KWAME NKRUMAH UNIVERSITY OF SCIENCE ANDTECHNOLOGY

DEPARTMENT OF CHEMISTRY

YEAR TWO (CHEM 269)

PRACTICAL CHEMISTRY III

TITLE: DETERMINATION OF MIXTURES OF HALIDES BYAN INDIRECT METHOD

NAME: OPOKU ERNEST

EMAIL: ernopoku@gmail.com

DATE: 28TH OCTOBER, 2013

DERTERMINATION OF MIXTURES OF HALIDES BY ANINDIRECT METHOD

AIMS AND OBJECTIVES

1. To expose students to the use of the electronic balance.2. To expose students to argentometric titration.3. To determine the quantities of halides present in a mixture using Mohr’s method.

THEORY AND INTRODUCTION

The Halogens

There are six elements in Group VIIA, the next-to-last column of the periodic table. Asexpected, these elements have certain properties in common. They all form diatomicmolecules (H2, F2, Cl2, Br2, I2, and At2), for example, and they all form negatively chargedions (H-, F-, Cl-, Br-, I-, and At-).

When the chemistry of these elements is discussed, hydrogen is separated from the others andastatine is ignored because it is radioactive. (The most stable isotopes of astatine have half-lives of less than a minute. As a result, the largest samples of astatine compounds studied todate have been less than 50 ng.) Discussions of the chemistry of the elements in Group VIIAtherefore focus on four elements: fluorine, chlorine, bromine, and iodine. These elements arecalled the halogens (from the Greek hals, "salt," and gennan, "to form or generate") becausethey are literally the salt formers.

None of the halogens can be found in nature in their elemental form. They are invariablyfound as salts of the halide ions (F-, Cl-, Br-, and I-). Fluoride ions are found in minerals suchas fluorite (CaF2) and cryolite (Na3AlF6). Chloride ions are found in rock salt (NaCl), theoceans, which are roughly 2% Cl- ion by weight, and in lakes that have a high salt content,such as the Great Salt Lake in Utah, which is 9% Cl- ion by weight. Both bromide and iodideions are found at low concentrations in the oceans, as well as in brine wells in Louisiana,California, and Michigan.

The Halogens in their Elemental Form

Fluorine (F2), a highly toxic, colorless gas, is the most reactive element known so reactivethat asbestos, water, and silicon burst into flame in its presence. It is so reactive it even formscompounds with Kr, Xe, and Rn, elements that were once thought to be inert. Fluorine is sucha powerful oxidizing agent that it can coax other elements into unusually high oxidationnumbers, as in AgF2, PtF6, and IF7.

Fluorine is so reactive that it is difficult to find a container in which it can be stored. F2

attacks both glass and quartz, for example, and causes most metals to burst into flame.Fluorine is handled in equipment built out of certain alloys of copper and nickel. It still reacts

with these alloys, but it forms a layer of a fluoride on the surface that protects the metal fromfurther reaction.

Fluorine is used in the manufacture of Teflon or poly(tetrafluoroethylene), (C2F4)n

which is used for everything from linings for pots and pans to gaskets that are inert tochemical reactions. Large amounts of fluorine are also consumed each year to make thefreons (such as CCl2F2) used in refrigerators.

Chlorine (Cl2) is a highly toxic gas with a pale yellow-green color. Chlorine is a very strongoxidizing agent, which is used commercially as a bleaching agent and as a disinfectant. It isstrong enough to oxidize the dyes that give wood pulp its yellow or brown color, for example,thereby bleaching out this color, and strong enough to destroy bacteria and thereby act as agermicide. Large quantities of chlorine are used each year to make solvents such as carbontetrachloride (CCl4), chloroform (CHCl3), dichloroethylene (C2H2Cl2), and trichloroethylene(C2HCl3).

Bromine (Br2) is a reddish-orange liquid with an unpleasant, choking odor. The name of theelement, in fact, comes from the Greek stem bromos, "stench." Bromine is used to prepareflame retardants, fire-extinguishing agents, sedatives, antiknock agents for gasoline, andinsecticides.

