EDTA Titration of Ca2+ and Mg2+ in Natural Waters

Introduction

The most common multivalent metal ions in natural waters are Ca2+ and Mg2+. In this

experiment, we will find the total concentration of metal ions that can react with

EDTA, and we will assume that this equals the concentration of Ca2+ and Mg2+. In a

second experiment, Ca2+ is analyzed separately after precipitating Mg(OH)2 with strong

base.

Background

One of the factors that establish the quality of a water supply is its degree of

hardness. The hardness of water is defined in terms of its content of calcium and

magnesium ions. Since an analysis does not distinguish between Ca2+ and Mg2+, and

since most hardness is caused by carbonate deposits in the earth, hardness is

usually reported as total parts per million calcium carbonate by weight. A water

supply with a hardness of 100 parts per million would contain the equivalent of 100

grams of CaCO3 in 1 million grams of water or 0.1 gram in one liter of water. In the

days when soap was more commonly used for washing clothes, and when people

bathed in tubs instead of using showers, water hardness was more often directly

observed than it is now, since Ca2+ and Mg2+ form insoluble salts with soaps and

make a scum that sticks to clothes or to the bath tub. Detergents have the distinct

advantage of being effective in hard water, and this is really what allowed them to

displace soaps for laundry purposes. Calcium in water will be analyzed this week by

EDTA titration and next week by atomic absorption analysis and the results compared.

Water hardness can be readily determined by titration with the chelating agent EDTA

(ethylenediaminetetraacetic acid). This reagent is a weak acid that can lose four

protons on complete neutralization; its structural formula is below.

The four acid sites and the two nitrogen atoms all contain unshared electron pairs,

so that a single EDTA ion can form a complex with up to six sites on a given

cation. The complex is typically quite stable, and the conditions of its formation can

ordinarily be controlled so that it contains EDTA and the metal ion in a 1:1 mole

ratio. In a titration to establish the concentration of a metal ion, the EDTA that is

added combines quantitatively with the cation to form the complex. The end point

occurs when essentially all of the cation has reacted.

In this experiment you will standardize a solution of EDTA by titration against a

standard solution made from calcium carbonate, CaCO3. You will then use the

EDTA solution to determine the hardness of an unknown water sample. Since both

EDTA and Ca2+ are both colorless, it is necessary to use a rather special indicator

to detect the end point of the titration. The indicator you will employ is called

Eriochrome Black T, which forms a rather stable wine-red complex, MgIn-, with the

magnesium ion. A tiny amount of this complex will be present in the solution during

the titration. As EDTA is added, it will complex free Ca2+ and Mg2+

ions leaving

the MgIn- complex alone until essentially all of the calcium and magnesium has

been converted to chelates. At this point, the EDTA concentration will increase

sufficiently to displace Mg2+ from the indicator complex; the indicator reverts to an

acid form, which is sky blue, and this establishes the end point of the titration.

The titration is carried out at a pH of 10, in an NH3-NH4+

buffer, which keeps the

EDTA (H4Y) mainly in the half-neutralized form, H2Y2-, where it complexes the

Group IIA ions very well but does not tend to react as readily with other cations

such as Fe3+ that might be present as impurities in the water. Taking H4Y and H3In

as the formulas for EDTA and Eriochrome Black T respectively, the equations for

the reactions that occur during the titration are as follows.

A little magnesium ion is already present in the solutions you will be using to form

the initial red complex with the indicator.

This experiment requires that you recall the concepts of stoichiometry, molarity and

dilutions. You may need to refresh your memory before you begin the prestudy and

the experiment.

REAGENTS

Buffer (pH 10): Add 142 mL of 28 wt % aqueous NH3 to 17.5 g of NH4Cl and dilute

to 250 mL with water.

Eriochrome black T indicator: Dissolve 0.2 g of the solid indicator in 15 mL of

triethanolamine plus 5 mL of absolute ethanol.

50 wt % NaOH: Dissolve 100 g of NaOH in 100 g of H2O in a 250-mL plastic bottle.

Store tightly capped. When you remove solution with a pipet, try not to disturb the

solid Na2CO3 precipitate.

PROCEDURE

1. Dry Na2H2 EDTA·2H2O (FM 372.24) at 80° C for 1 h and cool in the desiccator.

Accurately weigh out ~0.6 g and dissolve it with heating in 400 mL of water in a

500-mL volumetric flask. Cool to room temperature, dilute to the mark, and mix well.

2. Pipet a sample of unknown into a 250-mL flask. A 1.000-mL sample of seawater

or a 50.00-mL sample of tap water is usually reasonable. If you use 1.000 mL of

seawater, add 50 mL of distilled water. To each sample, add 3 mL of pH 10 buffer

and 6 drops of Eriochrome black T indicator. Titrate with EDTA from a 50-mL buret

and note when the color changes from wine red to blue. Practice finding the end

point several times by adding a little tap water and titrating with more EDTA. Save a

solution at the end point to use as a

color comparison for other titrations.

3. Repeat the titration with three samples to find an accurate value of the total Ca2+

+ Mg2+ concentration. Perform a blank titration with 50 mL of distilled water and

subtract the value of the blank from each result.

