Lab Final
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Experiment 6: precipitation titrations
In this experiment the main aim was to determine the chloride ion in an unknown solution by
using Mohr and Fajans method.
Mohr 's Method
· Add
at a time till bubbles are gone
· Titrations must be carried at a PH 7-8. If the
too low (pH<6)part of the chromate ion will change into
dichromate:
in this experiment, silver nitrate
contact with skin must be
avoided.If skin gets exposed to
silver nitrate solution it will turn
black and might peel of.
Essentially due to a
photographic reaction
In addition, more silver nitrate solution would be required to
reach the end point in the case of low PH.
.After adding.
, potassium chromate titrate it with standard silver nitrate solution till
color changes to pale brown
. Color change is due to the formation of silver chromate,
. Solubility of Silver chromate increases with rising temperature thus its sensitivity as an
indicator decreases. Hence Mohr'S method requires room temperature.
. Determine an indicator blank by using chloride free
-
Summary of Mohr’s Method
→ analyte: chloride
→ titrant: silver nitrate
→ indicator: sodium chromate
→ room temperature
→ PH between 7 or 8
→ precipitate:
-
In this titration, the anionic form of adsorption indicator dichlorofluorescein is used to detect
the end point. With the first excess of the titrant, indicator forms a complex with the silver
ions at the counter- ion layer and imparts a red color to the precipitate. The particles of the
precipitate should be kept in colloidal state in order to obtain satisfactory color change.
Therefore, dextrin is added to stabilize the silver chloride (AgCI) particles against coagulation.
Titrate this solution with standard AgNOs solution till the color of the solution turns into a
permanent pink color.
Before the end-point, the surface of the silver chloride particles are negatively charged due
to the adsorption of excess chloride ions present in the solution. Therefore the indicator anions
are repelled by this negatively charged surfaces and the color of the solution appears as
yellow- green. After all chloride ions present in the solution are titrated by silver ions, this
time, excess silver ions are adsorbed on the silver chloride particles immediately. A positively
charged layer is formed and it attracts the dichlorofluoresceinate ion displaying a pink-red
color.
Summary of Fajans Method
→ analyte: chloride
→ titrant: silver nitrate
→ adsorption indicator
→ indicator: dichlorofluoresein
→ room temperature
→dextrin is used to stabilize
AgCl particles against
coagulation
PRE -LAB STUDIES
1) What are the main differences between precipitation titrations and neutralization
titrations?
Neutralization is a reaction used to neutralize acids or bases and precipitation is a reaction in
which the new formed compound is insoluble in the solution.
2) What are the applications of precipitation titration?
Precipitation reactions can be used for making pigments, removing salts from water in water
treatment, and for qualitative chemical analysis.
3) Explain Mohr method by writing related reaction equations.
4) What causes the dark color that forms when AgCl is exposed to light?
6) Explain briefly the working principle of adsorption indicators? Give examples.
The precipitate AgCl adsorbs the Cl– ion when a chloride solution is titrated with a
solution of AgNO3 (a precipitate tends to adsorb its ions). This is known as the primary
adsorption layer, and it is retained by the secondary adsorption of oppositely charged
ions in the solution. When the equivalence point is achieved, there is an excess of Ag+
ions, which displaces Cl– ions from the primary layer and holds nitrate ions in the
secondary adsorption layer.
After equivalence with
the adsorption of chloride
particles positive surface
formed , which attracted
negatively charged
indicator. Hence pink-red
color formed.
POST-LAB STUDIES
1) Explain the importance of adjusting the pH of the titration medium in Mohr method.
Write the related reaction equations.
Titrations must be carried at a PH 7-8. If the too low (pH<6)part of the chromate ion will
change into dichromate.
2) Blank is used to dissipate the error brought about by impurities. What is the other
reason for using blank in Mohr method?
To eliminate the error caused by the indicator.
3) In Mohr method, why did we use CaCOs in the blank?