Iodine is an intensely colored solid with an almost metallic luster. This solid is relativelyvolatile, and it sublimes when heated to form a violet-colored gas. Iodine has been used formany years as a disinfectant in "tincture of iodine." Iodine compounds are used as catalysts,drugs, and dyes. Silver iodide (AgI) plays an important role in the photographic process andin attempts to make rain by seeding clouds. Iodide is also added to salt to protect againstgoiter, an iodine deficiency disease characterized by a swelling of the thyroid gland.

Some of the chemical and physical properties of the halogens are summarized in the tablebelow. There is a regular increase in many of the properties of the halogens as we proceeddown the column from fluorine to iodine, including the melting point, boiling point, intensityof the color of the halogen, the radius of the corresponding halide ion, and the density of theelement. On the other hand, there is a regular decrease in the first ionization energy as we godown this column. As a result, there is a regular decrease in the oxidizing strength of thehalogens from fluorine to iodine.

F2 > Cl2 > Br2 > I2

oxidizing strength

This trend is mirrored by an increase in the reducing strength of the corresponding halides.

I- > Br- > Cl- > F-

reducing strength

Some Properties of F2, Cl2, Br2, and I2

Melting Boiling Colour Natural 1st Electron Ionic Density

Point(C)

Point(C)

Abundance(ppm)

IonizationEnergy(kJ/mol)

Affinity(kJ/mol)

Radius(nm)

(g/cm3)

F2 -218.6 -188.1 colourless 544 1680.6 322.6 0.133 1.513

Cl2 -101.0 -34.0 pale green 126 1255.7 348.5 0.184 1.655

Br2 -7.3 59.5 dark red-brown 2.5 1142.7 324.7 0.196 3.187

I2 113.6 185.2 very darkvioletalmost black

0.46 1008.7 295.5 0.220 3.960

Methods of Preparing the Halogens from their Halides

The halogens can be made by reacting a solution of the halide ion with any substance that is astronger oxidizing agent. Iodine, for example, can be made by reacting the iodide ion witheither bromine or chlorine.

2 I-(aq) + Br2(aq) I2(aq) + 2 Br-(aq)

Bromine was first prepared by A. J. Balard in 1826 by reacting bromide ions with a solutionof Cl2 dissolved in water.

2 Br-(aq) + Cl2(aq) Br2(aq) + 2 Cl-(aq)

To prepare Cl2, we need a particularly strong oxidizing agent, such as manganese dioxide(MnO2).

2 Cl-(aq) + MnO2(aq) + 4 H+(aq) Cl2(aq) + Mn2+(aq) + 2 H2O(l)

The synthesis of fluorine escaped the efforts of chemists for almost 100 years. Part of theproblem was finding an oxidizing agent strong enough to oxidize the F- ion to F2. The task ofpreparing fluorine was made even more difficult by the extraordinary toxicity of both F2 andthe hydrogen fluoride (HF) used to make it.

The best way of producing a strong reducing agent is to pass an electric current through a saltof the metal. Sodium, for example, can be prepared by the electrolysis of molten sodiumchloride.

electrolysis

2 NaCl(l) 2 Na(s) + Cl2(g)

In theory, the same process can be used to generate strong oxidizing agents, such as F2.

Attempts to prepare fluorine by electrolysis, however, were initially unsuccessful. HumphryDavy, who prepared potassium, sodium, barium, strontium, calcium, and magnesium byelectrolysis repeatedly tried to prepare F2 by the electrolysis of fluorite (CaF2), and succeeded

only in ruining his health. Joseph Louis Gay-Lussac and Louis Jacques Thenard, whoprepared elemental boron for the first time, also tried to prepare fluorine and suffered fromvery painful exposures to hydrogen fluoride. George and Thomas Knox were badly poisonedduring their attempts to make fluorine, and both Paulin Louyet and Jerome Nickles died fromfluorine poisoning.

Finally, in 1886 Henri Moissan successfully isolated F2 gas from the electrolysis of a mixedsalt of KF and HF and noted that crystals of silicon burst into flame when mixed with thisgas. Electrolysis of KHF2 is still used to prepare fluorine today, as shown in the figure below.

