M+] = c

[H+] = Kw.c/Kb

pH = ½pKw - ½pKb+ ½pc

eg. For an 0.2M aqueous solution of ammonium chloride at 250C.

pKw=14.0pKb(ammonia)=4.74 (Kb is 1.8 x 10-5)pc =0.70 (for 0.2M)

pH of the solution =7.0 -2.37 + 0.35 = 4.98

Salt of weak acid and weak base

eg. NH4Ac + H2O NH4+ + Ac-

M+ + A- + H2O MOH + HAa-x a-x x x

Kh = [MOH][HA] = Kw = x2

[M+][A-][H2O] Kb.Ka (a-x)2

But from acid dissociation[H+] = Ka[HA]/[A-] =Ka.x/(a-x)

[H+] =KaKh = KaKw/KaKb = Kw.Ka/kb

pH = ½pKw + ½pKa -½pKb it’s independent of concentrations.

eg. For an 0.2M aqueous solution of ammonium methanoate at 250C.pKw=14.0pKb(ammonia)=4.74 (Kb is 1.8 x 10-5)pKa(methonoic acid) = 3.75 (Ka is 1.77 x 10-4)

pH of the solution =7.0 +1.88 –2.37 = 6.51

Titration curves (anEm`pn vkY) Plot of a variable parameter during a titration

against the titrant volume

For acid base titrations the parameter is pH.

pH

Volume of titrant added acid

base

Strong acid-strong base titrationspYbl am|l-pYbl x;~m anEm`pn

Net reaction H3O+ + OH- 2H2O

0.10M NaOH

0.10m HCl

Initial pH=-log[H+] = 1.0pH prior to equivalence point

+

At equivalence pointThe cation and anion does not get involve in reactions.pH = 7.00 (at 250C)

After equivalence point (over titration)

For the titration of 50.00 mL of 1.0M HCl with 1.0M NaOHVolume of NaOH added (mL)

pH of the flask

0.00 (initial)10.0030.0040.00 (prior to 49.00 endpoint)49.9550.00 (end point)50.0551.00 (after60.00 endpoint)80.00

0.000.180.600.951.993.307.0010.7011.9912.9613.36

The size of the vertical region will depend on the concentration of the acid and the base

If the base is the titrant the curve is a mirror image of the one discussed.

Weak acid strong base titrations qEbl am|l-pYbl x;~m anEm`pn

eg. Titration of 50.00 mL of 1.0M acetic acid (Ka= 1.0 x 10-5) with 1.0M NaOH.

Initial pHUse Oswald’s dilution law

HA A- + H+ c(1-x) cx cx

Ka = (cx)2/c(1-x)

[H+] = Ka.ca or pH=½pKa-½ log ca

**for above eg. pH =2.5(The pH is much higher than that of the strong acid and weaker the acid higher the pH)Titration of very weak acids pKa<10-10) are not feasible.

pH Before the equivalence point Reaction HA + OH- H2O + A-

Ka = [H+][A-] [HA]

[H+] = Ka[HA] [A-]

Since Ka is small can neglect the HA dissociationHence [HA] acid left & [A-] base reacted=salt formed

[H+] = Ka(CaVa- CbVb) CbVb

pH=pKa +log[A-]/[HA]Henderson-Hasselbalch equation

Total moles of acid left = (Ca.Va- Cb.Vb)/1000Moles of salt formed =CbVb/1000

For our example, Titration of 50.00 mL of 1.0M acetic acid (Ka= 1.0 x 10-5) with 1.0M NaOH.

When 5.00 mL of NaOH is added[H+] = Ka(CaVa- CbVb) CbVb

[H+] = 1 x10-5.(1.0x50.00-1.0x5.00) 1.0x5.00

= 4.04

Half equivalence point when [HA]=[A-]

Then, pH = pKa

pH=pKa +log[A-]/[HA]Henderson-Hasselbalch equation

At equivalence point All acid has reacted. And solution is like when a

salt is dissolved

pH = ½pKw +½pKa -½pc(or +½log c)

‘C’ is the concentration of salt at the equivalence point.which is equal to ca.Va/(Va+Vb)

For our eg. pH= 7.0 +2.5 -0.15 =9.35

After the equivalence point, the excess strong base governs the pH of the medium and calculations are same as in strong acid- strong base case.