4. For the determination Ca2+, pipet four samples of unknown into clean flasks

(adding 50 mL of distilled water if you use 1.000 mL of seawater). Add 30 drops of

50 wt % NaOH to each solution and swirl for 2 min to precipitate Mg(OH)2 (which

may not be visible). Add ~0.1 g of solid hydroxynaphthol blue to each flask. (This

indicator is used because it remains blue at higher pH than does Eriochrome black

T.) Titrate one sample rapidly to find the end point; practice finding it several times,

if necessary.

5. Titrate the other three samples carefully. After reaching the blue end point, allow

each sample to sit for 5 min with occasional swirling so that any Ca(OH)2 precipitate

may redissolve. Then titrate back to the blue end point. (Repeat this procedure if the

blue color turns to red upon standing.) Perform a blank titration with 50 mL of

distilled water.

6. Calculate the total concentration of Ca2+ and Mg2+, as well as the individual

concentrations of each ion. Calculate the relative standard deviation of replicate

titrations.

Analysis of Hydrogen Peroxide A Redox Titration

Introduction

Hydrogen peroxide is regarded as an "environmentally friendly" alternative to chlorine

for water purification and wastewater treatment. Because hydrogen peroxide

decomposes in the presence of heat, light, or other catalysts, the quality of a

hydrogen peroxide solution must be checked regularly to maintain its effectiveness.

The concentration of hydrogen peroxide can be analyzed by redox titration with

potassium permanganate.

Concepts

Redox reaction, Oxidizing and reducing agents, Titration, Half-reactions

Background

Titration is a method of volumetric analysis-the use of volume measurements to

analyze the concentration of an unknown. The most common types of titrations are

acid-base titrations, in which a solution of an acid, for example, is analyzed by

measuring the amount of a standard base solution required to neutralize a known

amount of the acid. A similar principle applies to redox titrations. If a solution

contains a substance that can be oxidized, then the concentration of that substance

can be analyzed by titrating it with a standard solution of a strong oxidizing agent.

The equation for an oxidation-reduction reaction can be balanced by assuming that it

occurs via two separate half-reactions. In this experiment, potassium permanganate

will be used as the titrant to analyze the concentration of hydrogen peroxide in a

commercial antiseptic solution. The permanganate ion acts as an oxidizing agent--it

causes the oxidation of hydrogen peroxide. The oxidation half-reaction shows that

two electrons are lost per molecule of hydrogen peroxide that is oxidized to oxygen

gas (Equation 1). The permanganate ion, in turn, is reduced from the +7 oxidation

state in MnO4- to the +2 oxidation state in Mn2+. The reduction half-reaction

shows a gain of five electrons (Equation 2).

H2O2(aq) → O2(g) + 2H+(aq) + 2e- Equation 1

MnO4-(aq) + 8H+(aq) + 5e- → Mn2+(aq) + 4H2O(l) Equation 2

Purple

Experiment Overview

The purpose of this experiment is to analyze the percent hydrogen peroxide in a

common "drugstore" solution by titrating it with potassium permanganate. Standard

potassium permanganate solution will be added via buret to the hydrogen peroxide

solution. As the dark purple solution is added, it will react with the hydrogen peroxide

and the color will fade. When all of the hydrogen peroxide has been used up, the

"last drop" of potassium permanganate that is added will keep its color. The endpoint

of the titration is the point at which the last drop of potassium permanganate added

to the solution causes it to turn pink.

Materials

Distilled or deionized water, 100 mL

Hydrogen peroxide, H2O2, commercial antiseptic solution, 3 mL

Potassium permanganate solution, KMnO4, 0.025 M, 75 mL

Sulfuric acid solution, H2SO4 6 M, 15 mL

Beaker, 100- or 150-mL Buret, 50-mL, and buret clamp

Erlenmeyer flasks, 125-mL, 2 Graduated cylinder, 10- or 25-mL

Labels and/or markers Pipet, volumetric or serological, I-mL

Pipet bulb Ring stand

Wash bottle Waste disposal beaker, 250 mL

Safety Precautions

Sulfuric acid solution is severely corrosive to eyes, skin, and other body tissues.

Always add acid to water, never the reverse. Notify your teacher and clean up all

acid spills immediately. Potassium permanganate solution is a skin and eye irritant

and a strong stain-it will stain skin and clothing. Avoid contact of all chemicals with

eyes and skin. Wear chemical splash goggles, chemical-resistant gloves, and a

chemical-resistant apron. Wash hands thoroughly with soap and water before leaving

the lab.

Procedure

1. Obtain about 75 mL of potassium permanganate standard solution in a small

beaker. Record the precise molarity of the solution in the data table.

2. Rinse a clean 50-mL buret first with dH2O with two 5-mL portions of potassium

permanganate solution.

3. Clamp the buret to a ring stand using a buret clamp and place a waste beaker

under the buret.

4. Fill the buret with potassium permanganate solution until the liquid level is just

above the zero mark.

5. Open the stopcock on the buret to allow any air bubbles to escape from the tip.

Close the stopcock when the liquid level in the buret is between the 0- and 5-mL

mark.

6. Record the precise level of the solution in the buret. This is the initial volume of

the potassium permanganate solution for Trial 1. Note: Volumes are read from the top

down in a buret. Always read from the bottom of the meniscus and remember to

include the appropriate number of significant figures.