ELECTROLYSIS2 KHF2(s) H2(g) + F2(g) + 2 KF(s)

Common Oxidation Numbers for the Halogens

Fluorine is the most electronegative element in the periodic table. As a result, it has anoxidation number of -1 in all its compounds. Because chlorine, bromine, and iodine are lesselectronegative, it is possible to prepare compounds in which these elements have oxidationnumbers of +1, +3, +5, and +7, as shown in the table below.

Common Oxidation Numbers for the Halogens

OxidationNumber Examples

-1 CaF2, HCl, NaBr, AgI

0 F2, Cl2, Br2, I2

+1 HClO, ClF+3 HClO2, ClF3

+5 HClO3, BrF5, BrF6-, IF5

+7 HClO4, BrF6+, IF7

General Trends in Halogen Chemistry

There are several patterns in the chemistry of the halogens.

1. Neither double nor triple bonds are needed to explain the chemistry of the halogens.

2. The chemistry of fluorine is simplified by the fact it is the most electronegative element inthe periodic table and by the fact that it has no d orbitals in its valence shell, so it can't expandits valence shell.

3. Chlorine, bromine, and iodine have valence shell d orbitals and can expand their valenceshells to hold as many as 14 valence electrons.

4. The chemistry of the halogens is dominated by oxidation-reduction reactions.

The Hydrogen Halides (HX)

The hydrogen halides are compounds that contain hydrogen attached to one of the halogens(HF, HCl, HBr, and HI). These compounds are all colorless gases, which are soluble in water.Up to 512 mL of HCl gas can dissolve in a single mL of water at 0oC and 1 atm, for example.Each of the hydrogen halides ionizes to at least some extent when it dissolves in water.

H2O

HCl(g) H+(aq) + Cl-(aq)

Several of the hydrogen halides can be prepared directly from the elements. Mixtures of H2

and Cl2, for example, react with explosive violence in the presence of light to form HCl.

H2(g) + Cl2(g) 2 HCl(g)

Because chemists are usually more interested in aqueous solutions of these compounds thanthe pure gases, these compounds are usually synthesized in water. Aqueous solutions of thehydrogen halides are often called mineral acids because they are literally acids prepared fromminerals. Hydrochloric acid is prepared by reacting table salt with sulfuric acid, for example,and hydrofluoric acid is prepared from fluorite and sulfuric acid.

2 NaCl(s) + H2SO4(aq) 2 HCl(aq) + Na2SO4(aq)

CaF2(s) + H2SO4(aq) 2 HF(aq) + CaSO4(aq)

These acids are purified by taking advantage of the ease with which HF and HCl gas boil outof these solutions. The gas given off when one of these solutions is heated is collected andthen redissolved in water to give relatively pure samples of the mineral acid.

The Interhalogen Compounds

Interhalogen compounds are formed by reactions between different halogens. All possibleinterhalogen compounds of the type XY are known. Bromine reacts with chlorine, forexample, to give BrCl, which is a gas at room temperature.

Br2(l) + Cl2(g) 2 BrCl(g)

Interhalogen compounds with the general formulas XY3, XY5, and even XY7 are formed whenpairs of halogens react. Chlorine reacts with fluorine, for example, to form chlorinetrifluoride.

Cl2(g) + 3 F2(g) 2 ClF3(g)

These compounds are easiest to form when Y is fluorine. Iodine is the only halogen thatforms an XY7 interhalogen compound, and it does so only with fluorine.

ClF3 and BrF5 are extremely reactive compounds. ClF3 is so reactive that wood, asbestos, andeven water spontaneously burn in its presence. These compounds are excellent fluorinatingagents, which tend to react with each other to form positive ions such as ClF2

+ and BrF4+ and

negative ions such as IF2- and BrF6

-.

2 BrF5(l) [BrF4+][BrF6

-](s)

Neutral Oxides of the Halogens

Under certain conditions, it is possible to isolate neutral oxides of the halogens, such as Cl2O,Cl2O3, ClO2, Cl2O4, Cl2O6, and Cl2O7. Cl2O7, for example, can be obtained by dehydratingperchloric acid, HClO4. These oxides are notoriously unstable compounds that explode whensubjected to either thermal or physical shock. Some are so unstable they detonate whenwarmed to temperatures above -40oC.

