the equilibrium constant (K) is a quantitative measurement of the
extent to which a chemical or physical reaction proceeds to completion. If two reactants of
known initial concentration are mixed and then the concentrations of the products are
measured once equilibrium has been established, then the equilibrium constant can be
calculated using an ICE table.
In this experiment you will measure the equilibrium constant for the formation of iron(III)
thiocyanate from the iron(III) cation and thiocyanate anion:
Fe3+(aq) + SCN-(aq) FeSCN2+(aq)
The product, FeSCN2+, is intensely red, which means that its concentration can be measured by
using an ultraviolet-visible (UV-vis) spectrophotometer. So long as a solution is not too
concentrated, the absorption of light at a particular wavelength by a chemical species is
described by Beer’s Law:
A = εℓ[X] where A = absorption of light (unitless number)
ε = molar absorptivity constant (cm-1 M-1)
ℓ = path length (cm)
[X] = concentration (M)
The path length for our cuvettes is 1.0 cm, and A will be what we record from the
spectrophotometer, so if we know ε, we can calculate the concentration of FeSCN2+. However,
we don’t know ε, so it needs to be experimentally determined first. The simplest way to
determine a molar absorptivity constant would be to take a known mass of pure solid FeSCN 2+
and dissolve it in water to a known volume of solution, and then measure the absorptivity of
this solution of known concentration. But there’s a problem with this method – as soon as any
solid FeSCN2+ is placed in water, some of it would decompose through the same equilibrium
reaction that we are hoping to study. We can bypass this problem, however, by taking
advantage of Le Chatelier’s Principle. If our thiocyanate solution is mixed with a huge excess of
iron (III) ions, then Le Chatelier’s Principle predicts that the equilibrium will shift so as to reduce
the iron concentration. In the process, nearly all of the thiocyanate will be converted into the
FeSCN2+ product, allowing its concentration to be calculated stoichiometrically without the need
for a Kc value or an ICE table

This requires creating a standard solution in which the molarity of the complex is known, and measuring its absorbance. Below are three ways of making solutions containing the FeSCN⁺² complex. For each one, explain whether or not it will result in a usable standard solution, and why. (Assume appropriate counterions are chosen that do not interfere with the chemistry or spectroscopy.) For any procedure that does work, what is the molarity of FeSCN⁺² complex?

The first part of this lab involves determining the molar absorptivity constant for the FeSCN⁺² complex. This requires creating a standard solution in which the molarity of the complex is known, and measuring its absorbance. Below are three ways of making solutions containing the FeSCN⁺² complex. For each one, explain whether or not it will result in a usable standard solution, and why. (Assume appropriate counterions are chosen that do not interfere with the chemistry or spectroscopy.) For any procedure that does work, what is the molarity of FeSCN⁺² complex?

0.500 moles of FeSCN⁺² in 1.00 L of water.

0.500 moles of Fe⁺³ and 0.500 moles of SCN^-1 in 1.00 L of water.

0.500 moles of Fe⁺³ and 0.00500 moles of SCN^-1 in 1.0 L of water.

Question

the equilibrium constant (K) is a quantitative measurement of the
extent to which a chemical or physical reaction proceeds to completion. If two reactants of
known initial concentration are mixed and then the concentrations of the products are
measured once equilibrium has been established, then the equilibrium constant can be
calculated using an ICE table.
In this experiment you will measure the equilibrium constant for the formation of iron(III)
thiocyanate from the iron(III) cation and thiocyanate anion:
Fe3+(aq) + SCN-(aq) FeSCN2+(aq)
The product, FeSCN2+, is intensely red, which means that its concentration can be measured by
using an ultraviolet-visible (UV-vis) spectrophotometer. So long as a solution is not too
concentrated, the absorption of light at a particular wavelength by a chemical species is
described by Beer’s Law:
A = εℓ[X] where A = absorption of light (unitless number)
ε = molar absorptivity constant (cm-1 M-1)
ℓ = path length (cm)
[X] = concentration (M)
The path length for our cuvettes is 1.0 cm, and A will be what we record from the
spectrophotometer, so if we know ε, we can calculate the concentration of FeSCN2+. However,
we don’t know ε, so it needs to be experimentally determined first. The simplest way to
determine a molar absorptivity constant would be to take a known mass of pure solid FeSCN 2+
and dissolve it in water to a known volume of solution, and then measure the absorptivity of
this solution of known concentration. But there’s a problem with this method – as soon as any
solid FeSCN2+ is placed in water, some of it would decompose through the same equilibrium
reaction that we are hoping to study. We can bypass this problem, however, by taking
advantage of Le Chatelier’s Principle. If our thiocyanate solution is mixed with a huge excess of
iron (III) ions, then Le Chatelier’s Principle predicts that the equilibrium will shift so as to reduce
the iron concentration. In the process, nearly all of the thiocyanate will be converted into the
FeSCN2+ product, allowing its concentration to be calculated stoichiometrically without the need
for a Kc value or an ICE table

This requires creating a standard solution in which the molarity of the complex is known, and measuring its absorbance. Below are three ways of making solutions containing the FeSCN⁺² complex. For each one, explain whether or not it will result in a usable standard solution, and why. (Assume appropriate counterions are chosen that do not interfere with the chemistry or spectroscopy.) For any procedure that does work, what is the molarity of FeSCN⁺² complex?

The first part of this lab involves determining the molar absorptivity constant for the FeSCN⁺² complex. This requires creating a standard solution in which the molarity of the complex is known, and measuring its absorbance. Below are three ways of making solutions containing the FeSCN⁺² complex. For each one, explain whether or not it will result in a usable standard solution, and why. (Assume appropriate counterions are chosen that do not interfere with the chemistry or spectroscopy.) For any procedure that does work, what is the molarity of FeSCN⁺² complex?

0.500 moles of FeSCN⁺² in 1.00 L of water.

0.500 moles of Fe⁺³ and 0.500 moles of SCN^-1 in 1.00 L of water.

0.500 moles of Fe⁺³ and 0.00500 moles of SCN^-1 in 1.0 L of water.

Accepted Answer

Given:
0.500 moles of FeSCN+2 in 1.00 L of water
0.500 moles of Fe+3 and 0.500 moles of SCN-1 in…