Acid–Base Titration.
Preparation and standardization of titrants.
1.
Titrimetric
method of the analysis: the basic concepts and classification.
2.
Technics of
titrimetric analysis.
3.
Types of
titrimetric determinations.
4.
Calculations
in titrimetric analysis
5. Ways of preparation of standard solutions
6. Establishment
of secondary standard concentration
7. Preparation rules of primary solutions and
definition of their titre
8. Requirements to reactions in titrimetric analysis.
9. Protolytometry: a theory,
reaction and titrants and determinate
substance.
Titrimetry, in which we measure the volume of a
reagent reacting stoichiometrically with the analyte, first appeared as an
analytical method in the early eighteenth century. Unlike gravimetry,
titrimetry initially did not receive wide acceptance as an analytical
technique. Many prominent late-nineteenth century analytical chemists preferred
gravimetry over titrimetry and few of the standard texts from that era include
titrimetric methods. By the early twentieth century, however, titrimetry began
to replace gravimetry as the most commonly used analytical method.
The acid-base
method of titration is very widely applied in the pharmaceutical analysis. Many
drugs are the bases or acids on the chemical nature and consequently their quantitative
definition is possible with the help acidimetry or alkalimetry.
Titrimetric
Methods of Analysis
Titrimetry, in which
we measure the volume of a reagent reacting stoichiometrically with the analyte,
first appeared as an analytical method in the early eighteenth century.
Unlike gravimetry, titrimetry initially did not
receive wide acceptance as an analytical technique. Many prominent late-nineteenth century
analytical chemists preferred gravimetry over titrimetry and few of the standard
texts from that era include titrimetric methods. By the early twentieth century, however, titrimetry began to
replace gravimetry as the most commonly used analytical method. Interestingly,
precipitation gravimetry developed in the absence of a theory of precipitation. The
relationship between the precipitates mass and the mass of analyte, called a gravimetric
factor, was determined experimentally by taking known masses of analyte (an external
standardization). Gravimetric factors could not be calculated using the precipitation
reactionÕs stoichiometry because chemical formulas and atomic weights
were not yet available! Unlike gravimetry, the growth and acceptance of titrimetry
required a deeper understanding of stoichiometry, thermodynamics, and chemical equilibria.
By the early twentieth century the accuracy and precision of titrimetric methods
were comparable to that of gravimetry, establishing titrimetry as an accepted analytical
technique.
Titrimetric
methods are classified into four groups based on the type of reaction involved. These groups are
acid–base titrations, in which an acidic or basic titrant reacts with an analyte that is a base or an
acid; complexometric titrations involving a metal–ligand complexation reaction;
redox titrations, where the titrant is an oxidizing or reducing agent; and precipitation
titrations, in which the analyte and titrant react to form a precipitate. Despite
the difference in chemistry, all titrations share several common features, providing the
focus for this section.
Based on the measurement of the amount of reagent that combined with an
analyte. Titrimetic methods are widely used for routine analysis because they
are rapid, convenient, accurate, and readily automated. There are such variants
of titrimetry:
1. Volumetric.
Volume of reagent solution required for a complete reaction.
2. Gravimetric.
Weight of reagent required for a complete reaction.
3. Coulometric.
Time/current required for complete oxidation or reducing of an analyte.
Requirements to chemical reaction used in
titrimetric methods of analysis:
1. Reaction
between reagent and analyte must be specific. Titrant can not react with
impurities or additions of analyte solution.
2. Reaction must
be stoichiometric.
3. Titrant must
react rapidly with the analyte so that the time required between additions of
reagent is minimised.
4. Titrant must
react more or less completely with the analyte so that satisfactory end points
are realised.
5. Undergo a
selective reaction with the analyte that can be described by simple balanced
equation. Equilibrium constant must have high value.
Titration
is a process in which a standard reagent (titrant) is added to a solution of an
analyte until the reaction between the analyte and reagent is judged to be
complete.
Titration
http://www.youtube.com/watch?v=9DkB82xLvNE
Titration can be:
1) direct
titration – titrant add to an analyte solution and react with
determined substrance;
2) back-titration –
is a process in which the excess of a standard solution used to react with an
analyte is determined by titration with a second standard solution.
Back-titrations are required when the reagent is slow or when the standard
solution lacks stability. For example:
CaCO3 + HCl = CaCl2
+ H2O + CO2
surplus (titrant 1)
HCl + NaOH
= NaCl + H2O
residue titrant
2
3) substitute-titration –
is a process in which a standard solution used to react with an additional
(substitute) substance, amount of which is equivalent an analyte amount.
Substitute-titrations are required when the analytes are unstable substance or
when is impossible to indicate the equivalent (end) point in direct reaction.
For example:
CrCl2 +
FeCl3 = CrCl3 + FeCl2
analyte
substitute
5FeCl2 + KMnO4 + HCl = 5FeCl3
+ KCl + MnCl2 + 4H2O
Equivalence
point is the point where sufficient titrant has been added
to be stoichiometrically equivalent to the amount of analyte. The equivalence
point of a titration is a theoretical point that can not be determined
experimentally, but can be determined experimentally the end point.