7. Using a volumetric or serological pipet, transfer 1.00 mL of the commercial

hydrogen peroxide solution into a 125-mL Erlenmeyer flask.

8. Add about 25 mL of distilled or deionized water to the flask.

9. Measure 5 mL of 6M sulfuric acid into a graduated cylinder and carefully add the

acid to the solution in the Erlenmeyer flask. Gently swirl the flask to mix the solution.

10. Position the flask under the buret so that the tip of the buret is within the flask

but at least 2 cm above the liquid surface. Place a piece of white paper under the

flask to make it easier to detect the endpoint.

11. Open the buret stopcock and allow 5-8 mL of the potassium permanganate

solution to flow into the flask. Swirl the flask and observe the color changes in the

solution.

12. Continue to add the potassium permanganate solution slowly, drop-by-drop, while

swirling the flask. Use a wash bottle to rinse the sides of the flask with distilled water

during the titration to ensure that all of the reactants mix thoroughly.

13. When a light pink color persists in the titrated solution while swirling the flask, the

endpoint has been reached. Close the stopcock and record the final volume of the

permanganate solution in the data table (Trial 1).

14. Subtract the initial volume of the permanganate solution from the final volume to

obtain the volume of KMnO4 added. Enter the answer in the data table.

15. Pour the titrated solution into a waste disposal beaker and rinse the flask with

distilled water. 16. Repeat the titration (steps 6-15) two more times (Trials 2 and 3).

Record all data in the data table.

17. Dispose of the solution in the waste beaker as directed by your instructor.

VOLUMETRIC DETERMINATION OF CARBONATE AND

BICARBONATE

Background

Acid base titrations with weak acids and bases represent well the challenges and

complexities of many biological, medical, environmental, and industrial measurements.

Further, mixtures represent well the demands faced in modern analytical science. In

the present laboratory, you will receive a mixture containing sodium carbonate (a

weak base), sodium bicarbonate (a weak acid and base) and other un-reactive

substances. Techniques used here will provide a sophisticated analytical challenge

with a demand to understand of weak acid-base equilibria and calculations involving

dissociation constants. The weak acid base character of these analytes and the

difficulties in resolution of composition using indicators will cause us to use some

maneuvers not commonly used in undergraduate laboratories. The central challenge in

this measurement is that we need to quantitatively determine both HCO3- or CO3

- and

then one of the components alone. The remaining constituent then can be calculated

by subtraction of the one component from total alkalinity determination (mindful of

stoichiometry).

A titration of your sample directly with strong acid will lead to reactions with both

carbonate and bicarbonate in stoichiometric proportions (see reactions on last page

of this laboratory). The sum of all acid moles or equivalents used is a useful piece of

information called total alkalinity in water quality or environmental studies; but, total

alkalinity alone does provide quantitative results for the mixture of carbonate and

bicarbonate. But, how can we measure HCO3- or CO3

- using volumetric methods?

Sadly, end points to distinguish individual constituents with mixtures such as NaHCO3

and Na2CO3 are difficult to recognize using indicators or dyes. Perhaps a

potentiometric titration might be possible using a pH indicating electrode, but that is

not our goal. The goal in this laboratory is to use the technique of back-titration with

analysis of a mixture (through a little help from precipitation chemistry).

One approach might be to titrate with strong base to determine HCO3- however, the

endpoint will be difficult to determine. Thus, our measurement challenge would be

simplified if we could remove or isolate the bicarbonate from the carbonate. But this

cannot be done easily or conveniently. Both of these complications can be solved

using the method of excess reagent and back-titration with strong reagents. What is

a back-titration? In a back-titration, an exact and excess amount of strong base is

added to solution to react with HCO3- (there will be no reaction between OH- and

CO3-) and some OH- will remain as residual base in solution. We may arrange a

clean back-titration of residual OH- with strong acid when all forms of carbonate are

removed from the sample. This can be accomplished when a soluble barium salt is

added to the solution before a titration is attempted, carbonate will be precipitated as

BaCO3 and will be eliminated as an interference.

Do you see the overall view of the analytical attack on this sample?

a. Determination of all HCO3- and CO3

- using a strong acid titration.

b. Determination of HCO3- using back-titration with strong acid of excess OH- after

eliminating CO3- (from original sodium carbonate and from the reacted bicarbonate)

by precipitation as BaCO3 from the solution. As with all other quantitative methods for

acid-base chemistry, solutions of acids and bases should never be accepted for use

in volumetric measurements without standardization of the solutions immediately

before use. Ideally, solutions should standardized and used the day of calibration. A

small and acceptable (for Chem 371) error will arise if solutions are left overnight or

another day. Solutions left for more than two days should be re-standardized as error

increases with time of storage. Thus, your measurements in this laboratory will involve

both calibration of strong acid/base solutions and the triplicate determination of your

sample for A) total alkalinity and B) bicarbonate using the back-titration method.

Procedures

A. Preparing and Standardizing solutions

Your HCl solution will be only approximate concentration regardless of how much

care you made in preparation. Fortunately, primary standards may be used to

calibrate the solution. Store your acid solution in a plastic bottle.

1. Use the information in the table on the inside cover of this text to calculate the

volume of concentrated (~37 wt%) HCl that should be added to 1 L (read as one

liter) of distilled water to produce ~0.1 M HCl, and prepare this solution.