For the titration of 50.00 mL of 1.0M HAc with 1.0M NaOHVolume of NaOH added (mL)

pH of the flask

0.00 (initial)10.0030.0040.00 (prior to 49.00 endpoint)49.9550.00 (end point)50.0551.00 (after60.00 endpoint)80.00

2.504.404.825.606.698.009.3511.0012.2913.2213.57

Usually the weak acid is the titrand. But if the acid is in the burette,

Weak acid

Strong base

“Till equivalence point is like strong acid-strong base case.”pH at equivalence point ???

After equivalence point[HA] =ca.Va

*/(Va+Vb)[salt]=[A-]= cb.Vb/(Va+Vb)

[H+]=Ka.[HA]/[A-]=Ka.(ca.Va*/ cb.Vb)

**Since the titrant is weak need time to dissociate and react.

Polyfunctional acids with strong base For a diprotic acid

H2A H+ + HA-

HA- H+ + A2-

For the 2 steps to be titrated separately (to get 2 well defined equivalence points) Ka1

Ka2

>103 (for 1 - .001M acids)

Strong acid weak base titrations pYbl am|l- qEbl x;~m anEm`pn Treat like previous case using Kb instead of Ka

eg. Titration of 50.00 mL of 1M NH4OH with 1M HCl NH4OH NH4

+ + OH-

c(1-x) cx cx Kb = [NH4+][OH-]

[NH4OH]

Initial pH[OH-]=Kb.cb

Before equivalence point pKb =pOH – log [NH4

+]/[NH4OH] ([salt]/[base])

But pH +pOH =pKw

Hence, pH = pKw- pKb - log[salt]/[base]

[salt]/[base]=amount of acid added/amount of base left = ca.Va/(cb.Vb-ca.Va)

At equivalence point Treat as hydrolysis of salt of weak base strong acid

pH = ½pKw - ½pKb+ ½pc

c= concentration of salt at equivalence point = cb.Vb/(Va +Vb)

**Equivalence point is in acidic pH’s & depend on the strength of the base.

Weak acid –weak base titrationsqRbl am|l-qRbl x;~m anRm`pn Feasibility depends on both dissociation

constants Ka&Kb.

Titration of anion of weak acid (conjugated base) with strong acid

eg. Acetate ion Ac- +H2O HAc + OH-

H+

NaAc(aq) + HCl(aq) HAc + Na+ + Cl-

eg. Borate ionB4O7

3- + 2H+ +5H2O 4 H3BO3

Ka of boric acid 6.4 x10-10

Titration of carbonate & bicarbonate with a strong acid

CO32- + H+ HCO3

-

HCO3- + H+ H2CO3 H2O + CO2

For carbonic acid Ka1=4.5 x10-7 molL-1 & Ka2=5.6 x10-11 molL-1

Hence for carbonate ion Kb1=1.8 x10-4 molL-1 & Kb2=2.2 x10-8 molL-1

Kb1/Kb2 104 and 2 separate equivalence points

Acid-base indicatorsam|l-x;~m qr\Xk Equivalence point detection’ in acid base titrations

is commonly done by visual detectors.

Spectrometric, potentiometric, thermometric detection is also used.

They are weak organic acids or bases.

Their color differs from the color of their conjugate base or acid.

o One color indicators eg. Phenolphthalein (colorless-magenta)

o Two color indicators eg. Methyl orange (red-yellow)

o They change color in a definite pH range called “transition range”.

o This range depends on the acidity constant(or basisity constant) of the indicator.

o In pH above the range indicator is predominantly in its base form & in lower pHs in its acidic form.

HIn H+ + In-

(acid form) (base form)

Indicator constant = Kin =[H+][In-]/[HIn]

pKIn= pH - log [In-]/[HIn]

pH = pKIn + log [In-]/[HIn]

**When both forms are present the eye detects the predominant color(usually taken as when 10 times more than the other color)

pH = pKIn 1

In order to be a good acid base indicator It should be weaker than the species reacting

Must be present in low concentrations (not significantly interfere with the equivalence point)

Must be high molecular weight molecules.

