sfc113 96-97
Download
Report
Transcript sfc113 96-97
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.
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.
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
to a wide variety of titrations.
In
general terms a visual indicator is a compound that
changes from one colour to another as its chemical form
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.
InA
Basics
An indicator constant is defined 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.
Acid-base
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.
A plot of Ecell against the volume of titrant is
obtained.
Precision of better than 0.2%.
Basics
Basics
The
Nernst Equation
aA + bB + …+ ne- = xX + yY + ...
x[Y]y...
[X]
RT
0
E=E ln
a[B]b...
[A]
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.
It
is a non-specific property.
Conductance
is defined as the reciprocal of resistance
(Siemans, -1).
Basics
A conductance cell consists of two
platinum electrodes of large surface area.
5-10 V at 50 -10,000 Hz is applied.
Control of temperature is essential.
Basics
Acid-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.
Primary
standards.
Secondary
standards.
Basics
Primary Standards
Available
in pure form, stable and easily dried
to a constant known composition.
Stable
High
in air.
molecular weight.
Readily
soluble.
Undergoes
stoichiometric and rapid reactions.
Basics
Acid-base
Na2CO3,
reactions.
Na2B4O7, KH(C8H4O4), HCl (cbpt.)
Complex
formation reactions.
AgNO3, NaCl
Precipitation
reactions.
AgNO3, KCl
Redox
reactions.
K2Cr2O7, Na2C2O4, I2
Basics
Secondary Standards
A
substance that can be used for
standardisations, and whose concentration of
active substance has been determined by
comparison to a primary standard.
Classification
Neutralisation Reactions
Complex Formation Reactions
Redox Reactions
Precipitation Reactions
Neutralisation Titrations
The neutralisation reactions between acids
and bases used in chemical analysis.
These reactions involve the combination of
hydrogen and hydroxide ions to form water.
Neutralisation Titrations
For any actual titration the correct end point
will be characterised by a definite value of
the hydrogen ion 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.
Neutralisation Titrations
A large number of substances called neutralisation indicators change colour according to the
hydrogen ion concentration of the solution.
The end point can also be determined
electrochemically by either potentiometric or
conductimetric 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.
The behaviour of an acid type indicator is
described by the equilibrium;
Theory of Indicator Behaviour
In- + H3O+
HIn + H2O
The behaviour of an base type indicator
is described by the equilibrium;
In + H2O
InH+ + OH-
Theory of Indicator Behaviour
The equilibrium constant takes the form:
[H3O+][In-]
= Ka
[HIn]
Rearranging:
-]
[HIn
[H3O+] = Ka
[In-]
Theory of Indicator Behaviour
pH (acid colour) = -log(Ka . 10) = pKa +1
pH (base colour) = -log(Ka / 10) = pKa -1
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.
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 Titrations
Strong 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
reaction of the solute with water.
The
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
NaOH,
KOH, Na2CO3
Standards:anhydrous
boiling HCl.
Na2CO3 and constant
Neutralisation Titrations
Weak acids and bases
Examples
Ethanoic
acid
Sodium cyanide
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.
After
additions of the titrant the solution behaves as a
buffer. The pH of each buffer can be calculated from
there analytical concentrations.
At
the equivilence point only salt is present and the pH is
calculated from the concentration of this product.
Beyond
the equivilence point the pH is governed largely
by the concentration of the excess titrant.
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.
The most useful of these complexing agents
are organic compounds with several electron
donor groups that can form multiple
covalent bonds with metal ions.
Complexometric Titrations
Most metal ions react with electron-pair
donors to form coordination compounds or
complex ions.
The donor species, or LIGAND, must have
at least one pair of unshared electrons
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.
Typical values are 2, 4 and 6.
The species formed as a result of
coordination can be electrically positive,
neutral or negative.
Complexometric Titrations
Complexometric methods have been
around for more than a century.
Rapid expansion in the 1940’s based on
a class of coordination compounds
called CHELATES.
Complexometric Titrations
A chelate is produced when a metal ion
coordinates to two or more donor groups
within a single ligand.
For example the copper complex of glycine.