[Alternatively, your TA might have prepared for you a solution (if so, collect ca. 500

ml of reagent for standardization and analysis)].

2. Dry primary standard-grade sodium carbonate for 1 h at 110°C and cool it in a

desiccator. Bring to constant weight. [A common batch of sodium carbonate may

exist in the oven. If so, remove enough for your uses and bring to constant weight.

Return the common batch to the oven]

3. Weigh four samples containing enough Na2CO3 to react with ~25 ml of 0.1 M HCl

and place each in a 125 ml flask. As you are ready to titrate each one, add ~25 ml

of distilled water, swirl the flask to dissolve the solid and add 3 drops of bromocresol

green indicator solution. Titrate one rapidly to a green color to find an approximate

end point. The stoichiometry of the reaction is shown here:

2HCl + Na2CO3 ∙.--------> CO2 + 2NaCl

4. Carefully titrate each of the other three samples to the blue to green end point.

During each titration, you should periodically tilt and rotate the flask to wash all liquid

from the walls into the bulk solution using deionized water from a plastic squeeze

bottle. When very near the end, you should deliver less than 1 drop of titrant at a

time. To do this, carefully suspend a fraction of a drop from the buret tip, touch it to

the inside wall of the flask, wash it into the bulk solution by careful tilting, and swirl

the solution (this is called splitting a drop). A slight refinement to correct for

equilibrium chemistry should be added now. When the end point is reached, warm

the solution slightly to expel CO2. The solution should return to a blue color.

5. Carefully add HCl from the buret until the solution turns green again (you may

have to split drops to get a refined end point).

6. A blank titration should also be performed, by using 3 drops of indicator in 50 mL

of 0.05 M NaCl.

7. Subtract the volume of HCl needed for the blank titration form the required to

titrate Na2CO3. Calculate the mean HCl molarity, standard deviation, and relative

standard deviation. All NaOH solutions come with impurities and contamination the

make each NaOH solution imprecise unless calibrated. Calibrated NaOH solutions

have limited shelf-life even when protected against atmospheric carbon dioxide.

Furthermore, store your solution only in a plastic bottle and insure that the cap is

free of solution between uses.

1. You will be provided 1 liter of a nominal 0.1M solution of NaOH or a concentrated

NaOH stock that must be diluted to ~0.1M. This must be standardized against the

HCl solution just standardized above.

2. Place (using a volumetric pipette) 25.00 ml of your dilute HCl solution into a 200

or 250 ml Erlenmeyer flask. Add 3 drops of phenolphthalein indicator solution, and

titrate to an end point with the NaOH solution. You should be able to anticipate what

volume will be require to reach the endpoint. So you may work fast and smart and

do not need a quick and dirty trial analysis. The end point is the first appearance of

faint pink color that persists for 15 s. (The color will slowly fade as CO2 from the air

dissolves in the solution.)

3. Make two more replicate determinations of 25.00 ml of NaOH solution using the

standardized HCl solution. Calculate the average molarity, standard deviation, and

%RDS of the NaOH solution using the formula VHClMHCl=VNaOHMNaOH. If you have

used some care the %RSD should be <0.2%. At this stage in the laboratory, you will

have standardized solutions of HCl and NaOH. Do not share or alter these solutions.

Always pour from your containers and never place pipets directly into reagent

solutions or analytical solutions.

B: Volumetric Determination of Carbonate and Bicarbonate

1. Solid unknown should be stored in a desiccator to keep dry; however, do not heat

the sample. Even mild heating at 50°-100°C converts NaHCO3 to Na2CO3. Accurately

weigh 2.0 to 2.5 grams of unknown into a clean 250 mL volumetric flask. A funnel

may be used to aid transfer of sample to the volumetric flask. Rinse the funnel

repeatedly with small portions of freshly boiled and cooled water to dissolve and

transfer the sample to the flask. Remove the funnel, mix well, and dilute to the mark.

This is your unknown solution.

2. Pipette a 25.00 mL aliquot of the unknown solution into a 250 mL Erlenmeyer

flask, add 3 drops of bromocresol green indicator solution and titrate (to a green

endpoint) using your standardized HCl solution (from above). You should do this

rapidly as a quick and dirty sample. Then, repeat in triplicate this procedure carefully

with 25.00 mL samples of unknown solution.

3. Pipette a 25.00 mL aliquot of the unknown solution and 50.00 mL of your

standardized NaOH solution into a 250-mL flask. Swirl thoroughly and add using a

graduated cylinder 10 mL of 10% wt BaCl2 solution. Swirl again to precipitate BaCO3,

add 2 drops of phenolphthalein indicator, and immediately titrate (the endpoint is a

faint pink) with your standardized HCl solution. Do all of this quickly without delay.

4. Repeat in triplicate this procedure carefully with 25.00 mL samples of your

unknown solution.

5. Calculate the total alkalinity in the unknown solution and calculate the bicarbonate

concentration in the unknown solution. Based upon volumetric amounts and dilution

volumes, calculate percent sodium carbonate and percent sodium bicarbonate in your

solid sample. You will have three independent determinations and calculations so you

may report average values for sodium carbonate (percent by weight), sodium

carbonate (percent by weight) and standard deviations for each.