End
point is the point in a titration when a physical change
that is associated with the condition of chemical equivalence occurs.
We can estimate its position by observing some
physical changes with various indicating techniques:
a) without
any special means. The visible changes occur in titrated solution – change of
titrant or analyte colour, turbidity arise, precipitation formation;
b)
with internal indicator using. The
special chemical substances called indicators are added to the analyte
solution. Typical indicator changes include the appearance or disappearance of
a colour, a change in colour, or the appearance or disappearance of turbidity;
c) with
instruments. This instruments respond to certain properties of the solution
that change in a characteristic way during the titration.
The
difference in volume between the equivalence point and the end point is the
titration error.
A standard
solution (or titrant) is a
reagent of exactly known concentration that is used in a titrimetric analysis.
Standard solutions are the main participants in all titrimetric methods of
analysis. The titrant solutions must be of known composition and concentration.
Ideally, we would like to start with a primary standard material.
Primary standard
is an highly purified compound that serves as the reference materials for a
titrimetric method of analysis. Important requirements for a primary standard
are:
1. High purity.
2. Stability
toward air.
3. Absence of
hydrate water so that the composition of the solid does not change with
variations in relative humidity.
4. Ready
availability at modest cost.
5. Reasonable
solubility in the titration medium.
6. Reasonable
large molar mass so that the relative error associated with weighing the
standard is minimised.
A secondary standard is compound whose purity has been
established by chemical analysis and serves as the reference material to a
titrimetric method of analysis.
stopcock buret for standartization
Gay-Lussac burette Buret with bottle of standard solution
Mohr burette Buret
with rubber shutter
Microburet:
à)
Shilov air-powered Buret; á) stopcock buret
The concentration of the standard solutions can be established by two
basic methods:
1. Direct method
– a carefully weighed quantity of a primary standard is dissolved in a suitable
solvent and diluted to an exactly known volume in a volumetric flask. A made
solution is referred to as a primary standard solution (titrant).
Volumetric flask — for preparing
liquids with volumes of high precision. It is a flask with an approximately
pear-shaped body and a long neck with a circumferential fill line.
2. Standardisation
– concentration of a volumetric solution (titrant) is detrmined by using to
titrate
1) a weighed
quantity of a primary standard,
where:
CN
and V are concentration and volume of secondary standard solution
m
and Em are mass and equivalent weight of primary standard
2) “standard
titrimetric substance” (primary standard),
More
often in an ampoule contains 0,1 mol (0,1 equivalents) of substances, it is
necessary for preparation of 0,1 mol/L solution.
3) a measured
volume of another standard solution.
where:
CN2
and V2 are concentration and volume of secondary standard
solution
CN1
and V1 are concentration and volume of primary standard
solution
A titrant that is standardised against a
secondary standard or against another standard solution is referred to as a secondary
standard solution (titrant).
Equivalence Points and End Points
For a titration to be accurate we must
add a stoichiometrically equivalent amount of titrant to a solution containing the
analyte. We call this stoichiometric mixture the equivalence point. Unlike
precipitation gravimetry, where the precipitant is added in excess, determining the exact
volume of titrant needed to reach the equivalence point is essential. The product of the
equivalence point volume, Veq, and the titrant’s concentration, CT, gives the moles of titrant reacting
with the analyte.
Knowing the stoichiometry of the titration reaction(s), we can calculate
the moles of analyte.
Unfortunately, in most titrations we usually have no obvious indication
that the equivalence point has been reached. Instead, we stop adding titrant
when we reach an end point of our choosing. Often this end point is indicated by
a change in the color of a substance added to the solution containing the analyte.
Such substances are known as indicators. The difference between the end point
volume and the equivalence point volume is a determinate method error, often called
the titration error. If the end point and equivalence point volumes
coincide closely, then the titration error is insignificant and can be safely
ignored. Clearly, selecting an appropriate end point is critical if a titrimetric
method is to give accurate results.
Units of concentration of standard
solutions
The concentration of standard solutions (titrants) are generally
expressed in units of either molarity (CM, or M) or normality (CN,
or N).
Molarity (M) – is the number of moles of a material per
liter of solution.
Normality (N) – is the number of species equivalents per
liter of solution.
Sometime is used also one unite of concentration – titer (T).
Titer established the relationship between volume of titrant and amount of
analysed substance present. The most commonly titer is in units of mg analysed
substance per ml of titrant. This system was developed to assist in doing
routine calculations. It reduces the amount of time and training for
technicians.
Equivalents law
Titrimetry is based on equivalents
law:
Na·Va
= Ns·Vs,
or
number of analyte equivalent present = number of standard reagent added,
or
one equivalent of one material will react exactly with one equivalent of
another
The weight of one equivalent of a compound depends on reference to a
chemical reaction in which that compound is a participant. Similarly, the
normality of a solution can never be specified without knowledge about how the
solution will be used. Equivalent value is based on the type of reaction and
the reactants:
1. One
equivalent weight of a substance participating in a neutralisation reaction is
that amount of substance that either react with or supplied one mol of hydrogen
ions in that reaction.
