ANALISA TITRIMETRI

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.


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.


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;


HIn + H2O

In- + H3O+

The behaviour of an base type indicator is described by the equilibrium; equilibrium;

In + H2O

InH+ + OH-


The equilibrium constant takes the form: form:

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

Rearranging:

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


pH (acid colour) = -log(Ka . 10) = pKa +1 10) pH (base colour) = -log(Ka / 10) = pKa -1 10) Therefore; Therefore; indicator range = pKa 1


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.



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


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


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.



Examples: HCl,

HNO3 KOH, Na2CO3 constant

NaOH,

Standards:anhydrous Na2CO3 and

boiling HCl.


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


Effect of Concentration Effect of reaction completeness Indicator choice; Feasibility of titration



Polyfunctional acids and bases Typified by more than one dissociation reaction.


Phosphoric acid

Yield multiple end points in a titration.



Sulphuric Acid Unusual because one proton behaves as a strong acid and the other as a weak acid (K2 = 1.20 x 10-2).


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.


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.


Inorganic Ligands Water Ammonia Halides

Organic Ligands Cyanide Acetate


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 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.


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


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.


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.


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


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)]


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.


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+


Typical indicators are: are: Murexide Solochrome Calmagite Bromopyrogallol Xylenol

black red

orange


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.


Used for the determination of many anions including: including: halides divalent

anions mercaptans certain fatty acids


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


Endpoint determination is by coloured indicators (usually back titrations) or turbidity methods. methods.

The most accurate is the VOLHARD METHOD. METHOD.


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

ANALISA TITRIMETRI

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