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Titrimetric Analysis
Quantitative chemical analysis carried out determining the volume of a solution accurately known concentration which required to react quantitatively with measured volume of the substance to determined. determined.
by of is a be
Classification
Neutralisation Reactions Complex Formation Reactions Redox Reactions Precipitation Reactions
Basics
Equivalence and end points Standards
Basics
Equivalence and end points Precise
and accurate titrations require the reproducible determination of the end point which either corresponds to the stoichiometric point of the reaction or bears a fixed and measurable relation to it. it.
Basics
Equivalence and end points Monitor
a property of the titrand which is removed at the end point. Monitor a property which is readily observed when excess titrant has been added. Two main methods Coloured indicators Electrochemical
techniques.
Basics
Colour Change Indicators Common In
to a wide variety of titrations. titrations.
general terms a visual indicator is a compound that changes from one colour to another as its chemical form changes. changes. = InB + nX where X may be H+, Mn+ or e-, and the colour is sensitive to the presence of H+, Mn+, oxidants or reductants. reductants.
InA
Basics
An indicator constant is defined as: as: KIn = [InB][X]n / [InA] [X]n = KIn ([InB] / [InA]) npX = pKIn + log10([InB] / [InA]) pH = pKa + log10 ([InB] / [InA])
Basics
Potentiometric Measurements Measuring the change in potential during the titration. titration. Acid-base Acid-
titrations. titrations. titrations. titrations.
Precipitation Redox
titrations. titrations.
Basics
Monitor the change of Ecell during the course of a titration where the indicator electrode responds to one of the reactants or the products. products. A plot of Ecell against the volume of titrant is obtained. obtained. Precision of better than 0.2%.
Basics
Basics The
Nernst Equation
aA + bB + + ne- = xX + yY + ...
[X]x[Y]y... 0 - RT ln E=E [A]a[B]b... nF
Basics
RT/F ln 10 = 0.059158 V thus:
E = E0 - (0.059 V/n) log10 And
[X]x[Y]y... [A]a[B]b...
E = E0 at unity concentrations
Basics
Conductimetric Indication The
electrical conductance of a solution is a measure of its current carrying capacity and is determined by its total ionic strength. is a non-specific property. non-
It
Conductance
is defined as the reciprocal of resistance (Siemans, ;-1).
Basics
A conductance cell consists of two platinum electrodes of large surface area. area. 5-10 V at 50 -10,000 Hz is applied. 10, applied. Control of temperature is essential. essential.
Basics
AcidAcid-base titrations especially at trace levels. Relative precision better than 1% at all levels. Rate of change of conductance as a function of added titrant used to determine the equivalence point. High concentrations of other electrolytes can interfere.
Basics
Basics
Standards Certain
chemicals which are used in defined concentrations as reference materials. standards. standards.
Primary
Secondary
Basics
Primary Standards Available
in pure form, stable and easily dried to a constant known composition. in air.
Stable High
molecular weight. soluble. and rapid reactions.
Readily
Undergoes stoichiometric
Basics Acid-base Acid-
reactions.KH(C8H4O4), HCl (cbpt.)
Na2CO3, Na2B4O7,
Complex
formation reactions. AgNO3, NaCl reactions. AgNO3, KCl reactions.
Precipitation
Redox
K2Cr2O7, Na2C2O4, I2
Basics
Secondary StandardsA
substance that can be used for standardisations, and whose concentration of active substance has been determined by comparison to a primary standard. standard.
Classification
Neutralisation Reactions Complex Formation Reactions Redox Reactions Precipitation Reactions
Neutralisation Titrations
The neutralisation reactions between acids and bases used in chemical analysis. analysis. These reactions involve the combination of hydrogen and hydroxide ions to form water. water.
Neutralisation Titrations
For any actual titration the correct end point will be characterised by a definite value of the hydrogen ion concentration. concentration. This value will depend upon the nature of the acid and the base, the concentration of the solution and the nature of the indicator. indicator.
Neutralisation Titrations
A large number of substances called neutraneutralisation indicators change colour according to the hydrogen ion concentration of the solution. solution. The end point can also be determined electrochemically by either potentiometric or conductimetric methods. methods.