6. Submit a laboratory report.

Purification methods by recrystallization

Introduction

To receive instruction for the use of common laboratory equipments and to learn

some separation techniques.

Background

Separation of Substances:

Every substance has a large number of physical and chemical properties. Physical

properties include color, smell, taste, solubility, density, electrical conductivity, heat

conductivity, melting and boiling points. When a physical change is observed, the

substance retains its chemical identity. In contrast to physical changes, when

chemical changes are observed, new substances are formed. Chemical properties

include decomposition by heating, and reactions of the substance with water, oxygen,

acids, bases, etc. So, physical changes are reversible, chemical changes are

irreversible (not reversible).

Recrystallization

Recrystallization is the method of choice for purifying organic solids. This technique

involves dissolving a solid in a solvent then inducing the solute to precipitate from

the solution. A more pure solid should precipitate if the impurities remain dissolved in

the solvent. The pure solid is collected by vacuum filtration and usually washed with

a minimal amount of ice-cold solvent.

All purification methods exploit some physical property of the substance.

Recrystallization depends on the solubility of a solid and the fact that solubility

increases with temperature. The success of any recrystallization relies on the fact that

at or near the boiling point, a given recrystallization solvent will dissolve more solid

than at room temperature. The result is that the crude solid can be dissolved in a

small amount of boiling solvent and as the solution cools, the solubility decreases

and eventually, the solid will precipitate (recrystallized) from solution.

The precipitation process is called crystallization. When an organic reaction produces

a compound that exists as a solid, the product may crystallize from the reaction

mixture or at some point after the solvents and by-products have been removed.

Frequently, this crystallized substance, the crude solid, contains impurities from the

reaction mixture and must be purified. A recrystallization attempts to separate the

compound from the remaining impurities which may be solvent residue, side-products

and salts, or some combination of all three, hence this is a purification technique.

In a typical procedure, the crude solid is dissolved in a recrystallization solvent by

heating the solvent/solid mixture or by adding boiling solvent to the crude. Once the

crude solid is completely dissolved, the solution is allowed to cool until the solution

becomes saturated and crystals begin to form. The flask is placed in an ice bath to

maximize the amount of solid that can precipitate from the solution. The recrystallized

solid is then collected by vacuum filtration and washed with cold solvent.

Occasionally, centrifugation is required to separate the recrystallized solid from the

recrystallization solvent.

In many cases, it appears that the crude solid does not dissolve completely. Adding

more solvent is not always the best practice since the result may be a solution that

will remain unsaturated even at lower temperatures; in other words, your compound

will not precipitate. If a crude sample does not dissolve completely, it may be that

some of the impurities are insoluble in the solvent being used. To separate the

suspended solids, a hot filtration step is performed. To perform a hot filtration, the

heterogeneous mixture is filtered quickly while still near the boiling point of the

solvent to remove the undissolved solids. The filtrate should still be hot enough to

keep your compound in solution. Frequently, the product may start to precipitate on

the filter paper during the hot filtration step. To prevent this, the glassware is warmed

by boiling a small amount (3-5 mL) of the recrystallizing solvent in a clean

Erlenmeyer flask with the filter funnel sitting on top. This Erlenmeyer flask will be

used to collect the filtrate. Once the glassware is hot from the solvent vapors, the

undissolved solids are filtered as quickly as possible. A ready supply of boiling

solvent should be on hand in case the compound begins to crystallize on the filter.

Add small amounts (1-2 mL) of boiling solvent to dissolve the compound so that it

passes into the filtrate. Keep in mind that you will most likely boil the filtrate to

evaporate the excess solvent used to re-dissolve the product. The filtrate is then

allowed to cool and the pure crystals are collected by vacuum filtration and washed

on the filter.

Occasionally, a crude solid will contain colored impurities which must be separated.

Activated charcoal (about the size of a pea) is added to the mixture and the contents

swirled to allow the colored substance to adsorb to the surface of the charcoal.

Afterwards, a hot filtration is needed to remove the charcoal.

When employing a hot filtration, the simplest operation to remove suspended solids,

either undissolved impurities or activated charcoal, is gravity filtration. Vacuum

filtration is not used during a hot filtration because the vacuum will cool the solvent

and crystallization of the product may occur. However, once recrystallization is

complete, vacuum filtration is the method of choice since solvent residue is quickly

removed from the pure crystals.

A word about washing purified solids on a filter: Swirl the contents of the flask and

try and get a suspension of solids then quickly pour the recrystallized material and

solvent into the filter funnel. Apply vacuum just long enough to draw the solvent into

the Erlenmeyer flask then remove the vacuum. The objective here is to produce wet

cake of recrystallized solid of uniform thickness. Next, gently add a shallow layer of

wash solvent over the wet cake, apply vacuum and draw the wash through the wet

cake. Repeat washings if necessary. Vacuum filtration is preferred over gravity

filtration in order to quickly separate the crystals from the solvent.

Choosing an appropriate solvent is key to a successful recrystallization. The organic

compound must have low solubility in the solvent at room temperature and moderate

solubility at the boiling point. Ideally, the solubility of the impurities is greater than the

organic compound at both high and low temperatures obviating the need for a hot

filtration step.