2. One
equivalent weight of a participant in an oxidation-reduction reaction is that
amount that directly or indirectly produces or consumer one mol of electrons.
3. The
equivalent weight of a participant in a precipitation or a complex-formation
reaction is that weight which or provides one mole of the univalent reacting
cation.
Volume as a Signal
Almost any chemical reaction can serve as a titrimetric method provided
that three conditions are met. The first condition is that all reactions
involving the titrant and analyte must be of known stoichiometry. If this is not the
case, then the moles of titrant used in reaching the end point cannot tell us how
much analyte is in our sample. Second, the titration reaction must occur rapidly. If
we add titrant at a rate that is faster than the reaction’s rate, then the end
point will exceed the equivalence point by a significant amount. Finally, a suitable
method must be available for determining the end point with an acceptable level
of accuracy. These are significant limitations and, for this reason, several
titration strategies are commonly used. A simple example of a titration is an analysis for Ag+
using thiocyanate, SCN–, as a titrant.
This reaction occurs quickly and is of known stoichiometry. A titrant of
SCN– is easily prepared using KSCN. To indicate the titration’s end point we add
a small amount of Fe3+ to the solution containing the analyte. The
formation of the redcolored Fe(SCN)2+ complex signals the end point.
This is an example of a direct titration since the titrant reacts with the analyte.
If the titration reaction is too slow, a suitable indicator is not
available, or there is no useful direct titration reaction, then an indirect analysis may be
possible. Suppose you wish to determine the concentration of formaldehyde, H2CO,
in an aqueous solution. The oxidation
of H2CO by I3–
is a useful reaction, except that it is too slow for a direct titration.
If we add a known amount of I3–, such that it is in excess, we can
allow the reaction to go to completion.
The I3– remaining can then be titrated with
thiosulfate, S2O32–.
This type of titration is called a back titration.
Calcium ion plays an important role in many aqueous environmental
systems. A useful direct analysis takes advantage of its reaction with the ligand
ethylenediaminetetraacetic acid (EDTA), which we will represent as Y4–.
Unfortunately, it often happens that there is no suitable indicator for
this direct titration. Reacting
Ca2+ with an excess of the Mg2+–EDTA complex
releases an equivalent amount of Mg2+. Titrating the released
Mg2+ with EDTA
gives a suitable end point. The amount of Mg2+ titrated provides an
indirect measure of the amount of Ca2+ in the original sample. Since the analyte
displaces a species that is then titrated, we call this a displacement
titration.
When a suitable reaction involving the analyte does not exist it may be
possible to generate a species that is easily titrated. For example, the sulfur
content of coal can be determined by using a combustion reaction to convert sulfur to sulfur
dioxide.
Passing the SO2 through an aqueous solution of hydrogen
peroxide, H2O2,
produces sulfuric acid, which we can titrate with NaOH,
providing an indirect determination of sulfur.
Titrimetric
Methods of Analysis are based on the measurement of the amount of reagent that
combined with an analyte. Titrimetic methods are widely used for routine
analysis because they are rapid, convenient, accurate, and readily automated.
There are such variants of titrimetry:
4. Volumetric.
Volume of reagent solution required for a complete reaction.
5. Gravimetric.
Weight of reagent required for a complete reaction.
6. Coulometric.
Time/current required for complete oxidation or reducing of an analyte.
Titrimetric
methods are classified into four groups based on the type of reaction involved.
These groups are
– acid–base titrations, in which an acidic or
basic titrant reacts with
an analyte that is a base or an acid;
– complexometric titrations involving a
metal–ligand complexation reaction;
– redox titrations, where the titrant is an
oxidizing or reducing agent;
– precipitation titrations, in which the analyte
and titrant react to form a precipitate.
Classification of titrimetric analysis methods
|
Method |
Technique |
Titrant |
|
Neutralisation (acid-basic
titration) |
Alkalimetry Acidimetry Halometry |
MeOH HAn HAn, MeOH |
|
Nonaqueous
titration |
HClO4
in acetic acid or nitrometane NaOH or CH3ONa
in methanol |
|
|
Redoximety (reducing-oxidising) |
Permanganatometry Iodometry Bromatometry Cerimetry Vanadatometry Titanometry Nitritimetry |
KMnO4 I2, Na2S2O3 KBrO3 Ce(SO4)2 NH4VO3 Ti2(SO4)3 NaNO2 |
|
Precipitation
titration |
Argentometry Mercurometry Rhodanometry |
AgNO3 Hg2(NO3)2 KSCN |
|
Complexometry (complex
compounds formation) |
Mercurimetry Fluorimetry Complexonometry |
Hg(NO3)2 NaF EDTA |
Requirements
to chemical reaction used in titrimetric methods of analysis:
6. Reaction
between reagent and analyte must be specific. Titrant can not react with
impurities or additions of analyte solution.