Oxyacids of the Halogens and Their Salts

Chlorine reacts with the OH- ion to form chloride ions and hypochlorite (OCl-) ions.

Cl2(aq) + 2 OH-(aq) Cl-(aq) + OCl-(aq) + H2O(l)

This is a disproportionation reaction in which one-half of the chlorine atoms are oxidizedto hypochlorite ions and the other half are reduced to chloride ions.

When the solution is hot, this reaction gives a mixture of the chloride and chlorate (ClO3-)

ions.

3 Cl2(aq) + 6 OH-(aq) 5 Cl-(aq) + ClO3-(aq) + 3 H2O(l)

Under carefully controlled conditions, it is possible to convert a mixture of the chlorate andhypochlorite ions into a solution that contains the chlorite (ClO2

-) ion.

ClO3-(aq) + ClO-(aq) 2 ClO2

- (aq)

The last member of this class of compounds, the perchlorate ion (ClO4-), is made by

electrolyzing solutions of the chlorate ion.

The names of the oxyanions of the halogens use the endings -ite and -ate to indicate low andhigh oxidation numbers and the prefixes hypo- and per- to indicate the very lowest and veryhighest oxidation numbers, as shown in the table below. Each of these ions can be convertedinto an oxyacid, which is named by replacing the -ite ending with -ous and the -ate endingwith -ic.

Oxyanions and Oxyacids of Chlorine

Oxyanions Oxyacids

Oxidation Stateof the Chlorine

Compound Name Compound Name

+1 ClO- hypochlorite HClO hypochlorous acid

+3 ClO2- Chlorite HOClO chlorous acid

+5 ClO3- chlorate HOClO2 chloric acid

+7 ClO4- perchlorate HOClO3 perchloric acid

All halogens are nonmetals with the general formula X2, where X denotes a halogen element(Fluorine, Chlorine, Bromine, Iodine, and Astatine). The halides are the anions derived fromthe halogens. The halides of the alkali metals are all water soluble salts. Precipitationtitrimetry is based on reactions that yield ionic compounds of limited solubility. By far themost widely used and most important precipitating reagent is silver nitrate, which is used forthe determination of the halides, the halide-like anions (SCN-, CN-, CNO-), mercaptans:-.,fatty acids, and several divalent and trivalent inorganic anions. Titrimetric methods based onsilver nitrate are sometimes called argentometric methods. There are three methods that areemployed in argentometric methods. These are;

1. Mohr’s method- uses chromate ions as indicators.2. Fajan’s method- uses iron (II) and iron (III) ions as indicators3. Volhard’s method- uses titration with thiocyanate solution.

As a condition, the two halides salts should either have a common anion or a common cation.In the Mohr Method, the sodium chromate can serve as an indicator for the argentometricdetermination of chloride, bromide and cyanide ions by reacting with silver ion to form abrick-red silver chromate (Ag2CrO) precipitate in the equivalence-point region.

CHEMICALS AND EQUIPMENT

1. Mixture of NaCl and KCl2. Distilled water3. Calcium carbonate4. 5% potassium chromate5. 0.1M standard silver nitrate solution.6. Burette7. 5ml measuring cylinder8. 250ml conical flask9. 100ml measuring cylinder10. Beaker11. Spatula

PROCEDURE

PROCEDURE OBSERVATION0.3g of the mixture was weighed into abeaker and dissolved in water. It wastransferred into a 100ml measuring cylinderand topped up to the mark.

A colorless solution of the mixture wasformed.

A pinch of calcium carbonate was added tothe solution.

There was no effervescence and the solutionchanged to a white.

2ml of 5% potassium chromate wasdissolved in 100ml of distilled water and thena pinch of calcium carbonate was added toprepare an indicator blank.

A deep brown solution was formed.

2ml of 5% potassium chromate was added tothe sample solution and titrated againststandard silver nitrate solution.