As with any purification step, some loss of product is expected. After collecting the

pure solid, some of the compound remains in solution in the filtrate. Using solubility

data, the amount of solid lost to the filtrate can be estimated. Solubility is often

reported as grams of solute per 100 mL of solvent (g/100 mL) at some temperature.

Single-solvent recrystallization

Typically, the mixture of "compound A" and "impurity B" are dissolved in the smallest

amount of hot solvent to fully dissolve the mixture, thus making a saturated solution.

The solution is then allowed to cool. As the solution cools the solubility of

compounds in solution drops. This results in the desired compound dropping

(recrystallizing) from solution. The slower the rate of cooling, the bigger the crystals

form.

In an ideal situation the solubility product of the impurity, B, is not exceeded at any

temperature. In that case the solid crystals will consist of pure A and all the impurity

will remain in solution. The solid crystals are collected by filtration and the filtrate is

discarded. If the solubility product of the impurity is exceeded, some of the impurity

will co-precipitate. However, because of the relatively low concentration of the

impurity, its concentration in the precipitated crystals will be less than its

concentration in the original solid. Repeated recrystallization will result in an ever

purer crystalline precipitate. The purity is checked after each recrystallization by

measuring the melting point, since impurities lower the melting point. NMR

spectroscopy can also be used to check the level of impurity. Repeated

recrystallization results in some loss of material because of the finite solubility of

compound A.

The crystallization process requires an initiation step, such as the addition of a

"seed" crystal. In the laboratory a minuscule fragment of glass, produced by

scratching the side of the glass recrystallization vessel, may provide the nucleus on

which crystals may grow. Successful recrystallization depends on finding the right

solvent. This is usually a combination of prediction/experience and trial/error. The

compounds must be more soluble at the higher temperature than at the lower

temperatures. Any insoluble impurity is removed by the technique of hot filtration.

Multi-solvent recrystallization

This method is the same as the above but where two (or more) solvents are used.

This relies on both "compound A" and "impurity B" being soluble in a first solvent. A

second solvent is slowly added. Either "compound A" or "impurity B" will be insoluble

in this solvent and precipitate, whilst the other of "compound A"/"impurity B" will

remain in solution. Thus the proportion of first and second solvents is critical.

Typically the second solvent is added slowly until one of the compounds begins to

crystallize from solution and then the solution is cooled. Heating is not required for

this technique but can be used.

The reverse of this method can be used where a mixture of solvent dissolves both A

and B. One of the solvents is then removed by distillation or by an applied vacuum.

This results in a change in the proportions of solvent causing either "compound A" or

"impurity B" to precipitate.

Hot filtration-recrystallization

Hot filtration can be used to separate "compound A" from both "impurity B" and

some "insoluble matter C". This technique normally uses a single-solvent system as

described above. When both "compound A" and "impurity B" are dissolved in the

minimum amount of hot solvent, the solution is filtered to remove "insoluble matter

C". This matter may be anything from a third impurity compound to fragments of

broken glass. For a successful procedure, one must ensure that the filtration

apparatus is hot in order to stop the dissolved compounds crystallizing from solution

during filtration, thus forming crystals on the filter paper or funnel.

One way to achieve this is to heat a conical flask containing a small amount of

clean solvent on a hot plate. A filter funnel is rested on the mouth, and hot solvent

vapors keep the stem warm. Jacketed filter funnels may also be used. The filter

paper is preferably fluted, rather than folded into a quarter; this allows quicker

filtration, thus less opportunity for the desired compound to cool and crystallize from

the solution.

Often it is simpler to do the filtration and recrystallization as two independent and

separate steps. That is dissolve "compound A" and "impurity B" in a suitable solvent

at room temperature, filter (to remove insoluble compound/glass), remove the solvent

and then recrystallize using any of the methods listed above.

Compounds Benzophenone Naphthalene

Solvent EtOH/H2O(2:1) MeOH EtOH

Procedure

Students will be assigned approximately 1 gram each of two compounds from the list

below. The recrystallizing solvent for each compound is also given.

1. Record the name of each compound and the solvent used in your notebook.

2. Record your observations such as

• time required for solid to dissolve

• small amount of impurities that did not dissolve

(or homogeneous solution obtained?)

• approx. time for crystals to form and approx. temp

(slightly hot, room temp, ice bath)

• physical characteristics of recrystallized solid

(needles, plates, amorphous, etc.)

3. Construct a simple data summary for each sample:

Mass of crude sample before recrystallization ___________

Amount of solvent used to dissolve your compound ___________

Amount of pure solid obtained after crystallization ___________

Percent recovery of your product ___________ (show your calculation!)

4. Check the melting point. The percent recovery was high but the melting point of

the dried solid showed that the soled was contaminated with impurities. Explain.

Purification methods by Distillation

Introduction

To receive instruction for the use of common laboratory equipments and to learn

some separation techniques.

Background

In this experiment the efficiencies of packed and unpacked columns are compared

at total reflux. The separation of a binary mixture by fractional distillation is studied

by using refractive-index measurements to analyze the distillate.

Fractional distillation is the separation of a mixture into its component parts, or

fractions, such as in separating chemical compounds by their boiling point by heating

them to a temperature at which several fractions of the compound will evaporate. It is

a special type of distillation. Generally the component parts boil at less than 25 °C

from each other under a pressure of one atmosphere (atm). If the difference in

boiling points is greater than 25 °C, a simple distillation is used.