To observe color change.
4) Explain the working principle of CrO,2 and write the importance of the concentration
of it in the titration medium.
Determination
Experiment 4: Gravimetric
Gravimetric Determination
Solution
of Nickel in an Unknown
Unknown Solution
Nickel (11) forms a precipitate with an alcoholic solution of dimethylgloxime (dmg),
In a slightly alkaline medium. The formation of the red solid Ni complex is illustrated below:
-
o
O
→ This reaction must be performed in a buffered
solution to keep pH >5 by using either ammonia or citrate
solution. Otherwise the equilibrium above favors the
dissolution of Ni complex back into mother liquor.
t
8
Hero
-
→ The bulky character of Ni complex limits the mass that can be analyzed efficiently and thus
the sample mass. Since the dimethylglyoxime is only slightly soluble in water (0.063 g in 100 ml
at 25°C ), an alcoholic solution of it used as the precipitating agent. However it is crucial to
avoid addition of excess amount of DMG as it may crystallize out with Ni complex.it is also
important that the Ni complex itself is also slightly soluble in alcoholic solutions. Thus by adding
precise amount of DMG, errors can be minimized.
PROCEDURE
PROCEDURE
• Always store crucibles in a desiccator.
• make the solution basic with ammonia.
• after make the solution slightly acidic with HCl. Slightly acidic means PH is 4-5.in this step
some minor Ni can be lost.
• Heat the solution to 80 - 60°C and then add alcoholic DMG solution.
• make solution basic again with ammonia.
• digest the sample at 60°C for an hour.
• perform chloride test with silver nitrate, AgNO3. If silver color precipitate forms, keep
washing the precipitate.
• Dry the precipitate at ~110°C
• do not heat the precipitate to temperatures over 130°C (Ni complex may decompose)
PRE-LAB STUDIES
1- Define coprecipitation and explain coprecipitation errors.
Coprecipitation is a process in which normally soluble compounds carried out of solution by a
precipitate. Coprecipitation can cause both + or - errors if contaminant is not the ion to be
determined, a positive error will occur.
2- What are the differences between Gooch crucibles and porcelain ones?
3-What are the differences between colloidal and crystalline suspension? Which factors
determine the particles size of precipitates? Explain briefly.
A colloidal precipitate consists of solid
particles with dimension less then 10
^-4 cm.
Colloidal precipitates remain suspended
in solution unless caused to
agglomerate.
Particle size
A crystalline precipitate consists of
solid particles with dimension greater
than 10 -4 cm.
Crystalline precipitates settle quickly.
Concentration of the solute
Relative Supersaturation=
Equilibrium solubility
When (Q-S)/S is large precipitate tends to be colloidal. When Q-S)/S is small precipitate tends
to be crystalline.
4- write at least three ways to maximize crystalline suspension.
~
~
~
~
minimize S
maximize Q
adjusting pH levels to precipitation medium.
digestion of crystalline precipitates.
5- describe the preparation of 25 ml 6.0 M HCI from concentrated HCI.
Formula
HCl, hydrochloric acid
%( w/w )
37
M. Weight (g/Mol)
Density ( g/cm3)
1.189
36.5
In 25 mL 6. 0 M solution
HCl
In 1 gram concentrated HCl:
= 0.0101 mol HCl
Concentrated solution is needed
= 12.45 ml solution is needed.
Hence, take 14.45 ml concentrated solution and dilute it to 25.00 ml.
POST-LAB STUDIES
1 - what is the importance of pH control in this experiment?
Equilibrium favors the dissolution of Ni complex back to its mother liquor form it ph
level >5.
2- why do we perform chloride test?
To make sure that our precipitate is free of chloride, it is pure.
3- write the name of the alcohol used in this experiment. Why the amount of alcohol
must be controlled?
DMG, dimethylglyoxime. The amount must be controlled since nickel complex is slightly
soluble in alcohol and also DMG can form precipitate with Ni
4- write the chelation reaction.