7. Reaction must
be stoichiometric.
8. Titrant must
react rapidly with the analyte so that the time required between additions of
reagent is minimised.
9. Titrant must
react more or less completely with the analyte so that satisfactory end points
are realised.
10. Undergo a
selective reaction with the analyte that can be described by simple balanced
equation. Equilibrium constant must have high value.
Titration
is a process in which a standard reagent (titrant) is added to a solution of an
analyte until the reaction between the analyte and reagent is judged to be
complete. Titration can be:
4) direct
titration – titrant add to an analyte solution and react with
determined substrance;
A simple example
of a titration is an analysis for Ag+ using thiocyanate, SCN–,
as a titrant.
AgNO3
+ KSCN = AgSCN¯ +
KNO3
This reaction
occurs quickly and is of known stoichiometry. A titrant of SCN– is
easily prepared using KSCN. To indicate the titration’s end point we add a
small amount of Fe3+ to the solution containing the analyte. The
formation of the redcolored Fe(SCN)2+ complex signals the end point.
5) back-titration –
is a process in which the excess of a standard solution used to react with an
analyte is determined by titration with a second standard solution.
Back-titrations are required when the reagent is slow or when the standard
solution lacks stability. For example: determination the concentration of
formaldehyde, H2CO, in an aqueous solution. The oxidation of H2CO by
I–
H2CO + 3NaOH + I2
= HCOONa + 2NaI + 2H2O
is a useful
reaction, except that it is too slow for a direct titration. If we add a known
amount of I2, such that it is in excess, we can allow the reaction
to go to completion. The I2 remaining can then be titrated with
thiosulfate, S2O32–.
I2
+ 2Na2S2O3 = Na2S4O6
+ 2NaI
6) substitute-titration
(displacement titration) – is a process in which a standard
solution used to react with an additional (substitute) substance, amount of
which is equivalent an analyte amount. Substitute-titrations (displacement
titration) are required when the analytes are unstable substance or when is
impossible to indicate the equivalent (end) point in direct reaction. For
example: calcium ion plays an important role in many aqueous environmental
systems. A useful direct analysis takes advantage of its reaction with the
ligand ethylenediaminetetraacetic acid (EDTA), which we will represent as Y4–.
Ca2+
+ Y4– «
CaY2–
Unfortunately, it often happens that there is no
suitable indicator for this direct titration. Reacting Ca2+ with an
excess of the Mg2+–EDTA complex
Ca2+
+ MgY2– «
CaY2– + Mg2+
releases an equivalent amount of Mg2+.
Titrating the released Mg2+ with EDTA
Mg2+
+ Y4– «
MgY2–
gives a suitable end point. The amount of Mg2+
titrated provides an indirect measure of the amount of Ca2+ in the
original sample.
When a
suitable reaction involving the analyte does not exist it may be possible to
generate a species that is easily titrated. For example, the sulfur content of
coal can be determined by using a combustion reaction to convert sulfur to
sulfur dioxide.
S(s) +
O2(g) ® SO2(g)
Passing the SO2 through an aqueous solution
of hydrogen peroxide, H2O2,
SO2(g)
+ H2O2 ® H2SO4
produces sulfuric acid, which we can titrate with NaOH,
H2SO4
+ 2NaOH ®
Na2SO4 + 2H2O
providing an indirect determination of sulfur.
Equivalence point is the point where sufficient titrant
has been added to be stoichiometrically equivalent to the amount of analyte.
The equivalence point of a titration is a theoretical point that can not be
determined experimentally, but can be determined experimentally the end point.
The product of the equivalence point volume, Veq, and the titrant’s concentration, CT, gives the moles of titrant reacting with the
analyte.
Moles titrant
= Veq ´ CT
Knowing the stoichiometry of the titration
reaction(s), we can calculate the moles of analyte.
Unfortunately,
in most titrations we usually have no obvious indication that the equivalence
point has been reached. Instead, we stop adding titrant when we reach an end
point of our choosing. End point
is the point in a titration when a physical change that is associated with the
condition of chemical equivalence occurs. We can estimate its position by
observing some physical changes with various indicating techniques:
d)
without any special means. The
visible changes occur in titrated solution – change of titrant or analyte
colour, turbidity arise, precipitation formation;
e) with
internal indicator using. The special chemical substances called indicators
are added to the analyte solution. Typical indicator changes include the
appearance or disappearance of a colour, a change in colour, or the appearance
or disappearance of turbidity;
f) with
instruments. This instruments respond to certain properties of the solution
that change in a characteristic way during the titration.
The difference between the end point volume and
the equivalence point volume is a determinate method error, often called the titration
error. If the end point and equivalence point volumes coincide closely,
then the titration error is insignificant and can be safely ignored. Clearly,
selecting an appropriate end point is critical if a titrimetric method is to
give accurate results.