The solution changed from yellow to deepbrown at the end point.

The experiment was repeated for two more values.

TABLE OF RESULTS

Indicator used: 5% potassium chromate (K2Cr2O7)Amount of indicator used: 2mlColour change: Lemon green to brick red

Burette readings/mL Final reading/mL Initial reading/ mL Titre value/ mL1 33.60 21.60 12.002 14.60 0.00 14.603 14.80 0.00 14.48

Average titre = (12.00+14.60+14.48 ) = 13.69 cm3

3

CALCULATION AND EVALUATION OF DATA

Calculating the masses of the NaCl and the KCl

Let the mass of NaCl = Xgthe mass of KCl = Ygthe mass of mixture sample = W1

the mass of standard silver nitrate solution = W2

from X + Y = W1

X +Y = 0.1g ------- 1169.87X + 169.87Y = W2 ------- 258.44 74.55

from C = n/Vwhere C= amount concentration of the standard silver nitrate

n= moles of standard silver nitrateV= volume of standard silver nitrate

n= C × V= 0.1×13.69=1.37molM(AgNO3)= 107.87 + 14 + 3(16) = 169.87g/molm(AgNO3)= 169.87×1.37/1000=0.232ghence from eq. 2169.87X + 169.87Y=0.232g58.44 74.552.91X + 2.28Y = 0.232g ------- 3from eq.1X= 0.1−Y ------- 4putting eq.4 into eq.32.91(0.1−Y) + 2.28Y= 0.232g0.291−2.91Y + 2.28Y= 0.232g-0.63Y= -0.059gY=0.094gputting Y=0.094g into eq.4X= 0.1-0.094g= 0.006g

Calculating the moles of the NaCl and the KCl

Let n(x) = moles of NaCl present in the samplen (y)= moles of KCl present in the sample

n(x) + n(y)= n(T),[ the moles of the mixture] ------ 1n(AgNO3)x + n(AgNO3)y = nT(AgNO3) ------- 2i.e. the sum of the amount of AgNO3 that reacted with both salts should be equal to the totalamount of AgNO3 that reacted.nT (AgNO3)= n(T)This implies thatn(AgNO3)x + n(AgNO3)y = n(T)m(AgNO3)x + m(AgNO3)y = n(T) ----- 3M(AgNO3) M(AgNO3)but the mass AgNO3 that reacted with NaCl = mass of NaCl

the mass of AgNO3 that reacted with KCl= mass of KClhence X + Y = n(T)

M(AgNO3)from eq.1X + Y = X + YM(AgNO3) M(NaCl) M(KCl)ButM(AgNO3) = 169.87g/molM(NaCl) = 58.44g/molM (KCl) = 74.55g/molhence X + Y = X + Y

169.87 58.44 74.55X + Y = 169.87X + 169.87Y

58.44 74.55

DISCUSSION

From the analysis above, it could be observed that the mass of both the KCl and the NaClwas found to be positive. This is an overt indication that there was a higher level of accuracyin weighing masses of the samples.

From the masses obtained it can be seen that the mass of the sample should be equal to themass of the silver nitrate produced. Thus, a decrease in one of the masses would lead to anegative mass and vice versa hence the result

PRECAUTION

1. Accurate measurement s was made.2. Low concentration of the chromate ions was used since a high concentration of the

chromate ions impart such an intense yellow color to the solution that formation of thered silver chromate is not readily detected. This could affect the titration.

ERROR ANALYSIS

Colloidal AgCl is sensitive to photodecomposition particularly in the presence of theindicator.

CONCLUSION

The actual masses of the constituents of the mixture could be determined because accuratemeasurements were made.

REFERENCES

Gary D. Christian Analytical chemistry sixth edition ( page 32-60) Practical Chemistry laboratory manual, KNUST (Year 2) pages 5-6. Statistics for Analytical Chemistry, J.C. Miller and J.N. Miller, 3rd edition, (page 16-20) Vogel’s hand text book of Quantitative chemical analysis ,fifth edition, G.H. Jeffery,

J.Bassett, J.Mendham.(pg 45-45)