With the permission of your instructor, carry out the...

With the permission of your instructor, carry out the following experiment. In a beaker, mix equal volumes of $0.001 \mathrm{M} \mathrm{NH}_4 \mathrm{SCN}$ and $0.001 \mathrm{M} \mathrm{FeCl}_3$ (the latter solution must be acidified with concentrated $\mathrm{HNO}_3$ at a ratio of 4 drops/L to prevent the precipitation of $\left.\mathrm{Fe}(\mathrm{OH})_3\right)$. Divide solution in half, and add solid $\mathrm{KNO}_3$ to one portion at a ratio of $4 \mathrm{~g}$ per $100 \mathrm{~mL}$. Compare the colors of the two solutions (see Color Plate 3), and explain why they are different. The relevant reaction is $$ \mathrm{Fe}^{3+}(a q)+\mathrm{SCN}^{-}(a q) \rightleftharpoons \mathrm{Fe}(\mathrm{SCN})^{2+}(a q) $$

With the permission of your instructor, carry out the following experiment. In a beaker, mix equal volumes of $0.001 \mathrm{M} \mathrm{NH}_4 \mathrm{SCN}$ and $0.001 \mathrm{M} \mathrm{FeCl}_3$ (the latter solution must be acidified with concentrated $\mathrm{HNO}_3$ at a ratio of 4 drops/L to prevent the precipitation of $\left.\mathrm{Fe}(\mathrm{OH})_3\right)$. Divide solution in half, and add solid $\mathrm{KNO}_3$ to one portion at a ratio of $4 \mathrm{~g}$ per $100 \mathrm{~mL}$. Compare the colors of the two solutions (see Color Plate 3), and explain why they are different. The relevant reaction is
$$
\mathrm{Fe}^{3+}(a q)+\mathrm{SCN}^{-}(a q) \rightleftharpoons \mathrm{Fe}(\mathrm{SCN})^{2+}(a q)
$$

Key Concepts

Chemical Equilibrium

Chemical equilibrium refers to the state in a reversible reaction where the rates of the forward and reverse reactions are equal, resulting in constant concentrations of the reactants and products. This principle is foundational in understanding how reactions respond when external conditions such as concentration, pressure, or temperature are changed, and it is used to predict the direction of the shift according to changes made in the system.

Le Chatelier's Principle

Le Chatelier's Principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract the change. This concept is critical for predicting how the addition of substances, like an inert salt or a common ion, will shift the equilibrium position to either favor the formation of products or reactants, thereby affecting observable properties such as color.

Ionic Strength

Ionic strength is a measure of the total concentration of ions in solution, taking into account both their concentrations and charges. Increased ionic strength, achieved by adding salts that do not directly participate in the chemical reaction, can alter activity coefficients of the species involved. This change modifies the effective concentrations and therefore influences the equilibrium position, which is essential in understanding variations in reaction outcomes.

Complex Ion Formation

Complex ion formation involves the bonding of a central metal ion with one or more molecules or anions (ligands). The resulting complexes often display distinct colors due to electronic transitions. Understanding the formation constants and behavior of these complexes is important in analytical chemistry, as the color intensity of a solution can serve as an indicator of the degree of complex formation at equilibrium.

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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