Titration
Curves
To find the end point we monitor some property of
the titration reaction that has a well-defined value at the equivalence point.
For example, the equivalence point for a titration of HCl with NaOH occurs at a
pH of 7.0. We can find the end point, therefore, by monitoring the pH with a pH
electrode or by adding an indicator that changes color at a pH of 7.0.
A titration
curve provides us with a visual picture of how a property, such as pH,
changes as we add titrant (Figure 9.1). We can measure this titration curve
experimentally by suspending a pH electrode in the solution containing the
analyte, monitoring the pH as titrant is added. As we will see later, we can
also calculate the expected titration curve by considering the reactions
responsible for the change in pH. However we arrive at the titration curve, we
may use it to evaluate an indicator’s likely titration error. For example, the
titration curve in Figure 9.1 shows us that an end point pH of 6.8 produces a
small titration error. Stopping the titration at an end point pH of 11.6, on the
other hand, gives an unacceptably large titration error.
Figure 9.1. Acid–base
titration curve for 25.0 mL of 0.100 M HCl with 0.100 M NaOH.
The titration
curve in Figure 9.1 is not unique to an acid–base titration. Any titration
curve that follows the change in concentration of a species in the titration
reaction (plotted logarithmically) as a function of the volume of titrant has
the same general sigmoidal shape.
Dependence of inflection points on concentration of defined substance
(0,1 mol/L and 0,01 mol/L)
Dependence of inflection points on force of acid
Dependence of inflection points on force of acid
Titration
curve for weak acid
(CH3COOH) by weak base (NH4OH)
Titration
curve for H3PO4 by NaOH
Concentration
is not the only property that may be used to construct a titration curve. Other
parameters, such as temperature or the absorbance of light, may be used if they
show a significant change in value at the equivalence point. Many titration
reactions, for example, are exothermic. As the titrant and analyte react, the
temperature of the system steadily increases. Once the titration is complete,
further additions of titrant do not produce as exothermic a response, and the
change in temperature levels off. The titration curve contains two linear
segments, the intersection of which marks the equivalence point.
A standard solution (or titrant) is a reagent of exactly known
concentration that is used in a titrimetric analysis. Standard solutions are
the main participants in all titrimetric methods of analysis. The titrant
solutions must be of known composition and concentration. Ideally, we would
like to start with a primary standard material.
Primary standard
is an highly purified compound that serves as the reference materials for a
titrimetric method of analysis. Important requirements for a primary standard
are:
1. High purity.
2. Stability
toward air.
3. Absence of hydrate
water so that the composition of the solid does not change with variations in
relative humidity.
4. Ready
availability at modest cost.
5. Reasonable
solubility in the titration medium.
6. Reasonable
large molar mass so that the relative error associated with weighing the
standard is minimised.
A secondary standard is compound whose purity has been
established by chemical analysis and serves as the reference material to a
titrimetric method of analysis.
The concentration of the standard solutions can be established by two
basic methods:
3. Direct method
– a carefully weighed quantity of a primary standard is dissolved in a suitable
solvent and diluted to an exactly known volume in a volumetric flask. A made
solution is referred to as a primary standard solution (titrant).
4. Standardisation
– concentration of a volumetric solution (titrant) is detrmined by using to
titrate
4) a weighed
quantity of a primary standard,
5) a weighed
quantity of a secondary standard,
6) a measured
volume of another standard solution.
A titrant that is standardised against a secondary standard or against
another standard solution is referred to as a secondary standard solution
(titrant).
Units of concentration of standard
solutions
The concentration of standard solutions (titrants) are generally expressed
in units of either molarity (CM, or M) or normality (CN,
or N).
Molarity (M) – is the number of moles of a material per
liter of solution.
Normality (N) – is the number of species equivalents per
liter of solution.
Sometime is used also one unite of concentration – titer (T).
Titer established the relationship between volume of titrant and amount of
analysed substance present. The most commonly titer is in units of mg analysed
substance per ml of titrant. This system was developed to assist in doing
routine calculations. It reduces the amount of time and training for
technicians.
Equivalents law
Titrimetry is based on equivalents
law:
Na·Va
= Ns·Vs,
or number of analyte
equivalent present = number of standard reagent added,
or one equivalent of one material
will react exactly with one equivalent of another
The weight of one equivalent of a compound depends on reference to a
chemical reaction in which that compound is a participant. Similarly, the
normality of a solution can never be specified without knowledge about how the
solution will be used. Equivalent value is based on the type of reaction and
the reactants:
4. One
equivalent weight of a substance participating in a neutralisation reaction is
that amount of substance that either react with or supplied one mol of hydrogen
ions in that reaction.
5. One
equivalent weight of a participant in an oxidation-reduction reaction is that
amount that directly or indirectly produces or consumer one mol of electrons.
6. The
equivalent weight of a participant in a precipitation or a complex-formation
reaction is that weight which or provides one mole of the univalent reacting
cation.