Titrations Based on Complex Formation
A complexation reaction involves a reaction between a metal ion M and another molecule or
ionic entity (a ligand) containing at least one atom with an unshared pair of electrons that can
be represented as:
Ligands are called unidentate when they can donate one pair of electrons and multidentate
when two or more pairs can be donated to form coordinate bonds with the metal ion. When a
multidentate ligand forms two or more coordinate bonds with the save acceptor metal ion, a
ring structure results, and the compound is said to be a chelate. The coordination number is
the number of electron pairs that a metal shares.
One of the most important ligands is EDTA, ethylene diamine tetra acetic acid.
Properties of EDTA:
→ weak tetraprotic acid
→ hexadante ligand
→ it combines with metal ions 1:1 regardless of the
charge on the cation
→ formation of metal-EDTA complex is dependent upon
the PH of the solution
→ for the titrations of magnesium and calcium ions,
alkaline medium is needed, pH 10.
Formation constant
→ end points of EDTA titrations can be established by using metal ion indicators.in this
experiment we will use eriochrome black T.it is an organic dye that forms intensely colored
chlates with metal ions.
In acidic moderately basic solutions, the predominant acid base equilibrium exhibited by the
indicator as follows:
Red
Blue
Eriochrome black-t
The metal complexes of erio-T is are generally red to observe color charge within this indicator,
then, it is necessary to adjust PH to 7 or above so that the blue form of the indicator
predominates in the solution. The end-point reaction is then
Red
Blue
Titration Methods Employing EDTA
A) direct titration: suitable for metals that can readily react rapidly with EDTA.
~ Mg
B) back titration: useful for cations that form very stable EDTA complexes and for which a
satisfactory indicator is not available.in this method a measured excess EDTA is added and the
excess EDTA determined by back-titration using Mg.
C) displacement titration: In this process on unmeasured amount excess EDTA in the form of
MgEDTA complex introduced. If the metal forms a more stable complex than that of magnesium
then displacement reaction takes place:
Where M represent the analyte cation. The liberated mg ion is then titrated with EDTA solution.
This method is suitable when no satisfactory indicator is available.
Determination of Water Hardness by Titration with EDTA
The primary components of water hardness is magnesium and calcium. Boil water to remove
CO2 Before determining the hardness remove CO3 and HCO3 present in water by adding HCl
Adjust the PH to 10 again before titration the sample with EDTA.
1) Define the followings and give examples to each one: Ligand, chelate, coordination number.
A chelate is produced when a metal ion coordinates with two or more donor groups of a single ligand to
form a five- or six-membered heterocyclic ring.
A ligand is an ion or a molecule that forms a covalent bond with a cation
or a neutral metal atom by donating a pair of electrons, which are then shared by the two.
The number of covalent bonds that a cation tends to form with electron donors is its coordination
number.
2) Draw the structure of EDTA and show the bonding sites.
Four carboxyl groups
Two amine groups
3) Write the name of a typical metal-ion indicator used for metal ions in EDTA titrations. Write
its dissociation equilibria.
Red
Blue
Eriochrome black-t
5) What do you understand from hard water?
The water that contains calcium carbonate molecules more than some extend.Rich in magnesium
and calcium ions.
6) What are the advantages of using EDTA as complexing agent?
The ability of EDTA to form complexes with metals is responsible for its widespread use as a
preservative in foods and in biological samples.
7) What are the advantages of multidentate ligands over their unidentate counterparts? '
Multidentate ligands are preferable to unidentate ligands for complexometric titrations because
they form more stable complexes. This is due to the chelate effect, which is a thermodynamic
phenomenon that occurs when a multidentate ligand binds to a metal ion, forming a ring-like
structure.
POST -LAB STUDIES:
1) Explain the importance of pH control in complexometric titrations. Why is it necessary to
buffer the solution, rather than simply adjust the pH with HCI or NaOH?