The 2 colors should be markedly different (ideally complementary)

The color/s should be intense to give a sharp end point.(the transition range should be small)

Choosing of a suitable indicator• The color change range (pKIn) & the vertical portion of the titration.• Ka,Kb, of the acids & bases•Concentration of acids & bases•Temperature

Effect on kw,Ka, Kb..Effect on Kin

Stability of colors•Solvent medium

Strength of acids & basesSolubility of indicator

Litmus Litmus is a water soluble dye extracted from certain

lichens and absorbed on to filter paper. The active ingredient of Litmus is called Erythrolitmin. Color change range 4.5- 8.3 (pKIn 6.5) Acid color red, base color blue (purple inbetween)

Phenolphthalein pKIn = 8.7 (range 8.3- 10.0) Colorless Pink (magenta) acidic basicacidic basic

*End point in basic medium

•If base in flask, the end point is when medium is colorless.colorless.

•If acid in flask, the end point is when medium is slightly pinkslightly pink

Methyl orange pkIn = 3.5 (range 3.2-4.4) Red Red orangeorange yellowyellow acidic basicacidic basic

*End point in acidic medium

•If base in flask, the end point is when medium is orangeorange. •If acid in flask, the end point is when medium is yellowyellow

Methyl red pkIn = 5.0 (range 4.2-6.2) Red Red orangeorange yellowyellow acidic basicacidic basic

Bromothymol blue pkIn = 7.3 (range 6.2-7.6) yellowyellow GreenGreen blueblue acidic basicacidic basic

Strong acid(1M)-strong base (1M) titration

Strong acid(1M)-weak base(1M) titration

Weak acid(1M)-strong base (1M) titration

Weak acid(1M)-weak base (1M) titration

CO32-(1M)- strong acid titration(1M)

Screened indicators To get a pronounce color at the end point add a dye

with the indicator.

eg. Screened methyl orange Red Red orangeorange yellowyellow acidic basicacidic basic

Add a Add a blueblue dye dye purplepurple greygrey greengreen

Buffers (s~v`r]k qY`vN) A mixture of a conjugate acid –base pair It tends to resist changes in pH when an acid or

base is added. Commonly used when pH must be maintained at

a relatively constant value and in many biological systems.

L 11

Effect of adding a acid to a buffereg. if add 10 mL of 1.0 M HCl to 100 mL of pure

water at pH 7.

if add 10 mL of 1.0 M HCl to 100 mL of a solution containing1.0M HA and 1.0M A- at pH 7. (pKa=7)

The added 10 mL of acid will react with the conjugate base,converting it to the acid.

So we would have 0.09 moles of base form and 0.11 moles of the acid form.

When 100mL of 1.0M HCl is added virtually all A- is converted to HA and the calculation of pH has to be done using KA.

At this point we have exceeded the buffer region of our system.The upper limit can be calculated too.?????

Buffer capacity(s~v`r]k {`rQw`v)

The number of moles of a strong acid or base that causes 1 liter of a buffer solution to undergo a pH change of 1.00.

Compelximetric titrations (sAkWr\NmQwQk anEm`pn)

Based on reactions forming a stable complex when titrant and titrand react.

The reaction has to be fast and stoicheometric

The complex has to be soluble in the medium (insoluble complex forming reactions are considered under precipitation titrations)

Usually involves a cation of a metal (metal complexes)

M + nL MLn

Metal ligand complex (complexing agent)

Kst (or KML) = [MLn] = formation constant of complex

[M][L]n (stability constant)

Classification of ligands Monodentate – eg. Cl, Br, CN, CO, NO2, H2O, NH3

(binds to the metal, donating one pair of electrons)

Bidentate – forms two bonds with the central atom eg. NH2CH2CH2NH2 (en)-ethylene diamine“Chelating ligands”

8-hydroxyquinoline

Dimethyl glyoxime (DMG)

1,10 phenanthroline (phen)

Diphenyl carbazide (carbazone)

Polydentate ligands EDTA(ethylenediamine tetraaetic acid)

Most common complexing agent. Insoluble in water. Need to add NaOH to dissolve. Usually use the disodium salt (Na2H2Y) Both are primary standards

•Forms 1:1 complexes with most of the metals (except group 1A)•Complexes are water soluble.