Calculations in titrimetric method of analysis
|
T
= |
N
= |
m
= |
|
mx(is)
= |
mx(al)
= |
ax
= |
|
T – titer
(g/ml); |
N – normality (number of equivalents/l); |
|
Nt
– nomality of used titrant (N); |
Vt – volume of used titrant (ml); |
|
m – mass of
substance (g); |
meq – mass of one equivalent (g); |
|
mx(al) – amount of analyte, determined as aliquot of sample
(g); |
mx(is) – amount of analyte, determined as individual sample
(g); |
|
meqx
– mass of one equivalent of analyte (g); |
W – dilution of analyte sample (ml); |
|
Vs
– aliquot of sample solution (ml); |
px – mass of sample (g); |
|
ax
– percentage of substance in sample (%) |
|
Indicators of Titrimetry Methods
Indicators are the chemical compounds,
which give some external effect
attached to concentrations of reactive species according to equivalence point. This external effect can be accompanied by change, appearance or
disappearance of colouring, and formation of
slightly soluble compounds
(precipitate formation).
On
appliance technique indicators are external
and internal.
Internal indicators are introduced into titrated solution. An end point install on changes of colour of analysed
mixture.
The external indicators
are used when internal indicators using is
impossible. Reaction with external indicators runs out of analysed mixture.
Some drops of analysed solution put on peace of filter paper, impregnated with
indicator, or mix with drop of indicator solution on porcelain plate.
For effect the reactions appearance indicators are reversible and
unreversible.
Reversible indicators – changes the colour can be repeated
many times as changes the system state.
Unreversible indicators – colour
changes ones with destruction of indicator molecule. The unreversible
indicators are less comfortable and
thinly use.
Acid-Basic
Titration
The earliest acid–base titrations involved the determination of
the acidity or alkalinity of solutions, and the purity of carbonates and
alkaline earth oxides. Before 1800, acid–base titrations were conducted using H2SO4,
HCl, and HNO3 as acidic titrants, and K2CO3 and
Na2CO3 as basic titrants. End points were determined
using visual indicators such as litmus, which is red in acidic solutions and
blue in basic solutions, or by observing the cessation of CO2
effervescence when neutralizing CO32–. The accuracy of an
acid–base titration was limited by the usefulness of the indicator and by the
lack of a strong base titrant for the analysis of weak acids.
Titration of an acid solution of unknown concentration
with a base solution of known concentration
http://www.onlinemathlearning.com/acid-base-titration.html
http://library.thinkquest.org/3310/lographics/video/dtitrate.html
The utility of acid–base titrimetry improved when NaOH was first
introduced as a strong base titrant in 1846. In addition, progress in
synthesizing organic dyes led to the development of many new indicators.
Phenolphthalein was first synthesized by Bayer in 1871 and used as a visual
indicator for acid–base titrations in 1877. Other indicators, such as methyl
orange, soon followed. Despite the increasing availability of indicators, the
absence of a theory of acid–base reactivity made selecting a proper indicator
difficult.
Developments in equilibrium theory in the late nineteenth century led to
significant improvements in the theoretical understanding of acid–base
chemistry and, in turn, of acid–base titrimetry. Sørenson’s
establishment of the pH scale in 1909 provided a rigorous means for comparing
visual indicators. The determination of acid–base dissociation constants made
the calculation of theoretical titration curves possible, as outlined by
Bjerrum in 1914. For the first time a rational method existed for selecting
visual indicators, establishing acid–base titrimetry as a useful alternative to
gravimetry.
This is a quick and accurate method for determining acidic or basic
substances in many samples. This method enable to determine some inorganic and hundred of organic acids and bases of
different types; frequently organic compounds
are titrated in waterless environment. The used titrant is typically a strong
acid or base. The sample species can be either a strong or weak acid or base.
The neutralisation method based on acid-basic reactions (exchange reactions by
protons), which one can be expressed by general scheme:
HA + BOH = B+ + A– + H2O
Titrations according to the applied
titrant are
1) acidimetric (titrants are the acids solutions) – uses for determination of
strong and weak bases, salts of strong bases and weak acids and organic compounds;
2) alkalimetric (titrants are solutions of bases) – uses for titration of strong and weak acids, sour
salts, salts of strong acids and weak bases, organic compounds having acidic disposition (acids, phenols).
Standard Titrants
I.
Bases.
NaOH is the most common although KOH can be serve
the same purpose. There are not primary standards.
Primary standards for bases
standardisation are weak acids:
oxalic acid H2C2O4∙2H2O,
benzoic acid C6H5COOH, succinic acid HOOC(CH2)2COOH,
potassium hydrogen phthalate KHC8H4O4,
potassium hydrogen iodate KH(IO3)2, potassium hydrogen
tartrate KHC4H4O6.
II. Acids.
More frequently are used HCl and H2SO4.
