Analysis of CoFe2O4 by Complexometric Titration by Sushanta K Sahu

Analysis of CoFe2O4 by Complexometric Titration F. Romero Advanced Integrated Lab Chemistry 431 October 14, 1994 Abstract: A Complexometric titration was performed to analyze Fe(III) and Co(II) in a cobalt ferrite sample. Direct titration with EDTA, a polycarboxylic acid, was used to determine the amount of Fe(III) with the aid of an indicator, Tiron. The analysis of Co(II) was done by titrating excess EDTA and back titrating the solution with standard Zn2+ solution with the indicator Eriochrome Black T. The mole ratios obtained are as follows: moles of Fe3+ to moles of CoFe2O4 was 1.95, moles of Co2+ to moles of CoFe2O4 was 1.02 and moles of Fe3+ to moles of Co2+ moles was 1.91. The expected values, are moles of Fe3+ to moles of CoFe2O4 is 2, moles of Co2+ to moles of CoFe2O4 is 1, and moles of Fe3+ to moles Co2+ is 2. Complexometric titration is an accurate way to determine the amount of Fe(III) and Co(II) in the sample. INTRODUCTION: Complexometric Reaction: In this experiment we quantitatively analyzed cobalt ferrite. To accomplish this, EDTA complexes Fe(III) and Co(II) by complexometric reactions which can be used for titrimetric analyses. A complexometric titration is a quantitative reaction of a Lewis acid and a Lewis base to form a complex. The Lewis base is referred to as a ligand which has at least one lone pair of electrons to donate, while the Lewis acid is referred to as a metal ion which will receive at least a lone pair of electrons. Complexometric reactions must fulfill the same requirements as all other volumetric methods, these are: I. Complex formation reaction must be rapid. II. Must proceed according to exact stoichiometry. III. Have large equilibrium constants so the titration curve has a sharply defined equivalence point.

CHEMISTRY OF EDTA: EDTA (ethylenediaminetetraacetic acid) is used in many complexometric reactions. In fact, ethylenediaminetetraacetic acid (EDTA) is a suitable titrant for more than 95% of all analytical applications for the use of complexometric titrations. EDTA is an aminopolycarboxylic acid and is not very water soluble, therefore, the disodium salt is used to measure solutions of EDTA for complexometric titration. The pH of the disodium salt is about 4.42. EDTA is conveniently represented as H4Y and has a structure shown in Figure 1.

Figure 1 Structure of EDTA Because EDTA is tetraprotic acid, there are four ionization reactions these are: H4Y (aq) + H2O

H3O+ (aq) + H3Y- (aq) Ka1 = 1.0X10-2 (1)

H3Y- (aq) + H2O

H3O+ (aq) + H2Y2- (aq) Ka2 = 2.16X10-3 (2)

H2Y2- (aq) + H2O

H3O+ (aq) + HY3- (aq) Ka3 = 6.92X10-7 (3)

HY3- (aq) + H2O

H3O+ (aq) + Y4- (aq) Ka4 = 5.5X10-11 (4)

There are seven different species of EDTA, these depend on the pH of the solution. The dominate concentration of the EDTA species and there corresponding pH range are shown in table 1. Species 2+

H6Y

pH Range 0-1

H5Y1+

1 - 1.6

H4Y

1.6 - 2

2 - 2.7

H3Y

H2Y2HY

2.7 - 6.2 6.3 - 10.2 10.2 - Above

Table 1 EDTA species and there corresponding pH's. Ref. Skoog and West. EDTA Metal Complex:

A complex ion is a metal cation which is covalently bonded to one or more electron donating groups. Usually, the metal cation is called the central atom, where as the term ligand is used to designate the electron-donating group (EDTA in this case). In order for a complex ion to be formed from one or more ligands and a central atom, each ligand must possess at least one unshared pair of electrons, and the central atom must be able to accept an electron pair from each ligand. Thus, a ligand shares a pair of electrons with the central atom in the formation of a covalent bond. Multidentate ligands usually form much more stable complexes than do chemically similar monodentate ligands, so that well defined titration curves are obtainable. EDTA is a sexidentate ligand, it can coordinate with a metal ion through its two nitrogens and four carboxyl groups. All EDTA species will react as a hexadentate ligand with metal ions of 2+ or greater to give a one-to-one stoichiometry as shown: Mn+ (aq) + H4Y (aq)

MYn-4 (aq) + 4H+ (aq) (5)

Mn+ (aq) + H3Y1- (aq)

MYn-4 (aq) + 3H+ (aq) (6)

Mn+ (aq) + H2Y2- (aq)

MYn-4 (aq) + 2H+ (aq) (7)

Mn+ (aq) + HY3- (aq) Mn+ (aq) + Y4- (aq)

MYn-4 (aq) + H+ (aq) (8) MYn-4 (aq) (9)

Most metal cations in the periodic table form very stable, one-to-one complexes with EDTA. One-to-one stoichiometry of Metal-EDTA complexes arises from the fact that the EDTA ion possesses a total of six electron donating groups, four of which are carboxylate groups and two amine groups which can occupy four, five, or six coordination positions around a central metal atom. Metal-EDTA complexes gain particular stability from the five member planar rings which make the complex especially stable. The geometry of the metal ion and the EDTA complex is octahedral, as shown in Figure 2.

Figure 2 Complex of EDTA and Cobalt (II) We can write the general reaction for the formation of Metal-EDTA complexes as Mn+ + Y-4

MYn-4. (10)

Where Mn+ denotes a hydrated metal cation and Y-4 is the tetrasodium salt of EDTA. The reaction equilibrium expression is: (11) in which KMY is the formation constant for the MYn-4 complex. For example, table 2 has some formation constants. Cation Log Kf Ag+

7.3

16.1

7.8

Ca2+

10.7

16.3

18.8

Fe2+

14.3

25.1

Hg2+

21.8

Ni2+

3.0

Sr2+ U

18.0 8.6 25.5 16.5

Table 2 Formation constants of EDTA and metal ions. It is important to note that the larger the cation's charge the larger the KMY value. This point is very important in the titration of Fe(III). For Fe(III) is a more stable complex with EDTA in acidic solution than that of Co(II) and EDTA complex. This will help in the titration of Fe(III) with EDTA because the titration will occur in acidic solution. Therefore there will be no interference of Co(II) complexing with the EDTA. Buffers: Why is a buffer needed in a complexometric titration? First, any increase of hydrogen ions will dissociate the metal complex, for this reason it is important to use a buffer, which will keep the pH constant. Second another important reason why a buffer is used is because the indicators are acid-base indicators as well as complexometric titration indicator. It is important to have a color change for the equivalence point of the complexometric reaction and not the color change due to addition of acid to the indicator. Lastly, the buffer prevents the formation of metal hydroxides. Typically, the solution to be titrated is buffered at pH 4 to 5 using an acetic acidsodium acetate at buffer, or with an ammonium chloride-ammonia buffer at pH 9 to 10. Since the predominant form of EDTA under these conditions is HY-3 (aq) or H2Y-2 (aq), the titration process leading to the formation of a Metal-EDTA complex might be formulated as: Mn+ (aq) + H2Y2- (aq) Mn+ (aq) + HY3- (aq)

MYn-4 (aq) + 2H+ (aq) (12) MYn-4 (aq) + H+ (aq) (13)

Note that the protons liberated in these reactions could change the pH sufficiently to cause a significant amount of reverse reaction. MY2- (aq) + H+ (aq) MY2- (aq) + 2H+ (aq)

M2+ (aq) + HY3- (aq) (14) M2+ (aq) + H2Y2- (aq) (15)

