Gaseous iodine pentafluoride, IF 5 , can be prepared by the reaction of solid iodine and gaseous fluorine: I 2 (s) + 5 F 2 (g) → 2 IF 5 (g) A 5.00-L flask containing 10.0 g of I 2 is charged with 10.0 g of F 2 , and the reaction proceeds until one of the reagents is completely consumed. After the reaction is complete, the temperature in the flask is 125 °C. (a) What is the partial pressure of IF5 in the flask?

Hi everyone here we have a question asking us to consider the reaction bro. Mean liquid plus chlorine gasses forms bromine chloride gashes and we're told that a 2.0 liter flask containing 5.0 g of roaming is charged with 2.50 g of chlorine. After the reaction is complete, the temperature in the flask is 150 degrees Celsius, assuming that the reaction proceeds until one of the re agents is completely consumed. Calculate the partial pressure of bromine chloride in the flask. So first we need a balanced reaction. So we have browning liquid plus chlorine. Gashes forms roaming chloride gashes And if we look at our reactive side, we see that we have to browning and only one on our product side. So we're going to put a two in front of their. Now we have 10 chlorine on our product side. So we're going to put a five in front of our chlorine on a reacting side. To balance that out. So now we have 10 on both sides. So now we can calculate the molds of bromine chloride produced using both react ints. So we need the molar mass. So for browning we have to Times 79. equals 1 59.8 g per mole in chlorine. We have two tons, 35. Equal 70.9 g per mole. So now we're going to start with our 5. brands of browning. Times one mole of browning Divided by .8 g of grooming times two moles. A roaming chloride. And that is from our balanced reaction over one mole. A roaming and that equals 0.0 - six moles of browning chloride. Now if we look here are grams of brown meat are canceling out, our moles of bromine are canceling out, leaving us with moles of bromine chloride, which is exactly what we want. So now we're going to do the exact same thing with our chlorine. So we have 2.50 g of chlorine times one more of chlorine over 70.9 g of chlorine. Times two moles. A roaming chloride over five moles. The chlorine and our grams of chlorine are canceling out our molds of chlorine are canceling out and that equals 0. moles of roaming. Cool ride. And if we compare the two, we see that we get less when we use chlorine. So chlorine is our limiting reactant and it gives us our theoretical yield. So now we can calculate the pressure of roaming chloride using the ideal gas law. The ideal gas law gives us the formula PV equals NRT. Our volume equals 2.00 L. Our number of moles equals 0. moles. Our gas constant equal 0. to 06. Leaders times atmospheres over moles time kelvin and our temperature Equals 150. And that's in C. and we want to change it to Kelvin. So we're gonna add 273.15 and that equals .15 Kelvin. So now we're going to rearrange our PV equals NRT to solve for pressure. So we're going to divide both sides by B, which gives us P equals NRT over V. So P equals 0.0141 moles Times 0.08206 L times atmospheres over moles times kelvin times 423.15 Kelvin, divided by two leaders Equals 0.245 Atmospheres. And that is our final answer. Thank you for watching. Bye.

Ch.10 - Gases

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Brown 14th Edition

Ch.10 - Gases

Problem 126a

Chapter 10, Problem 126a

Gaseous iodine pentafluoride, IF₅, can be prepared by the reaction of solid iodine and gaseous fluorine: I₂(s) + 5 F₂(g) → 2 IF₅(g) A 5.00-L flask containing 10.0 g of I₂ is charged with 10.0 g of F₂, and the reaction proceeds until one of the reagents is completely consumed. After the reaction is complete, the temperature in the flask is 125 °C. (a) What is the partial pressure of IF5 in the flask?

Verified step by step guidance

Calculate the molar mass of iodine (I2) and fluorine (F2) to determine the number of moles of each reactant. The molar mass of I2 is approximately 253.8 g/mol and for F2 it is about 38.0 g/mol.

Use the molar masses to convert the mass of each reactant to moles. For I2, divide 10.0 g by its molar mass. For F2, divide 10.0 g by its molar mass.

Determine the limiting reactant by comparing the stoichiometric ratio from the balanced chemical equation with the mole ratio calculated. The balanced equation requires 1 mole of I2 for every 5 moles of F2.

Calculate the moles of IF5 produced based on the moles of the limiting reactant, using the stoichiometry of the balanced equation, which shows that 2 moles of IF5 are produced for each mole of I2 reacted.

Use the ideal gas law, PV = nRT, to find the partial pressure of IF5. Convert the temperature from Celsius to Kelvin by adding 273.15 to the Celsius temperature. Use the value of R (0.0821 L·atm/mol·K) and the volume of the flask (5.00 L) along with the moles of IF5 calculated.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Stoichiometry

Stoichiometry is the calculation of reactants and products in chemical reactions based on the balanced chemical equation. It allows us to determine the amount of product formed or reactant consumed by using mole ratios derived from the coefficients in the balanced equation. In this case, understanding the stoichiometric relationship between iodine (I2) and fluorine (F2) is essential to find out how much IF5 can be produced.

Ideal Gas Law

The Ideal Gas Law relates the pressure, volume, temperature, and number of moles of a gas through the equation PV = nRT. This law is crucial for calculating the partial pressure of gases in a mixture. In this scenario, after determining the moles of IF5 produced, the Ideal Gas Law can be applied to find its partial pressure in the 5.00-L flask at the given temperature of 125 °C.

Partial Pressure

Partial pressure is the pressure exerted by a single component of a gas mixture. According to Dalton's Law of Partial Pressures, the total pressure of a gas mixture is the sum of the partial pressures of each individual gas. In this question, calculating the partial pressure of IF5 requires knowing the total number of moles of IF5 produced and applying the Ideal Gas Law to find its contribution to the overall pressure in the flask.