There are not primary standards too. Primary standards for acids
standardisation are weak bases:
borax Na2B4O7´10H20,
TRIS (hydroxymethyl-aminomethane) (HOCH2)2CNH2,
sodium carbonate Na2CO3, mercury oxide HgO, sodium
oxalate Na2C2O4, potassium iodate KIO3.
Standardization of ÍÑl solution on
sodium tetraborate.
§
Weigh exact shot of Na2B4O7×5H2O or
Na2B4O7×10H2O and
place it in a measured flask, dissolve in hot
water, after a solution is cooled and diluted of solution by water to necessary
volume and it is mixed.
§
In a flask for titration place an aliquot
of prepared primary standard solution Na2B4O7×5H2O or
Na2B4O7×10H2O,
add some drops of the methyl orange. The received solution is titrated
by solution of ÍÑl to
change of colour with yellow to orange with a rose shade.
§
By
3-4 results of titration calculate average volume of used titrant and
calculate concentration of hydrochloric acid.
Standardization of HCl
solution on sodium carbonate
§
In a flask for titration place exact shot
of sodium carbonate, dissolve in necessary volume of
water, add some drops methyl orange and titrate this
solution by chloric acid.
§
Such titration repeat for 3-4 times. Each time
calculate concentration of HCl:
§
By 3-4 results of titration calculate
average concenration of chloric acid.
Standardization of NaOH
solution on oxalic acid.
§
Weigh exact shot of H2C2O4×2H2O and
place it in a measured flask, dissolve in hot
water, after a solution is cooled and diluted of solution by water to necessary
volume and it is mixed.
§
In a flask for titration place an aliquot
of prepared primary standard solution H2C2O4×2H2O ,
add some drops of the phenolphthalein.
The received solution is titrated by solution of NaOH to
change of colourless to rose (or red).
§
By
3-4 results of titration calculate average volume of used titrant and
calculate concentration of NaOH.
Titration Curves
For an acid–base titration, the equivalence point is characterized by a
pH level that is a function of the acid–base strengths and concentrations of
the analyte and titrant. The pH at the end point, however, may or may not
correspond to the pH at the equivalence point. To understand the relationship
between end points and equivalence points we must know how the pH changes
during a titration.
Acid-base property of titration system changes accordingly to proportion
(ratio) of protolytes in mixture. Dependency on correlation of protolytes force
the equivalence point can be in neutral, alkaline or acidic environment. Change
the pH value during titration process, or dependence the pH value on
concentration of titrated electrolytes, show the titration curves. A titration
curve is a graph of the pH as a function of the amount of titrant (acid or
base) added. This still results in four types of titration for simple acids or
bases:
Strong acid vs. strong base
Strong
acid vs. weak base
Strong
base acid vs. strong acid
Strong
base vs. weak base
Strong
Acid-Strong Base Titrations
Here is an example of a titration curve, produced when a strong base is
added to a strong acid. This curve show how pH varies as 0,100 M NaOH is added
to 50,0 ml of 0,100 M HCl.
For the reaction of a strong base with a strong acid the only
equilibrium reaction of importance is
NaOH
+ HCl = NaCl + H2O
or, in ionic form
H+
+ OH– = H2O (1)
The
first task in constructing the titration curve is to calculate the volume of
NaOH needed to reach the equivalence point. At the equivalence point we know
from reaction that
Moles
HCl = moles NaOH
or
MaVa =
MbVb
where the subscript ‘a’
indicates the acid, HCl, and the subscript ‘b’ indicates the base, NaOH. The
volume of NaOH needed to reach the equivalence point, therefore, is
The equivalence point of the titration is the point at which exactly
enough titrant has been added to react with all of the substance being titrated
with no titrant left over. In other words, at the equivalence point, the number
of moles of titrant added so far corresponds exactly to the number of moles of
substance being titrated according to the reaction stoichiometry. (In an
acid-base titration, there is a 1:1 acid:base stoichiometry, so the equivalence
point is the point where the moles of titrant added equals the moles of
substance initially in the solution being titrated.)
Notice that the pH increases slowly at first, then rapidly as it nears
the equivalence point. At the equivalence point, the pH = 7.00 for strong
acid-strong base titrations. However, in other types of titrations, this is not
the case.
Titrations
Involving a Weak Acid
There are three major differences between this curve (in blue) and the
one we saw before (in black):
1.
The weak-acid solution has a higher
initial pH.
2.
The pH rises more rapidly at the
start, but less rapidly near the equivalence point.
3.
The pH at the equivalence point does
not equal 7.00. The equivalence point for a weak acid-strong base titration has
a pH > 7.00.
For a strong acid-weak base or weak
acid-strong base titration, the pH will change rapidly at the very beginning
and then have a gradual slope until near the equivalence point. The gradual
slope results from a buffer solution being produced by the addition of the
strong acid or base, which resists rapid change in pH until the added acid or
base exceeds the buffer's capacity and the rapid pH change occurs near the
equivalence point.
Titration
curve of a weak acid being titrated by a strong base:
Here,
0,100 M NaOH is being added to 50,0 ml of 0,100 M acetic acid.
