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)
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 FeSCN2+ 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.
A = 0.344 (absorption of light)
ℓ = 1.0 cm
Stock solutions:
KSCN: 2.00 mM (concentration) and 3 mL (volume)
Fe(NO3)3 : 200 mM and 5 mL
Final volume: 50.00 mL
Calculations for the molar absorptivity constant for FeSCN2+: [In this section, you present the calculations that are pertinent for determining the molar absorptivity constant (ε) for FeSCN2+. Include units! . Explain the chemical principles that were involved in deciding to manipulate the numbers the way you did]
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