As an example, consider the distillation of a mixture of water and ethanol. Ethanol

boils at 78.4 °C while water boils at 100 °C. So, by heating the mixture, the most

volatile component will concentrate to a greater degree in the vapor leaving the liquid.

Some mixtures form azeotropes, where the mixture boils at a lower temperature than

either component. In this example, a mixture of 96% ethanol and 4% water boils at

78.2 °C, being more volatile than pure ethanol. For this reason, ethanol cannot be

completely purified by direct fractional distillation of ethanol-water mixtures.

The apparatus is assembled as in the diagram. (The diagram represents a batch

apparatus, as opposed to a continuous apparatus.) The mixture is put into the round

bottomed flask along with a few anti-bumping granules (or a Teflon coated magnetic

stirrer bar if using magnetic stirring), and the fractionating column is fitted into the

top. The fractional distillation column is set up with the heat source at the bottom on

the still pot. As the distance from the stillpot of the distillation column increases a

heat gradient is formed in the column where it is coolest at top and hottest at the

bottom. As the mixed vapor ascends the temperature gradient, some of the vapor

condenses and revaporizes along the temperature gradient. Each time the vapor

condenses and vaporizes, the composition of the more volatile liquid in the vapor

increases. This distills the vapor along the length of the column, and eventually the

vapor is composed solely of the more volatile liquid. The vapor condenses on the

glass platforms, known as trays, inside the column, and runs back down into the

liquid below, refluxing distillate. The efficiency in terms of the amount of heating and

time required to get fractionation can be improved by insulating the outside of the

column in an insulator such as wool, aluminium foil or preferably a vacuum jacket.

The hottest tray is at the bottom and the coolest is at the top. At steady state

conditions, the vapor and liquid on each tray are at equilibrium. Only the most volatile

of the vapors stays in gaseous form all the way to the top. The vapor at the top of

the column then passes into the condenser, which cools it down until it liquefies. The

separation is more pure with the addition of more trays (to a practical limitation of

heat, flow, etc.) The condensate that was initially very close to the azeotrope

composition becomes gradually richer in water. The process continues until all the

ethanol boils out of the mixture. This point can be recognized by the sharp rise in

temperature shown on the thermometer.

Typically the example above now only reflects the theoretical way fractionation

works. Normal laboratory fractionation columns will be simple glass tubes (often

vacuum jacketed, and sometimes internally silvered) filled with a packing, often small

glass helices of 4 to 7 mm diameter. Such a column can be calibrated by the

distillation of a known mixture system to quantify the column in terms of number of

theoretical plates. To improve fractionation the apparatus is set up to return

condensate to the column by the use of some sort of reflux splitter (reflux wire,

gago, Magnetic swinging bucket, etc.) - a typical careful fractionation would employ

a reflux ratio of around 10:1 (10 parts returned condensate to 1 part condensate take

off).

In laboratory distillation, several types of condensers are commonly found. The

Liebig condenser is simply a straight tube within a water jacket, and is the simplest

(and relatively least expensive) form of condenser. The Graham condenser is a spiral

tube within a water jacket, and the Allihn condenser has a series of large and small

constrictions on the inside tube, each increasing the surface area upon which the

vapor constituents may condense.

Alternate set-ups may utilize a "cow" or "pig" which is connected to three or four

receiving flasks. By turning the "cow" or "pig", the distillates can be channeled into

the appropriate receiver. A Perkin triangle is versatile piece of apparatus that can also

be used to collect distillation fractions which does not require a "cow" or "pig"

adapter. A Perkin triangle is most often used where the distillates are air-sensitive or

where the fractions distill and are collected under reduced pressure, but can be used

for a simple and fractional distillation.

Vacuum distillation systems operate at reduced pressure, thereby lowering the boiling

points of the materials. Note that the use of anti-bumping granules will not work at

reduced pressures.

Procedure

Place 30 ml of an unknown liquid mixture (15ml A + 15 ml B) that is to be purified

by simple distillation and for which the boiling point range is to be determined.

Assemble the assigned distillation apparatus (fractional). Transfer the unknown liquid

to a 50 mL round bottom flask (this will be the distilling pot). Add one boiling stone,

and proceed to distill the liquid into a 10 mL graduated cylinder (this will be the

receiver). Check the position of the thermometer (the bulb of the thermometer must

be below the arm of the distillation head) and make sure that the bottom of the

distillating pot touching the heating surface of the heating mantel. Securely attach a

piece of condenser tubing to each condenser outlets. Securely connect the other end

of the “water in” tubing to the water jet in the sink (or hood). Place the other end of

the “water out” tubing in the sink (or back of the hood). Plug in the heating mantle

and before heating your distillation apparatus or turning on the water for cooling the

condenser, have your laboratory instructor check your distillation apparatus.

After your laboratory instructor has checked your apparatus, slowly turn on the water

for condenser, and begin heating. Adjust the heating mantle to maintain a distillation

rate of one drop per second. As the lower boiling component is distilled, the boiling

point of the mixture in the distillation flask will increase.

Record the temperature after the first drop is collected and again after every 2 ml of

distillate is collected. After the 10 ml of distillate has been collected, you will have to

empty the graduated cylinder into a test tube as it fills. Cover and label the test tube

first fraction (component A). KEEP IT.