Titrations
Involving a Weak Base
Titration curve of a weak base being titrated by a strong acid:
Here, 0,100 M HCl is being added to 50,0 ml of 0,100 M ammonia solution.
As in the weak acid-strong base titration, there are three major
differences between this curve (in blue) and a strong base-strong acid one (in
black): (Note that the strong base-strong acid titration curve is identical to
the strong acid-strong base titration, but flipped vertically.)
The weak-acid solution has a lower initial pH.
1.
The pH drops more rapidly at the
start, but less rapidly near the equivalence point.
2.
The pH at the equivalence point does
not equal 7.00. The equivalence point for a weak base-strong acid titration has
a pH < 7.00.
Titrations of
Polyprotic Acids
An
example of a polyprotic acid is H2CO3 which neutralises
in two steps:
H2CO3 + OH– « H2O
+ HCO3–
HCO3– + OH– « H2O
+ CO32–
The titration curve for these reactions will look like this, with two
equivalence points.
Uses of Titrations
Use titration data or a titration curve to calculate reaction quantities
such as the concentration of the substance being titrated.
The most common use of titrations is for measuring unknown
concentrations. This is done by titrating a known volume of the unknown
solution with a solution of known concentration (where the two react in a
predictable manner) and finding the volume of titrant needed to reach the
equivalence point using some method appropriate to the particular reaction.
Then, the volume and concentration of titrant can be used to calculate the
moles of titrant added, which, when used with the reaction stoichiometry, gives
the number of moles of substance being titrated. Finally, this quantity, along
with the volume of substance being titrated, gives the unknown concentration.
For acid-base titrations, the
equivalence point can be found very easily. A pH meter is simply placed in the
solution being titrated and the pH is measured after various volumes of titrant
have been added to produce a titration curve. The equivalence point can then be
read off the curve.
In the same way, knowing the equivalence point can also be used to
calculate other unknown quantities of interest in acid base reactions, such as
concentration of titrant or volume of solution being titrated, provided that
enough other information is known to perform the calculations.
http://freevideolectures.com/Course/2550/Chemistry/50
Nonaqueous Titrations
Nonaqueous titration is the special technique of
the acid-base titration. Acids with dissociation constant value less than 1×10–7
(pKa > 7) and bases with dissociation constant value less than 1×10–7
(pKb > 7) can not be titrated in water solutions. The ionisation
degree of such weak acid and weak bases is comparable with indicator ionisation
degree. During titration of aqueous solutions of these compounds can not be
indicated equivalent point.
For titration of weak acid and weak bases are
used waterless (nonaqueous) solvents, which intensify its acidic/basic
properties. Indeed, water is the most common solvent in acid–base titrimetry.
When considering the utility of a titration, however, the solvent’s influence
cannot be ignored.
Autoprotolysis of solvents
Many solvents autoprotolysise
like as water:
2C2H5OH
« C2H5OH2+
+ C2H5O–
2H2N–(CH2)2–NH2
«
R–NH3+ + R–NH–
2(CH3)2SO
« (CH3)2SOH+
+ CH3–SO–CH2–
The
dissociation, or autoprotolysis constant for a solvent, SH, relates the
concentration of the protonated solvent, SH2+, to that of
the deprotonated solvent, S–. For amphoteric solvents, which can act
as both proton donors and proton acceptors, the autoprotolysis reaction is
2SH
«
SH2+ + S–
with an equilibrium constant
of
Ks = [SH2+][S–]
Remember, that water
autoprolysis constant KS = KW = 1×10–14.
You
should recognize that Kw
is just the specific form of Ks
for water. The pH of a solution is now seen to be a general statement about the
relative abundance of protonated solvent
pH
= –log[SH2+]
where the pH of a neutral
solvent is given as
Perhaps the most obvious limitation imposed by Ks is the change in pH during a
titration. To see why this is so, let’s consider the titration of a 50 mL
solution of 10–4 M strong acid with equimolar strong base. Before
the equivalence point, the pH is determined by the untitrated strong acid,
whereas after the equivalence point the concentration of excess strong base
determines the pH. In an aqueous solution the concentration of H+
when the titration is 90% complete is
= 5.3´10–6
M
corresponding
to a pH of 5.3. When the titration is 110% complete, the concentration of OH–
is
or a pOH of 5.3. The pH,
therefore, is
pH
= pKw – pOH = 14.0 –
5.3 = 8.7
The
change in pH when the titration passes from 90% to 110% completion is
DpH
= 8.7 – 5.3 = 3.4
If
the same titration is carried out in a nonaqueous solvent with a Ks of 1.0´10–20,
the pH when the titration is 90% complete is still 5.3. However, the pH when
the titration is 110% complete is now
pH
= pKs – pOH = 20.0 –
5.3 = 14.7
In this case the change in pH
of
DpH
= 14.7 – 5.3 = 9.4
is
significantly greater than that obtained when the titration is carried out in
water.
