Question

2. Ethylenediaminetetraacetic acid (EDTA) is used for a range of applications because of its ability to strongly bind metal ions via its six ligating groups. EDTA is used in the paper industry to bind Mn2+ (and other metal) ions from catalyzing the decomposition of hydrogen peroxide (H2O2), which is an important chemical for bleaching the pulp. In this example, we are going to focus on the chelation of free Mn2+ ions in a solution at pH = 7. a) If there are 20.0 L of 0.00025 M Mn2+ in the solution before the addition of EDTA, how much EDTA is needed to complex all of the Mn2+ in moles? b) You have a 0.050 M solution of EDTA, how many L of the solution is needed to reach the equivalence point for the solution of Mn2+ described in part a? c) What is the concentration of Mn2+ ions at the equivalence point? d) Instead of just going to the equivalence point, excess EDTA is added to further reduce the concentration of metal ions and to ensure that the H2O2 is available to bleach the pulp. If 0.20 L of EDTA is added (in total), what will the concentration of free Mn2+ be? e) How much total EDTA would be needed to reduce the concentration of Mn2+ to less than 5*10-12 M? Give this answer in L of 0.050 M EDTA to 1 significant figure. [Compared to what you expected, is this a lot? Think about the titration curve, where on the titration curve are you? Does this make sense?]

Answer 1

Part a.

The moles of EDTA needed = 20 L * 0.00025 mol/L = 0.005 mol

Part b.

The volume of EDTA needed = 0.005 mol/(0.05 mol/L) = 0.1 L

Part c.

The concentration of bound EDTA-bound Mn²⁺ = 0.005 mol/(20+0.1) L = 2.49*10^-4 M

Part d.

The concentration of free Mn²⁺ = 0 M

Part e.

The volume of EDTA needed = 0.005 mol/(5*10^-12 mol/L) = 1*10⁹ L