Solution must be alkaline since EDTA is a weak acid. Solution must be buffered since additional
EDTA can also chance pH.
2) Can you determine Mg by displacement or back titration? (Hint: You are not restricted to
use the reagents used in your experiment)
3) You need to know the amount of EDTA added to solution in back titration method but in
displacement titration you do not need to know the amount of MgEDTA added. Why? Explain
your answer by giving related reactions.
4) What are the effects of water hardness in our life?
Hard water can interfere with the action of soaps and detergents and can result in deposits of calcium
carbonate, calcium sulphate and magnesium hydroxide (Mg(OH)2) inside pipes and boilers, causing lower
water flows and making for less efficient heating , and dry skin.
5) What is the minimum pH value needed for satisfactory titrations of Ca?* and Mg2+ cations
with EDTA?
10?
6) Why is a small amount of MgY often added to a water specimen that is to be titrated
for water hardness?
The absence of magnesium causes slow reaction. To speed up MgY is added.
The
The Q-test:
Q-test:
An outlier is a result that is quite different from the others in the data set and might be due
to a gross error. The Q test is a simple, widely used statistical test for deciding whether a
suspected result should be retained or rejected. In this test, the absolute value of the
difference between the questionable result and its nearest neighbor is divided by the spread w
of the entire set to give the quantity Q:
If
reject
Interval:
Finding
Confidence Interval:
the Confidence
Finding the
The
The Coefficient
Coefficient of
of Variation(%RDS
Variation(%RDS ):
):
While choosing t
value we should
consider the
degrees of
freedom
NEUTRALIZATION TITRATIONS
A standard solution is a reagent of known concentration. Standard solutions are used in
titrations and in many other chemical analyses.
Standard solutions play a central role in all volumetric methods of analysis. The ideal standard
solution for a volumetric method should:
• be sufficiently stable so that it will not be necessary to determine its
concentration (to standardize it) so frequently.
• react rapidly with the analyte so that the time required between each addition
of the titrant is minimized.
• react more or less completely with the analyte so that satisfactory end points
are realized.
• undergo a selective reaction with the analyte that can be described by a simple
balanced equation.
primary standard is a highly purified compound that serves as a reference material in all
titrimetric methods. Important requirements for a primary standard are:
• High purity
• Stability in air
• Absence of hydrate water so that the composition doesn't change with variations in
relative humidity
• Reasonable solubility in the titration medium
• Reasonably large formula weight so that the relative error associated with weighing
is minimized
• Readily available at a modest cost
In a standardization, the concentration of a volumetric solution is determined by titrating it
against a carefully measured quantity of a primary or secondary standard or an exactly known
volume of another standard solution.
The Effect of Carbon Dioxide on Standard Base Solutions
distilled water is sometimes supersaturated with the CO2 gas and thus contains sufficient
carbonic acid to cause detectable errors. In order to test the distilled water to be used in
neutralization titrations, take about 500 mL from the source of production. Add 5 drops of
phenolphthalein and titrate with ~0.10 M NAOH. Less than 0,2 0 0,3 ml of base should be
used to see the first pink color.
CO2 (from the atmosphere) reacts with the hydroxides of sodium, potassium, and barium
(solution or solid), producing carbonates:
Absorption of CO2 by a standardized solution of sodium or possium hydroxide leads to a
negative systematic error (called carbonate error), in analyses when a basic range indicator is
used. Here, each carbonate ion (produced from two hydroxide ions) reacts with only one
hydronium ion when the indicator changes its color:
However, no systematic error occurs when an acidic range indicator is used. Each carbonate ion
produced will react with two hydronium ions of the acid:
The amount of hydronium ion consumed by this reaction equals the amount of hydroxide lost
during formation of the carbonate ion, thus, no error occurs.