The formation of hydrogen ions as a reaction product would cause the pH of the solution to decease and the desired titration reaction to cease or even reverse itself, unless the solution is buffered against the pH change. This happens because there is an equilibrium occurring in the formation of the Metal-EDTA complex, so if one amount is larger on one side (product or reactant) the reaction will proceed in the opposite direction. Therefore a buffer system such as acetic acid-sodium acetate or ammonia-ammonium chloride, is used to keep the pH at the desired value for the particular application. Indicator Species: Metachromic indicators are used to describe ligand cation complexes, because they form stable, bright colored complex with most metal ions of analytical interest. Tiron is the first indicator to help

determine the equivalence point of EDTA complexing with Fe(III). The structure of Tiron is shown in figure 3:

Figure 3 1,2- dihydroxybenzene-3,5-disulfonic acid, disodium salt. When the reaction of Fe(III) and EDTA occurs, Fe3+-Tiron

Fe3+ + Tiron (16)

Blue-Green Yellow the blue-green Fe3+-Tiron complex forms. Note, this reaction (16) is an equilibrium reaction. When the concentration of Fe3+ is decreased, the equilibrium shifts to the right and Tiron is produced to give a yellow solution. The second indicator to be used in the titration is Eriochrome Black T. Eriochrome Black T has one -SO3H group which dissociates completely, while the other two hydrogens have Ka values of: Ka HIn2- 5.00X10-7 In3- 2.82X10-12 The structure of Eriochrome Black T is:

Figure 4 Eriochrome Black T In the second half of the experiment EBT is used to determine the equivalence point in the titration with Zn(II). During the back titration with standard Zn(II), EBT is placed in the solution with excess EDTA, the reaction is : EBT (aq) + Zn(II) blue - green red

Zn(EBT) (17)

When all of the Zn(II) has complexed with the EDTA, excess Zn(II) begins to react with EBT. The equilibrium shifts to the right and Zn(EBT) forms and the color of the solution changes from Blue to Red. Titration Methods: There are two titration methods which will be used to determine the amount of Fe(III) and Co(II) in sample: direct titration and back titration. Direct titration: As the name suggests, one carries out a direct determination of a metal ion by adding standard EDTA titrant to the sample solution until an appropriate end point is observed. Usually, a buffer of sodium acetate / acetic acid is included in the sample medium to control the pH and prevent precipitation of the metal hydroxide. Fe(III) with EDTA is a direct titration which will be performed in this experiment to determine the amount of Fe(III) in the sample. Back Titration: One adds to the sample solution a known volume of standard EDTA in excess, to combine with the desired metal cation. It may be necessary to heat the solution or to alter the solution conditions so that the reaction between metal and EDTA reaches completion rapidly. Finally, the excess EDTA is back titrated under appropriate conditions with a standard solution of Zn(II) and usually a buffer of ammonium chloride, ammonia is needed. The back titration technique will be used to determine the amount of Fe(III) and Co(II) combined. EXPERIMENTAL: Materials and Apparatus: EDTA (disodium ethylenediaminetetraacetate dihydrate) Molecular Weight = 372.24 Tiron (1,2-dihydroxydenzene-3,5-disulfonic acid disodium salt), Eriochrome Black T (3-hydroxy-4-[1-hydroxy-2-naphthyl)az0]-7-nitro-1naphthalene-sulfonic acid sodium salt), Molecular Weight = 461.38 Concentrated aqueous ammonia 1:3 dilute aqueous ammonia Ammonium chloride Concentrated acetic acid Sodium acetate