Collect the next 10 ml of distillate, again recording the temperature after every 1 ml

of distillate. After the second 10 ml of distillate has been collected, you will have to

empty the graduated cylinder into a test tube as it fills and DISCARD IT in organic

waste container.

Collection of last portion of distillate should continue until the temperature remains

constant. If the distillation flask is approaching dryness, remove the heat source

immediately and after cooling, transfer the distillate and any remaining liquid from the

flask to the third test tube (component B). KEEP IT.

Determine the boiling point range of the first fraction of the collected liquid and the

third portion of the collected liquid. Identify the unknowns by their boiling points using

the possible boiling points of compounds.

Construct a table like that given below, to record the temperature at the distillation

"head" as a function of volume distilled. You will record your data in report form. Plot

distillate temperature (℃) vs. volume of distillate (ml) collected for the mixture with

and without the fractionating column and use the graph to determine the boiling

points of the two compounds in the mixture and identify the compound in the

mixture.

Both sets of data will be plotted on the same graph, using different symbols

(colors). Label the two curves. Check the refractive-index.

Colligative Properties:

Freezing-Point Depression and Molar Mass

Purpose

To become familiar with colligative properties and to use them to determine the molar

mass of a substance.

Introduction

Solutions are homogeneous mixtures that contain two or more substances. The major

component is called the solvent, and the minor component is called the solute. Since

the solution is primarily composed of solvent, physical properties of a solution

resemble those of the solvent. Some of these physical properties, called colligative

properties, are independent of the nature of the solute and depend only upon the

solute concentration, measured in molality, or moles of solute per kilogram of solvent.

Background

The colligative properties include vapor pressure lowering, boiling point elevation,

freezing point depression, and osmotic pressure. The vapor pressure is the escaping

tendency of solvent molecules. When the vapor pressure of a solvent is equal to

atmospheric pressure, the solvent boils. At this temperature, the gaseous and liquid

states of the solvent are in dynamic equilibrium, and the rate of molecules going

from the liquid to the gaseous state is equal to the rate of molecules going from the

gaseous state to the liquid state.

The phase diagram below illustrates the effect of adding a solute to a pure

substance.

As is demonstrated by the phase diagram above, adding a solute to a solvent lowers

the freezing point and raises the boiling point; it also lowers the vapor pressure. The

new freezing point of a solution can be determined using the colligative property law:

∆Tf = kf m

The change in freezing point is equal to the molal freezing-point constant times the

molality of the solution. The molal freezing-point constant used is the constant for

the solvent, not the solute.

In this experiment, the molar mass of sulfur will be determined using the colligative

property law. The freezing point of naphthalene will be determined experimentally;

then a controlled solution of naphthalene and sulfur will be made, and the freezing

point of that solution will be determined. The difference in freezing point can be used

in the colligative property law to determine the experimental molality of the

solution, leading to a calculation of molecular weight.

The freezing temperature is difficult to ascertain by direct visual observation because

of a phenomenon called supercooling and also because solidification of solutions

usually occurs over a broad temperature range. Temperature-time graphs, called

cooling curves, reveal freezing temperatures rather clearly. The cooling curve will look

like the one below in figure 19.2:

In order to minimize supercooling, the solution will be stirred while freezing.

To determine the molar mass of a substance, one must simply divide the grams of

substance by the number of moles of substance present. All of these values will be

determined experimentally.

Procedure

Part A – Cooling Curve for Pure Naphthalene

1. A large test tube was weighed to the nearest .01 g using a standard laboratory

balance. Approximately 15 to 20 grams of naphthalene was added to the test tube.

The test tube was weighed again using the standard balance.

2. The following apparatus was assembled using the labeled parts.

3. The 600-mL beaker was nearly filled with water. It was heated to about 85°C. The

test tube was clamped in the water bath as shown in Figure 19.3 above. When most

of the naphthalene had melted, the stopper containing the thermometer and stirrer

was placed into the test tube. The thermometer was not allowed to touch the bottom

or sides of the

test tube.

4. When all of the naphthalene had melted, the test tube was removed from the

beaker of boiling water. The test tube was placed into a wide-mouthed bottle with

some paper towels at the bottom. The temperature reading from the thermometer was

recorded every

30 seconds. The naphthalene was stirred using the wire stirrer to ensure even

freezing. When the temperature remained constant for several readings, the

naphthalene was allowed to cool without further temperature readings.

Part B – Determination of the Molar Mass of Sulfur

1. Using weighing paper and an analytical balance, approximately 1.2 to 1.5 grams of

sulfur were weighed.

2. The test tube containing the naphthalene from Part A was placed into the water

bath from Part A and heated until all of the naphthalene had melted.

3. The stopper was removed and the sulfur was poured in. The stopper was replaced

and the solution was stirred gently until all of the sulfur had dissolved in the

naphthalene.

4. The test tube was removed from the water bath and placed into a wide-mouth

glass bottle with paper towels at the bottom. The temperature reading from the

thermometer was recorded every 30 seconds. The solution was stirred using the wire

stirrer to ensure even freezing. When the temperature remained constant for several

readings, the solution was allowed to cool without further temperature readings.

5. The test tube containing solid solution was handed to the laboratory instructor for

proper disposal of chemicals.