The molecule contains six donor groups

Forms a hexadentate complex with metals

pKa1=1x10-2 pKa2=2.1x10-3

pKa3= 6.9 x10-7 pKa4=7.4 x10-11

2 strong acid groups, major ion is H2Y2-

Mn+ + H2Y2- MY(n-4) + 2H+

•Higher the acidity lower the formation constant (stability of complex formed)•Only strong complexes can be used at low pH. buffering is very essential in EDTA titrations

pH stability of M-EDTA complexesMinimum pH of complex metal ion

1-3 4-68-10

Zr4+, Hf4+, Th4+, Bi3+, Fe3+,Al3+,Pb2+,Cu2+,Zn2+,Co2+,Ni2+,Mn2+,Fe2+,Cd2+,Sn2+

Ca2+, Sr2+, Ba2+, Mg2+,

Titration curve similar to acid base titration curves. pM (for pH) is plotted against volume of EDTA added.EDTA is usually the titrant

pM = -log[M]

EDTA

metal

eg. 100 mL of 0.1M metal(KMY=1.0x1011) titrated with 0.1M EDTA

Initial pMpM = -log[M] =1.0

pM before the equivalence point[M] = unreacted metal ions = CmVm- CLVL

total volume Vm+VL

pM at equivalence pointAll metal ions are complexed.[complex] =[MY]= CmVm

Vm+VL(at endpoint)

But, KMY=[MY]/[M][Y] and [M]=[Y]

[M] =[MY]/KMYStronger the complex, lower the [M], higher is pM which leads to a larger vertical portion.

After the equivalence pointAmount of free ligand from dissociations of complex is negligible.[Y] = extra ligand added = CLV*L

total volume Vm+VL

[MY] = initial metal in solution = CmVm

total volume Vm+VL

KMY=[MY]/[M][Y] [M] = CmVm

KMY.(CLV*L)

The vertical portion of the curve is dependent on Stability constant of the complex pH of medium Temperature Metal and ligand concentrations

Indicators Visible and instrumental methods are used to detect

the endpoint. Visible indicators are called metalochromic indicators or

metal indicators. These form weaker complexes with metals, than EDTA. The complexed and uncomplexed(free) forms are

different in color.

L 12

M + In MIn free m-complex

Since it is easier to react with free metal EDTA will not react with MIn till free metal is almost over.

EBT (Eriochrome Black T) A triprotic acid and can act as acid base indicator

too. The metal-In complexes are wine red in color. The color of the free indicator is pH dependent.

H2In- HIn2- In3-

Red Blue OrangepH 6-7pKa1 6.3

pH 11-12pKa1 11.6

Only useful in the range of pH 7-11

Other metal indicatorsCalmagiteSame as EBT, useful range = 8.1-12.4

NASRed violet at very acidic pH & red orange at pH 3.5 Metal complex pale yellowpale yellow with Cu, Zn & Pb

DMG for detection of Ni

8 hydroxyquinolineFor Mg, Zn, Cu, Cd, Pb, In, Al, Bi, Ga, Th, Zr,

Potton & Reeders (HHSNNA)(2 hydroxy-1-(2 hydroxy-4-sulpho-1-napthylazo)-3-naphthoic acid) for CaColor change wine red to bluepH range extends beyond 13.

Other compleximetric ligands

NITA or NTA (nitrilotriacetic acid)

CDTA or DCTA (trans-1-2 diaminocyclohexane N,N,N’,N’ tetraacetic acid)

EGTA (2,2’ ethylenedioxybisethyliminodiacetic acid)

TTHA (triethylene tetraamine N,N,N’,N’’,N’’’,N’’’ hexaacetic acid)

Types of EDTA titrations1. Direct titrations Direct addition of EDTA to a metal containing solution. eg. Ca2+, Zn2+ EDTA

Metal analyte

2. Back titrations Excess EDTA is added to the analyte

solution and the unreacted EDTA is titrated with a standard metal solution.

eg. Ni2+

Used for analysis of slow reactions, buffer effected complexes,& when no indicator is available

Standard metal

+ excess EDTA

3. Replacement titrationseg. Hg2+, Ag+

For metals that does not give a clear color to the indicator.M +MgY MY + Mg

4. indirect titrations Determination of anions that form precipitates with metal

ions

eg. CO32-, CrO4

2-, S2-, SO42-

The precipitate formed is filtered, washed and excess EDTA is added. The unreacted EDTA is titrated with Mg2+.