6 M NaOH Zinc metal Concentrated HCl 95% ethanol Two 50 mL burets 250 and 400 mL beakers 250 and 500 mL volumetric flasks pH meter Standard pH 7 buffer Hot Plate Medicine dropper White sheet of paper Volumetric pipettes Watch glass Rubber policeman Funnels Magnetic stirrer Magnetic bar Procedure: Preparation of Buffers and Indicators: .9880 g of Tiron was dissolved in 20 mL of water in a 50 mL beaker. Eriochrome Black T (EBT) was already prepared for this experiment. The acetic / acetate buffer was prepared by 24.929 g of sodium acetate and dissolving the sodium acetate by adding up to 250.0 mL of acetic acid. The ammonia / ammonium chloride buffer was prepared already (11/5/91) by dissolving 14.015 g of ammonium chloride, 90.0 mL of aqueous ammonia, and 250.0 mL of water. Preparation of Standard 0.1 M Zn(II):

0.3267 g of reagent grade zinc solid (Zn) was dissolved by adding 6 M HCl. The Zn(II) solution was heated, then cooled, and quantitatively transferred to a 500.00 mL volumetric flask and diluted with distilled water to the calibration mark. Preparation of Standard EDTA Solution: 1.8531 g of EDTA disodium salt was dissolved in 150 mL of water. Then the solution was heated and cooled. The solution was quantitatively transferred to a 500.00 mL volumetric flask. *Excess water was placed over the volumetric flask. Approximately 1 mL of water was placed. This was determined by how many drops of solution was over the volumetric line. 20 drops = 1 mL. Preparation of the Cobalt Ferrite Sample: 0.1997 g of Cobalt Ferrite was dissolved in 5.0 mL of concentrated HCl. The solution was gently dissolved in a hot plate. Then the solution was quantitatively transferred to a 250.00 mL volumetric flask and diluted with distilled water to the calibration mark. Observation and Data: Analysis of Fe(III): 5.00 mL of Co(II) and Fe(III) solution was pipetted in to a 250.00 mL Erlenmeyer flask. 100.0 mL of distilled water was added to the flask along with 5 mL of the sodium acetate / acetic acid buffer. Then 5 to 7 drops of Tiron indicator were added to the flask and the solution turned blue-green. EDTA (9.94X10-3 M) solution was added from a buret to the flask until a color change of yellow occurred. Two other 5.00 mL portions of the same Co (II) - Fe(III) solution were titrated following the same procedure. . Table 3 shows the volumes used: Sample Run Initial Volume (mL) Final Volume (mL) Change Volume(mL) 1

0.000

16.57

0.000

16.78

0.000

16.79

Ave: 16.73 mL Table 3 Volume of EDTA used for Fe(III) Analysis. Analysis of Cobalt (II) and Fe(II): 10.00 mL of the solution of Co (II) and Fe (III) was volumetrically pipetted in to a 500.0 mL Erlenmeyer flask. 150.0 mL of ethanol was added to the flask. Immediately after the addition of ethanol, the solution turned yellow. Before the experiment I made sure that the pH meter was working properly. Then after the addition of ethanol, the pH of the solution was adjusted to two by adding 6 M HCl. EDTA was added to the flask, which was stirred with a magnetic stirrer. 40.00

mL of EDTA was added in the flask. After addition of EDTA, the yellow color disappeared. The addition of a 1:3 aqueous ammonia solution to neutralized the solution. A pH meter was used to adjust the solution pH to 7. Then a slight excess aqueous ammonia was added to make the solution slightly basic. Addition of 10 mL of ammoniaammonium chloride buffer was added which turned the pH of the solution to about 8. The solution was blue-green after Eriochrome Black T indicator was added. 9.995 X 10-3 M Zn(II) solution was titrated to the flask until a red-wine color emerged. the volumes of the Zn(II) used for the titration are in table 4: Sample Run Initial Volume (mL) Final Volume (mL) Change Volume(mL) 1