Complexometric Titrations
NH2
Cu2+ + 2 H
C
C
OH
H
O
C
O
O
O
C
Cu
C N
H2 H2
O
+ 2H +
N C
H2 H2
Complexometric Titrations
A ligand with a single donor group is
called unidentate.
Glycine is bidentate.
Tri, tetra, penta and hexadentate
chelating agents are also known.
Complexometric Titrations
Multidentate ligands have two advantages
over unidentate ligands.
They react more completely with cations to
provide a sharper endpoint.
The reaction is a single step process.
Complexometric Titrations
Tertiary amines that also contain carboxylic acid
groups form remarkably stable chelates with many
metal ions.
Ethylenediaminetetraacetic Acid EDTA
HOOCCH2
CH 2CO O H
N
HOOCCH2
CH2
CH2
N
CH 2CO O H
Complexometric Titrations
EDTA can complex a large number of metal ions.
Approximately 40 cations can be determined by
direct titration.
EDTA is usually used as the disodium salt,
Na2H2EDTA
H2EDTA2- + M2+ [M(EDTA)]2- + 2H+
Complexometric Titrations
Because EDTA complexes most cations, the
reagent might appear at first glance to be totally
lacking in selectivity.
However, great control can be acheived by pH
regulation and the selection of suitable indicators.
Complexometric Titrations
Indicators are generally complexing agents which
undergo a colour change when bonded to a metal
ion.
H2EDTA2- + [M(Ind)] [M(EDTA)]2- + Ind2- + 2H+
Complexometric Titrations
Typical indicators are:
Murexide
Solochrome
black
Calmagite
Bromopyrogallol
Xylenol
orange
red
Complexometric Titrations
Typical applications:
Determination
of cations
Hardness of water
Redox Titrations
Basics
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
Specific:
Starch
Oxidation
/ Reduction Indicators
Redox Titrations
Redox Titrations
Potassium Permanganate
MnO4- + 8H+ + 5e- Mn2+ + 4H2O
Standardisation
Sodium
oxalate or arsenic (III) oxide
Many Analyses
Redox Titrations
Hydrogen Peroxide:
2MnO4- + 5H2O2 + 6H+ 2Mn2+ + 5O2 + 8H2O
Nitrites:
2MnO4- + 5NO2- + 6H+ 2Mn2+ + 5NO3- + 3H2O
Redox Titrations
Persulphates:
Add
an excess of iron (II)
S2O82- + 2Fe2+ + 2H+ 2Fe3+ + 2HSO4 The
excess iron (II) is determined by back titration against
standardised permangenate.
MnO4- + 8H+ + 5Fe2+ Mn2+ + 5Fe2+ + 4H2O
Redox Titrations
Potassium Dichromate
CrO72- + 14H+ + 6e- 2Cr3+ + 7H2O
Standardisation
Against
metallic iron
1 mole K2CrO7 = 6 moles Fe
Redox Titrations
Iron (II):
CrO72- + 14H+ + 6Fe2+ 2Cr3+ + 6Fe3+ 7H2O
Indicators
include N-phenylanthranilic acid and sodium
diphenylamine 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- 2I-
Standardisation
Standardised
sodium thiosulphate or
arsenic (III) oxide
Many Analyses
Redox Titrations
Hydrogen Peroxide:
Thiosulphates:
H2O2 + 2I- + 2H+ I2 + 2H2O
2S2O32- + I2 S4O62- + 2I-
Hydroxyl Groups:
2OH- + I2 IO- + H2O + 2I-
Redox Titrations
Others:
Copper
Dissolved
oxygen
Chlorine
Arsenic
(V)
Sulphides
etc..........
Redox Titrations
Cerium (IV) Sulphate
Very strong oxidising agent (1.43V)
Ce4+ + e- Ce3+
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.
Titrimetric methods based upon the use
of silver nitrate are sometimes called
Argentometric titrations.
Precipitation Titrations
Used for the determination of many
anions 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.
The most accurate is the
VOLHARD METHOD.
Precipitation Titrations
VOLHARD METHOD
A back titration of thiocyanate ions
against the excess silver ions using an
iron (II) salt as the indicator.
Precipitation Titrations
Ag+ + SCN-
AgSCN
Fe3+ + SCN-
FeSCN2+
Blood Red