CHEM 26.1 REVIEW (FINALS)

Experiment 6. Complexometric EDTA titration Water hardness = measures the total metal ions present in water which is primarily calcium and magnesium ions EDTA, Ethylenediaminetetraacetic acid (H4Y) = hexadentate and tetraprotic ligand which forms a stable, soluble,

stoichiometric, 1:1 complex with many metals Ca2+ + Y4- CaY2-, hence at equivalence point: moles Ca2+ = moles EDTA CaCO3 is used as primary standard for standardization of EDTA

Mstock CaCO3 = (g CaCO3/FW CaCO3)(%purity of primary std.)/L solutionMworking std = (Mstock)(Vstock)/Vworking std

MEDTA = (Mworking std)(Vworking std)/VEDTA

Eriochrome Black T (H3In) = an organic dye with acid-base indicator properties; it undergoes color changes when a proton is added to and from the dye and can also perform complexes with metalsH2In- (Red) HIn2- (Blue) In3-(Orange) pH=6.3 pH =11.5HIn2-(Blue) Min-(Red) + H+

pH must be buffered at 10 to get the desired indicator color change (EBT must be in the form HIn2-). Higher pH favors complexation since it is the deprotonated form of EDTA which binds the metal but too high a pH can precipitate the metal hydroxides (at pH = 13, Ca2+ ions can be preferentially titrated since Mg will precipitate as Mg(OH)2)

Calculations:(1) MEDTA = (Mworking std)(Vworking std)/VEDTA

(2) T = MEDTAxFWCaCO3

(3) ppm CaCO3 = (TxmL EDTA)/Vsample

to compare with theoretical: ppm Ca2+ convert to ppm CaCO3

* (ppm Ca2+)(1g/1000mg)(1mol Ca/40g Ca)(1mol CaCO3/1mol Ca)(100g CaCO3/1mol CaCO3)(1000mg/1g)=> ppm Mg2+ convert to ppm CaCO3

*(ppm Mg2+)(1g/1000mg)(1mol Mg/24g Mg)(1mol MgCO3/1mol Mg)(1molCaCO3/1mol MgCO3) (100g CaCO3/1mol CaCO3 )(1000mg/1g) MgCl2•6H2O crystal is added to EDTA to give a sharp endpoint by ensuring the presence of Mg2+ for satisfactory

indicator action. NaOH is added to dissolve the disodium salt of EDTA (Na2H2Y) completely in water. Trend: (Stability of Complexes) CaIn-<MgIn-<MgY2-<CaY2-

Buffer is added to maintain the pH of the solution. NH3-NH4+ buffer is used since we are working with alkaline

condition. Adding too much buffer can cause a less sharp endpoint since NH3 also complexes with metal ions. A violet endpoint indicated the presence of Fe ions. This can be avoided by adding KCN after the buffer is added.

Experiment 7: Winkler Method Indicator used is soluble starch. The deep blue color that develops in the presence of iodine is believed to arise

from the absorption of iodine into the helical structure of beta-amylose. Addition of indicator is not done in the start of the titration process to avoid or prevent the extensive hydrolysis (breakdown) of starch which may give a premature end point.

Stoichiometry: 1 mol IO3- = 3mol I3

- = 6 mol S2O32-

Mthiosulfate = [(mass KIO3/MWKIO3)(6mol Na2S2O3/1mol KIO3)x%purity]/(Vol thiosulfate) x (10/100) ppm O2 = [(MthiosulfateVthiosulfate)(1 mol O2/4 mol S2O3

2-)(MWoxygen)]/(Vsample)