In a dilute nitric acid solution, Fe^3+ reacts with thiocyanate ion ( SCN^- ) to form a dark-red

In a dilute nitric acid solution, Fe^3+ reacts with...

In a dilute nitric acid solution, $\mathrm{Fe}^{3+}$ reacts with thiocyanate ion ( $\mathrm{SCN}^{-}$ ) to form a dark-red complex: $\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{3+}+\mathrm{SCN}^{-} \rightleftharpoons$ $\mathrm{H}_{2} \mathrm{O}+\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5} \mathrm{NCS}\right]^{2+}$ The equilibrium concentration of $\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5} \mathrm{NCS}\right]^{2+}$ may be determined by how darkly colored the solution is (measured by a spectrometer). In one such experiment, $1.0 \mathrm{~mL}$ of $0.20 \mathrm{M} \mathrm{Fe}\left(\mathrm{NO}_{3}\right)_{3}$ was mixed with $1.0 \mathrm{~mL}$ of $1.0 \times 10^{-3} \mathrm{M} \mathrm{KSCN}$ and $8.0 \mathrm{~mL}$ of dilute $\mathrm{HNO}_{3} .$ The color of the solution quantitatively indicated that the $\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5} \mathrm{NCS}\right]^{2+}$ concentration was $7.3 \times 10^{-5} M$. Calculate the formation constant for $\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5} \mathrm{NCS}\right]^{2+}$

In a dilute nitric acid solution, $\mathrm{Fe}^{3+}$ reacts with thiocyanate ion ( $\mathrm{SCN}^{-}$ ) to form a dark-red complex:
$\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{3+}+\mathrm{SCN}^{-} \rightleftharpoons$
$\mathrm{H}_{2} \mathrm{O}+\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5} \mathrm{NCS}\right]^{2+}$
The equilibrium concentration of $\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5} \mathrm{NCS}\right]^{2+}$
may be determined by how darkly colored the solution is (measured by a spectrometer). In one such experiment, $1.0 \mathrm{~mL}$ of $0.20 \mathrm{M} \mathrm{Fe}\left(\mathrm{NO}_{3}\right)_{3}$ was mixed with $1.0 \mathrm{~mL}$ of $1.0 \times 10^{-3} \mathrm{M} \mathrm{KSCN}$ and $8.0 \mathrm{~mL}$ of dilute
$\mathrm{HNO}_{3} .$ The color of the solution quantitatively indicated that the $\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5} \mathrm{NCS}\right]^{2+}$ concentration was $7.3 \times 10^{-5} M$. Calculate the formation constant for $\left[\mathrm{Fe}\left(\mathrm{H}_{2} \mathrm{O}\right)_{5} \mathrm{NCS}\right]^{2+}$

Key Concepts

Chemical Equilibrium

This concept refers to the state in a chemical reaction wherein the forward and reverse reactions occur at equal rates, resulting in constant concentrations of reactants and products. Understanding equilibrium is crucial for calculating how much of each species is present in solution at a given point when a system has reached a steady state.

Formation Constant

Also known as the stability constant, the formation constant is an equilibrium constant that quantifies the formation of a complex ion from its constituent parts in solution. It is defined as the ratio of the concentration of the complex ion to the product of the concentrations of the individual reactants at equilibrium, providing insight into the stability and affinity of the complex formation.

Reaction Stoichiometry

Stoichiometry involves the quantitative relationships between reactants and products in a chemical reaction, dictated by the balanced chemical equation. It allows for the determination of how many moles of a product are formed from given moles of reactants, which is essential when setting up equilibrium calculations.

Dilution and Concentration Calculations

When multiple solutions are mixed, the total volume changes, which affects the concentration of each solute. Dilution calculations are used to determine the new concentrations after mixing, and these altered concentrations are then utilized in equilibrium expressions to evaluate the system's behavior.

Spectrophotometric Analysis

Spectrophotometry is an analytical technique that measures the intensity of light absorbed by a solution to determine the concentration of a particular species. The color intensity, often related through the Beer-Lambert law, provides a quantitative means of assessing the concentration of colored complexes in equilibrium studies.

Transcript

00:01 In this question we're going to be looking into the dissociation, that is, we have f .e .h2o 63 plus, that is reacting with scn to form f .e .h2o, 5ncs2 plus and h2o.

00:25 So what we want to determine here is the equilibrium constant, which is kf.

00:30 And this is the concentration of the products divided by the concentration of the reactants and the products that come into the kf expression is this ion that we are forming here and not h2o so we're going to be having the concentration of let's just call this f and this is s and this is let's just call this h so this is going to be the concentration of f divided by the concentration of h multiplied by the concentration of s and all of these are evaluated at equilibrium so the main task here will be to determine the concentrations that we're going to be having at equilibrium so we're going to have the initial and the change and we're going to finally calculate the concentration at equilibrium by adding the initial and the respective changing so for this one the initial concentration is 0 .02 and the initial concentration here is 1 .0 by 10 to the power negative form so before the reaction actually starts we won't be having any of this we are ignoring this part when we are dealing with this concentration so looking at the change if this changes by negative if this changes by 7 .3 by 10 to the power negative 5 it also means that this is also going to change by 7 .3 by 10 to the power negative 5 we are looking at the ratio of moles here one more of each is forming one more of this so looking at the stoichiometric coefficients, that is 1 is to 1.

02:08 If these two change by this factor, it means this is also going to increase by that factor of 7 .3 by 10 to the power negative 5.

02:18 So what we are going to have at equilibrium, this is going to be 7 .3 by 10 to the power negative 5...

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