0.000

28.72

0.000

29.62

0.000

29.65

Ave: 29.66 mL Table 4 Volume of Zn(II) for the determination of Co(II) experiment. Calculations and Results: Preparation of Solutions: EDTA = 372.24 g/mole Sample of EDTA = 1.8531 g Weighted directly from the reagent bottle Moles of EDTA =

moles of EDTA

Molarity of EDTA =

= 9.94X10-3 M of EDTA

*The amount of water should have been .500 liters. When the water was added to the volumetric flask, the amount went past the calibration mark. The excess water was picked up by a pipette and the drops were counted and the excess solution was exactly one mL. CoFe2O4 = 234.623 g/mole Sample of CoFe2O4 = 0.1997 = 8.512X10-4 moles CoFe2O4

Moles of CoFe2O4 = Zn = 65.37 g/mole Sample of Zn = 0.3267 grams

Moles of Zn =

= 4.998X10-3 moles of Zn

Molarity of Zn2+ solution =

= 9.995X10-3 M of Zn2+

Analysis of Fe3+ Average amount of EDTA used in titration = 16.73 mL for a 25.00 mL sample of CoFe2O4. Moles of EDTA titrated = (9.94X10-3 M of EDTA)*(.01673 L of EDTA) Moles of EDTA titrated = 1.66X10-4 moles of EDTA used in titration for 25 .00 mL of CoFe2O4 solution EDTA complex with Fe3+ in a one-to-one reaction: Moles of EDTA = Moles of Fe3+ in 25.00 mL sample of CoFe2O4 Moles of Fe3+ = 1.66X10-4 moles Fe3+ in 25 .00 mLsample of CoFe2O4 For a 10.0 mL sample of CoFe2O4; how many moles of Fe3+?

x = 6.65X10-5 moles of Fe3+ in 10.00 mL sample of CoFe2O4 Total moles of Fe3+ in the sample solution of 250.00 mL. 25.00 mL, Fe3+ = 1.66X10-4 moles Fe3+ (25.00 mL).

x = 1.66X10-3 moles Fe3+ in 250.00 mL of solution. Analysis of Co2+: 10.00 mL sample of CoFe2O4 Solution: Average amount of EDTA used for Co2+ analysis = 40.00 mL EDTA Moles EDTA = (.04000L)*(9.94X10-3 M EDTA) = 3.976X10-4 moles of EDTA used for cobalt (II) titration. Moles of Total EDTA = 3.976X10-4 moles M used of Zn(II) to Titrate Excess EDTA = 9.995X10-3 M of Zn(II) Average Volume used for titrating of Zn(II) to EDTA = 29.66 mL Moles of Zn(II) used = (9.995X10-3 M Zn)*(.02966L) = 2.965X10-4 moles of Zn(II)

[Total moles of EDTA (Fe3+ + Co2+ + excess EDTA) ] - excess moles EDTA - Moles of Fe3+ = 3.976X10-4 - 2.965X10-4 - 6.65X10-5 = 3.460X10-5 moles of Co2+ in 10.00 mL sample of CoFe2O4 Amount of Co2+ moles in 250.00 mL Sample of CoFe2O4. = X = 8.65X10-4 moles Co2+ in total amount of solution of CoFe2O4 The results are as follows: moles of Fe3+ to moles of CoFe2O4 was 1.95, moles of Co2+ to moles of CoFe2O4 was 1.02 and moles of Fe3+ to moles of Co2+ moles was 1.91. References: 1. Fischer, Robert B. and Dennis G. Peters, Quantitative Chemical Analysis, 3rd edition, Saunders Co., Philadelphia, 1968, p. 414-448. 2. Peters, Dennis G., J.M. Hayes, and G.M. Hieftje, Chemical Separations and Measurements, Theory and Practice of Analytical Chemistry, Saunders, Philadelphia, 1974, p. 161-194. 3. Pickering, W.F., Modern Analytical Chemistry, Dekker Inc., New York, N.Y., 1971, p. 358-367. 4. Reilley and Schmid, Anal. Chem., 30, p. 947, 1958. 5. Skoog, Douglas A., and Donald M. West, Analytical Chemistry, 3rd edition, Holt, Rinehart, and Winston, New York, N.Y., 1988, p. 247-267.