Masking “Exclusion of the action of an interfering substance,

by adding appropriate reagent” By precipitation, oxidation, reduction complexation or a mix

of these actions.

EDTA titrationsAddition of anion of higher binding strength to the metal ion than EDTA.Masked ions does not participate in the titration.

eg. Mg2+ is titrated at pH 10, but Ba2+ can also complex.Masking agent- sodium sulfate BaSO4 will ppt. And do not even have to filter off. “precipitation masking”

CN- as masking agentWill form stable complexes with Fe, Cu, Hg, Co, Ni, Cd, Zn and noble metals

but complexes with Zn and Cd are weak.hence can “de-mask” by adding formaldehyde.

Think!- how can you get the concentration of Ca2+, Zn2+ and Ni2+ in a mixture by using EDTA titrations?

The Fe cyano-complex is colored and will interfere the end point detection. Hence H3PO4 is added to mask Fe.

OH- (high pH, above 11) will mask Mg2+

used in determination of Ca2+ in the presence of Mg2+

NH3 can mask Ag against Cl- but not I-

Oxidation will mask Cr from precipitation as hydroxide. (Cr3+ to Cr7+)

Precipitation titartions(avk~@;~pNmQwQk anEm`pn) The analyte forms a sparingly soluble complex

with the titrant. Not all precipitation reactions can be used. They have to be

• Fast reactions• Having reproducible product composition • Low solubility of precipitate• Having a method locate the endpoint

Precipitations are generally slow, co-precipitation and colored dispersed precipitates making endpoint detection hard limits precipitation titrations to Argentimetry (X & CN) and BaSO4.

Titration curve For Cl- titration with AgNO3 The plot of pCl with the concentration of

AgNO3 added

AgNO3

Cl-

eg. 100 mL of 0.1M NaCl titrated with 0.1M AgNO3 (Ksp(AgCl)=1.0x10-10)

Initial pClpCl = -log[Cl-] =1.0

pCl before the equivalence point[Cl-] = unreacted Cl- ions = CClVCl- CAgVAg

total volume VCl+VAg

pCl at equivalence pointAll Cl- ions are reacted.but due to sparingly solubilityKsp=[Ag+][Cl-]

[Cl-] =[Ag+]= Ksp

Lower the solubility product larger the vertical portion of the curve

After the equivalence pointAmount of free Ag+ from dissociation of precipitate

is negligible.[Ag+] = extra AgNO3 added = CAgV*Ag

total volume VCl+VAg

Ksp= [Ag+][Cl-]

[Cl-] = Ksp.(VCl+VAg) CAgV*Ag

The size of the vertical portion depends on the Ksp, concentration of reaction species, solvent medium and the temperature.

Side reactions(presence of impurities) and pH effects the vertical portion on some cases.

Endpoint detection Potentiometric detection use “ion selective electrode” (like pH meter)Ag/AgCl electrode is sensitive to changes in

[Ag+] and[Cl-] concentrations

•Mohr methodUse CrO4

2- ion to detect the end point. The Ksp difference between

AgCl and Ag2CrO4 will not ppt Ag2CrO4 till the endpoint. (see fractional precipitation)

Ag+ + Cl- AgCl (titration reaction)2Ag+ + CrO4

2- Ag2CrO4(at end point) brick-red

Volhard methodExcess Ag+ is added and AgCl is filtered off. Excess

Ag+ is titrated with SCN-.Fe3+ acts as indicator and give red Fe(SCN)2+ at the

end point.

• Fajans methodUse a adsorption indicator like dichlorofluorscene(for Cl-)Detect the change in primary adsorbed layer.Until the equivalence point Cl- is the primary adsorbed ion and the

outer surface is negative. Which repel the indicator.After equivalence point Ag+ is the primary ion and attracts

the indicator.

Absorbance & colorimetry