Procedure:
• Dried potassium hydrogen phthalate, KHC8H404 (KHP) as primary standard
• Dried sodium carbonate, Na2CO3 as primary standard
Preparation and Standardization of 0.10 M NaOH
Titrate the solution with 0.10 M NaOH prepared until a faint pink color appears and persists at
least 30 seconds. The net reaction during titration can be written as follow:
Where, HP is hydrogen phthalate ion coming from KHP, OH is hydroxide ion coming from NaOH
and p2 is phthalate ion.
3)What is the difference between primary standard and standard solution?
A primary standard is an
ultrapure compound that
serves as the reference
material for a titration or
for another type of
quantitative analysis
A standard solution is a
reagent of known
concentration. Standard
solutions are used in titrations
and in many other chemical
analyses.
4)Define and compare end-point and equivalence point.
The equivalence point is
the point in a titration
when the amount of
added standard reagent is
equivalent to the amount
of analyte.
The end point is the
point in a titra- tion
when a physical change
occurs that is associated
with the condition of
chemical equivalence.
5)Why can weak acids not be used as a titrant in the titrations?
Weak acids and weak bases react incompletely with analytes.
1)Explain the effect, if any, of each of the following sources of error on the molarity
of the base as determined in the experiment; i.e., would the experimental value for
molarity be too high or too low, and why?
a)The Erlenmeyer flask in which the titration was performed contained several mL of
distilled water from the rinsing at the time the KHP was weighed into it.
•
+ error
b) If a bubble appeared in the tip during the titration.
•.
+ error
c)If liquid splashed from the titration beaker before the end point had been
reached.
•.
+error
d) If the burette was not rinsed with the NaOH solution following the rinsing with
distilled water.
4)Why do we boil and then cool distilled water before the preparation of NaOH solution?
Why is this process not necessary for the preparation of HCI?
no systematic error occurs when an acidic range indicator is used. Each carbonate ion produced
will react with two hydronium ions of the acid.The amount of hydronium ion consumed by this
reaction equals the amount of hydroxide lost during formation of the carbonate ion, thus, no
error occurs.
5) In the titration of a weak and strong acid mixture with a strong base, which acid is
titrated first? Write the related titration reactions by drawing the corresponding titration
curve.
Equivalence point for acidic-acid
Equivalence - point for HCl
Strong acid titrated first
.........
Kjedahl
Method
Kjedahl
Method
Digestion
Distillation
=
In this method, a certain substance or sample is
heated in the presence of sulphuric acid. The
acid breaks down the organic substance via
oxidation and reduced nitrogen in the form of
ammonium sulphate,an inorganic salt is liberated.
Potassium sulphate is usually added to increase
the boiling point of the medium. Catalysts like
mercury, selenium, copper, or ions of mercury
or copper are also used in the digestion process.
The sample is fully decomposed when we obtain
a clear and colorless solution.
Addition of potassium sulfate can increase the
boiling point of sulfuric acid to 390°C, which is
330°C without the potassium sulfate. This
significantly increases the rate of reaction.
If the temperature
goes too high, over
400°C, volatile
nitrogen
components may be
lost to the
atmosphere.
Distillation is adding excess base to the acid
digestion mixture to convert NH4+ to NH3,
followed by boiling and condensation of the
NH3 gas in a receiving solution.
→ raise the pH by NaOH
separating the nitrogen away from the
digestion mixture by distilling the ammonia
(converting it to a volatile gas, by raising the
temperature to boiling point) and then
trapping the distilled vapors in a special
trapping solution of boric acid (HBO3). The
ammonia is bound O to the boric acid in the
form of ammonium borate complex.
TITRATION
Then titrate the ammonia solution from the
distillation step by HCl
If salt/ acid ratio is
too high, a
considerable amount
of material will
"saltout". A certain
amount of salting can
be managed by adding
warm water.
Summary
Catalyst: Se
Analyte: nitrogen
Titrant:HCl
Boric acid: traps NH3
Potassium sulfate: increases the
boiling point of sulfuric acid
NaOH changes ammonium ion to
ammonia (g)