1) Determine the equilibrium constant, Keq, at 25°C for the reaction

2Br– (aq) + I2(s) ⇌ Br2(l) + 2I– (aq)

My wrong answer: 7.6e-10

2) Aluminum metal is formed by the electrolysis of Al2O3 in molten cryolite. How many minutes are required to form 15.6 g of Al using a current of 9 A?

My wrong answer: 620.


I NEED HELP ANSWERING THESE TWO QUESTIONS STEP BY STEP.

Sure! I'd be happy to help you step-by-step with these questions.

1) Determining the equilibrium constant, Keq:

To find the equilibrium constant for the given reaction, you need to use the concentration or pressure values of the reactants and products at equilibrium. In this case, you are given a reaction involving ions in solution, so we will use the concentrations of the ions.

Step 1: Write the balanced chemical equation:
2Br– (aq) + I2(s) ⇌ Br2(l) + 2I– (aq)

Step 2: Set up the equilibrium expression using the law of mass action:
Keq = [Br2] / ([Br–]^2 * [I2])

Step 3: Determine the concentrations of the species involved:
You have not provided any concentration values, so I cannot give you an accurate answer. Please provide the concentrations of Br2, Br–, and I2 in the given reaction.

2) Determining the time required to form aluminum:

To find the time required to form a certain amount of Al using a given current, you need to use Faraday's law of electrolysis and the concept of mole. The equation you'll use is:

moles of Al = (current × time)/(Faraday's constant)

Step 1: Write down the specific equation for the electrolysis of Al2O3:
Al2O3 → Al + O2

Step 2: Calculate the number of moles of Al:
moles of Al = mass of Al / molar mass of Al

Given information: mass of Al = 15.6 g

Molar mass of Al: You can look up the molar mass of aluminum (Al) in the periodic table, which is approximately 26.98 g/mol.

Step 3: Calculate the charge (in coulombs) required to produce the given moles of aluminum:
charge = moles of Al × Faraday's constant

Faraday's constant is the charge of one mole of electrons and has a value of approximately 96,485 C/mol.

Step 4: Determine the time required using the given current:
current = 9 A (amperes)

time = (charge) / (current)

Step 5: Substitute the values and calculate the time needed:
time = [(moles of Al × Faraday's constant) / current]

Please substitute the calculated values of moles of Al into the equation to obtain the correct time required to form 15.6 g of Al.

Note: It is important to double-check your calculations and make sure all the units are consistent. Also, keep in mind any necessary conversions between grams, moles, and ampere-seconds (coulombs) to obtain the final result.