Theory of Indicator Behaviour
An acid/base indicator is a weak organic acid or a weak organic base whose undissociated form differs in colour from its conjugate base or conjugate acid form. form. The behaviour of an acid type indicator is described by the equilibrium; equilibrium;
Theory of Indicator Behaviour
HIn + H2O
In- + H3O+
The behaviour of an base type indicator is described by the equilibrium; equilibrium;
In + H2O
InH+ + OH-
Theory of Indicator Behaviour
The equilibrium constant takes the form: form:
[H3O+][In-] = Ka [HIn]
Rearranging:
[HIn-] [H3O+] = Ka [In-]
Theory of Indicator Behaviour
pH (acid colour) = -log(Ka . 10) = pKa +1 10) pH (base colour) = -log(Ka / 10) = pKa -1 10) Therefore; Therefore; indicator range = pKa 1
Theory of Indicator Behaviour
The human eye is not very sensitive to colour change in a solution containing In- and HIn. HIn. Especially when the ratio [In-] / [HIn] is greater than 10 or less than 0.1. Hence the colour change is only rapid within the limited concentration ratio of 10 to 0.1.
Theory of Indicator Behaviour
Neutralisation Titrations
Strong acids and bases Weak acids Weak bases Polyfunctional acids Applications
Neutralisation TitrationsStrong acids and bases. When both reagent and analyte are strong electrolytes, the neutralisation reaction can be described by the equation:
H3O+ + OH-
2H2O
Neutralisation Titrations
The H3O+ concentration in aqueous solution comprises of two components. The The
reaction of the solute with water. dissociation of water.
[H3O+] = CHCl + [OH-] = CHCl [OH-] = CNaOH + [H3O+] = CNaOH
Neutralisation Titrations
Using these assumptions you can calculate the pH of a titration solution directly from stoichiometric calculations and therefore simulate the titration curves. This is useful in determining the correct indicator for a new titration.
Neutralisation Titrations
Neutralisation Titrations
Examples: HCl,
HNO3 KOH, Na2CO3 constant
NaOH,
Standards:anhydrous Na2CO3 and
boiling HCl.
Neutralisation Titrations
Weak acids and bases Examplesacid Sodium cyanide Ethanoic
Four types of calculation are required to derive a titration curve for a weak acid or base.
Neutralisation Titrations Solution
contains only weak acid. pH is calculated from the concentration and the dissociation constant. additions of the titrant the solution behaves as a buffer. The pH of each buffer can be calculated from there analytical concentrations. the equivilence point only salt is present and the pH is calculated from the concentration of this product. the equivilence point the pH is governed largely by the concentration of the excess titrant.
After
At
Beyond
Neutralisation Titrations
Effect of Concentration Effect of reaction completeness Indicator choice; Feasibility of titration
Neutralisation Titrations
Neutralisation Titrations
Polyfunctional acids and bases Typified by more than one dissociation reaction.
Neutralisation Titrations
Phosphoric acid
Yield multiple end points in a titration.
Neutralisation Titrations
Neutralisation Titrations
Sulphuric Acid Unusual because one proton behaves as a strong acid and the other as a weak acid (K2 = 1.20 x 10-2).
Neutralisation Titrations
Applications: Determination
of the concentration of analytes which are either acid or bases. Determination of analytes which can be converted to acids or bases.
Complexometric Titrations
Titrations between cations and complex forming reagents. reagents. The most useful of these complexing agents are organic compounds with several electron donor groups that can form multiple covalent bonds with metal ions. ions.
Complexometric Titrations
Most metal ions react with electron-pair electrondonors to form coordination compounds or complex ions. ions. The donor species, or LIGAND, must have LIGAND, at least one pair of unshared electrons available. available.
Complexometric Titrations
Inorganic Ligands Water Ammonia Halides
Organic Ligands Cyanide Acetate
Complexometric Titrations
The number of bonds a cation forms with an electron donor is called the COORDINATION NUMBER. NUMBER. Typical values are 2, 4 and 6. The species formed as a result of coordination can be electrically positive, neutral or negative. negative.
Complexometric Titrations
Complexometric methods have been around for more than a century. century. Rapid expansion in the 1940s based on 1940s a class of coordination compounds called CHELATES. CHELATES.
Complexometric Titrations
A chelate is produced when a metal ion coordinates to two or more donor groups within a single ligand. ligand. For example the copper complex of glycine. glycine.
Complexometric TitrationsNH2 Cu2+ + 2 H C H O C O Cu C N H2 H2 N C H2 H2 O C OH O + 2H + O
C
Complexometric Titrations
A ligand with a single donor group is called unidentate. unidentate. Glycine is bidentate. bidentate. Tri, tetra, penta and hexadentate chelating agents are also known. known.
Complexometric Titrations
Multidentate ligands have two advantages over unidentate ligands. ligands. They react more completely with cations to provide a sharper endpoint. endpoint. The reaction is a single step process. process.
Complexometric Titrations
Tertiary amines that also contain carboxylic acid groups form remarkably stable chelates with many metal ions. ions.
