Complexometric Determination of Water Hardness

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CHM152LL LAB MANUAL

COMPLEXOMETRIC DETERMINATION OF WATER HARDNESS

Complexometric Determination of Water Hardness
Introduction
Complex ions When a neutral molecule or anion (a Lewis base) donates electron pairs and attaches itself to a metal ion center (a Lewis acid), the resulting cluster, or complex, of atoms becomes a single complex ion. When such complexes form, the electron donating groups (called ligands) form coordinate covalent bonds through empty orbitals on the metal ion. An example appears in Figure 1, where the copper(II) ion complexes with four ammonia molecules (the electron donor ligands) to form the square planar, copper-ammine complex cation. NH3 | H3N— Cu2+ —NH3 | NH3 Figure 1. The complex ion tetraamminecopper(II) [Cu(NH3)4]2+. Some polyatomic ligands have multiple lone pairs of electrons available for bonding to the central metal ion. When such a ligand (with more than one binding site) forms a complex with a metal ion, we call the process chelation, and the ligand used in the complex the chelating agent. Chemists often employ chelation to make the metal more soluble–or less soluble–in a solvent of choice. For example, toxic heavy-metal waste cleanup and remediation uses a selected chelating agent to form complexes with heavy-metal ions, which eases the metal’s removal from soil. Doctors use a similar method to chelate metal toxins in the human body for therapeutic applications. A common chelating agent is ethylenediaminetetraacetic acid (EDTA). EDTA acts as a chelating agent because each nitrogen and one oxygen from each of the four carboxylic acid groups has an electron pair to donate to a metal ion center, making EDTA a hexadentate (“six-toothed”) ligand which forms an octahedral complex.

Fig. 2. Ethylenediaminetetraacetic acid (EDTA).

Fig. 3. EDTA4– chelating a calcium ion.

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COMPLEXOMETRIC DETERMINATION OF WATER HARDNESS

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CHM152LL LAB MANUAL

COMPLEXOMETRIC DETERMINATION OF WATER HARDNESS

The Hardness of Water Water from rainfall picks up impurities by dissolving salts present in the soil. These salts contain ions of sodium, magnesium, calcium, iron and other metals. When this water evaporates or boils away, the difficult-to-dissolve metal salts remain as a scaly residue, sometimes also reacting with soap molecules to form an insoluble “scum”. If you’ve ever cleaned a shower stall or a coffee pot, you’re familiar with these deposits and how difficult they can be to remove. Water’s hardness arises from the presence of metal ions–particularly metal ions with a charge of +2 or higher. Calcium ions typically make the most significant contribution to water hardness, so by convention we report hardness as mg CaCO3/L of solution, as if all the hardness came from calcium carbonate. You will also see hardness reported in units of parts per million (or ppm), since one mg of solute has one millionth of the mass of a liter of water or dilute aqueous solution. For practical purposes, we consider water with a hardness value less than 60 ppm “soft” water, while water with more than 200 ppm CaCO3 (or equivalent) we consider “hard”. Complexometric/Chelometric Titration with EDTA In today’s lab, you will use the disodium salt of EDTA to determine the concentration of M2+ metal ion impurities in hard water by a complexometric or chelometric titration. An indicator called Eriochrome Black T will enable you to detect when your EDTA has completely chelated the metal impurities. In the presence of a metal cation, Eriochrome Black T forms a pink complex (see Equation 1). H2In- represents the anion of the free, solvated indicator, and M2+ represents Mg2+ or Ca2+. H2In–(aq) + M2+(aq) ⇌ MIn–(aq) + 2H+(aq) Equation 1 blue pink

As you add EDTA solution from a buret, the metal ions preferentially complex to the EDTA, leaving the indicator solvated, as in Equation 2. EDTA4–(aq) + MIn–(aq) + 2H+(aq) → H2In–(aq) + MEDTA2–(aq) pink blue

Equation 2

Consequently, a color...
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