Discussion-equilibrium lab from lab flow Dr. Rashmi PDF

Preview

Summary

discussion for the 3rd lab on equilibrium through lab flow with dr rashmi...

Description

Discussion This lab began with a reference test tube with the colourless solution KSCN, in this solution Fe (NO3)3 was added. This addition resulted in a colour change which caused the solution to have a transparent orange hue. This is due to the formation of the iron complex FeSCN2+. This complex has red colour and the darker this colour gets; the higher concentration is present withing the solution. Therefore, in this first reaction there was not a high concentration of the iron (III) thiocyanate complex as the colour was not intense. The equation of this reaction is as follows. Fe3+(aq) + SCN−(aq) ⇄ FeSCN2+(aq) Once this step was complete, the first test tube was considered a reference and KSCN was added to test tube #2. This addition causes the solution to become more red/brown than the original solution. This is caused due to a disassociation that occurs during the reaction. The compounds ions disassociate into K+ and SCN-, the SCN- ion cause there to be more reactants present. According to Le Chatelier’s Principle, if a dynamic equilibrium is disturbed, the position of the equilibrium will shift to offset the change[1]. As a result when the KSCN was added to the reaction, the equilibrium was forced to move to the products side in order to maintain the balance due to the increase in reactants and subsequently increased the concentration of FeSCN2+. Similarly, when more Fe (NO3)3, the ions disassociated causing a surplus in reactants, forcing the dynamic shift in the equilibrium. However, when NaF was added to the solution, the colour faded and became reactant favoured. This is due the removal of Fe3+ ions as Fe3+ and F form FeF. Due to the removal, the products had a higher concentration and the equilibrium shifted to the left. When these test tubes were subjected to heat, the solution was less red/brown than the reference, showing that it is reactant favoured. On the contrary, when placed in an ice bath, the solution became a darker red/brown indicating a product favoured reaction. From this information, a conclusion can be drawn that this reaction is exothermic, and that heat can be considered a product. The second part of this reaction focused on the dissolution of tetrachlorocobaltate(II) complex ion, [CoCl4]2-, in alcohol with water. The [CoCl4]2 ion originally has a blue colour, while [Co(H2O)6]2+ has a pink colour. Similar to the first part of this experiment, if the solution is bluer there is a higher concentration of [ CoCl4]2- and if it is more pink then it has a higher concentration of [Co(H2O)6]2+. The equation for this reaction is as follows. [CoCl4] 2-(alc) + 6 H2O (alc) ⇄ [Co(H2O)6]2+ (alc) + 4 Cl−(alc)

In this reaction the addition of HCl caused the test tube to become bluer, the HCl was able to react with the Cl- ions to form a stronger product side. The equilibrium then had to shift to make up for the imbalance making this addition reactant favoured. The addition of DI water and AgNO3, shifted the equilibrium to the product side as the H2O increased with the addition of water. Furthermore, the silver ions were able to form a precipitate with chlorine which decreases the concentration of chlorine on the product side. This equation is shown below. Ag+(aq) + Cl-(aq) → AgCl(s) This reaction can also be considered a exothermic reaction, as when placed in a hot water and ice bath, the solutions reacted the same as the ones in the previous part. In both these parts of the experiment, the manipulation of the concentration of products and reactants will not affect the equilibrium constant or the K value. This is due to the fact that, the concentrations used are not based off of the starting concentrations allowing the equilibrium constant to actually be a constant value [2]. However, what can change the equilibrium constant is the temperature because the definition of equilibrium is a condition resulting from rates of forward and reverse reactions being equal. When the temperature changes, there will also be a change in the reaction rates causing a change in the K value [3]. In both parts of this experiment the reaction was exothermic, because as the temperature increased, the value of K would decrease as heat is acting as a product and vice versa for the ice bath. There was also the use of buffer solutions in this experiment. The buffer solutions used, contained (CH3COOH) and (CH3COONa or CH3COO-). A buffer is a solution that is able to resist changes in pH values when acidic or basic solutions are added. It is also bale to neutralize small amounts of acids and bases, allowing it to maitian a stabe pH of the solution[4]. Buffer solutions also have a working pH range and capacity that allows for the buffer to dictate how much acid or base can be neutralized before the pH changes. For this experiment, the buffer capacity was determined by the amount of acid and base that could be added before the change in pH was more than 1.0. While determining buffer capacities in this experiment, it can be observed that when the concentration of CH3COOH was higher than the concentration of CH3COO−, the acid buffer capacity was 10 drops of acid; since CH3COOH is a neutral ion, when the concentration of it is higher than that of CH3COO−, it is easier for the hydrogen ions (dissociated from HCl) to make the solution acidic. However, if the concentration of CH3COO− was higher, the H+ ions would react with CH3COO− to create CH3COOH, a neutral/ non-acidic ion.

Reference:

[1] Libretexts. (2020, August 15). Le Chatelier's principle. Chemistry LibreTexts. Retrieved November 10, 2021, from https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_ Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Equilibria/Le_Chat eliers_Principle. [2] EQUILIBRIUM CONSTANTS and LE CHATELIER'S PRINCIPLE. Equilibrium constants and changing conditions. (n.d.). Retrieved November 10, 2021, from https://www.chemguide.co.uk/physical/equilibria/change.html.

[3] Libretexts. (2020, August 15). The effect of changing conditions. Chemistry LibreTexts. Retrieved November 10, 2021, from https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_ Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Equilibria/Le_Chat eliers_Principle/The_Effect_of_Changing_Conditions. [4] Libretexts. (2020, August 15). Buffers. Chemistry LibreTexts. Retrieved November 10, 2021, from https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_ Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Acids_and_Bases/B uffers....