Ethylenediaminetetraacetic Acid EDTACH 2COO H N CH2 CH2 N CH 2COO H
H O O C CH 2 HOOCCH2
Complexometric Titrations
EDTA can complex a large number of metal ions. ions. Approximately 40 cations can be determined by direct titration. titration. EDTA is usually used as the disodium salt, Na2H2EDTA H2EDTA2- + M2+ p[M(EDTA)]2- + 2H+ p[M(EDTA)]
Complexometric Titrations
Because EDTA complexes most cations, the reagent might appear at first glance to be totally lacking in selectivity. selectivity. However, great control can be acheived by pH regulation and the selection of suitable indicators. indicators.
Complexometric Titrations
Indicators are generally complexing agents which undergo a colour change when bonded to a metal ion. ion.
H2EDTA2- + [M(Ind)] p[M(EDTA)]2- + Ind2- + 2H+
Complexometric Titrations
Typical indicators are: are: Murexide Solochrome Calmagite Bromopyrogallol Xylenol
black red
orange
Complexometric Titrations
Typical applications: applications: Determination
of cations Hardness of water
Redox TitrationsBasics Potassium Permanganate Potassium Dichromate Cerium IV Iodine
Redox Titrations
Basics Electrode Indicators
Potentials
Redox Titrations
Electrode Potentials Derived
from Nernst equation. Calculations of cell potentials leads to theoretical titration curves. EOX = ERED = Esystem
Redox Titrations
Redox Titrations
Indicators Potentiometric Coloured EIn
indicators
= EOX = ERED = Esystem Starch / Reduction Indicators
Specific:
Oxidation
Redox Titrations
Redox Titrations
Potassium Permanganate MnO4- + 8H+ + 5e- p Mn2+ + 4H2O
Standardisation Sodium
oxalate or arsenic (III) oxide
Many Analyses
Redox Titrations
Hydrogen Peroxide:
2MnO4- + 5H2O2 + 6H+ p 2Mn2+ + 5O2 + 8H2O
Nitrites:
2MnO4- + 5NO2- + 6H+ p 2Mn2+ + 5NO3- + 3H2O
Redox Titrations
Persulphates: Add
an excess of iron (II) S2O82- + 2Fe2+ + 2H+ p 2Fe3+ + 2HSO4 The
excess iron (II) is determined by back titration against standardised permangenate. MnO4- + 8H+ + 5Fe2+ p Mn2+ + 5Fe2+ + 4H2O
Redox Titrations
Potassium Dichromate CrO72- + 14H+ + 6e- p 2Cr3+ + 7H2O
Standardisation Against
metallic iron 1 mole K2CrO7 = 6 moles Fe
Redox Titrations
Iron (II): CrO72- + 14H+ + 6Fe2+ p 2Cr3+ + 6Fe3+ 7H2O Indicators
include N-phenylanthranilic acid and sodium Ndiphenylamine sulphonate.
Chlorates: Reduced The
with an excess of iron (II)
excess iron (II) is determined by back titration against standardised dichromate.
Redox Titrations
Iodine Iodometric Titrations I2 + 2e- p2Ip2I
Standardisation Standardised
sodium thiosulphate or arsenic (III) oxide
Many Analyses
Redox Titrations
Hydrogen Peroxide:
H2O2 + 2I- + 2H+ pI2 + 2H2O pI
Thiosulphates:
2S2O32- + I2 p S4O62- + 2I-
Hydroxyl Groups:
2OH- + I2 p IO- + H2O + 2I-
Redox Titrations
Others: Copper Dissolved Chlorine Arsenic
oxygen
(V) Sulphides etc..........
Redox Titrations
Cerium (IV) Sulphate Very strong oxidising agent (1.43V) Ce4+ + e- pCe3+ pCe
Standardisation Sodium
oxalate or arsenic (III) oxide
Many Analyses
Precipitation Titrations
Titrations between analytes and reagents resulting in the formation of a precipitate. The most useful of these precipitating reagents is silver nitrate. nitrate. Titrimetric methods based upon the use of silver nitrate are sometimes called Argentometric titrations. titrations.
Precipitation Titrations
Used for the determination of many anions including: including: halides divalent
anions mercaptans certain fatty acids
Precipitation Titrations
Precipitation titrations are based on the SOLUBILITY PRODUCT of the salt, KSP. The smaller KSP, the less soluble the silver salt and the easier it is to determine the endpoint
Precipitation Titrations
Endpoint determination is by coloured indicators (usually back titrations) or turbidity methods. methods.
The most accurate is the VOLHARD METHOD. METHOD.
Precipitation Titrations
VOLHARD METHOD A back titration of thiocyanate ions against the excess silver ions using an iron (II) salt as the indicator. indicator.
Precipitation TitrationsAg+ + SCNFe3+ + SCNAgSCN FeSCN2+
Blood Red