Ch.10 - Gases

Problem 117a

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Chapter 10, Problem 117a

Gaseous iodine pentafluoride, IF5, can be prepared by the reaction of solid iodine and gaseous fluorine: I21s2 + 5 F21g2¡2 IF51g2 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?

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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.

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Textbook Question

Ammonia and hydrogen chloride react to form solid ammonium chloride: NH31g2 + HCl1g2¡NH4Cl1s2 Two 2.00-L flasks at 25 °C are connected by a valve, as shown in the drawing. One flask contains 5.00 g of NH31g2, and the other contains 5.00 g of HCl(g). When the valve is opened, the gases react until one is completely consumed. (a) Which gas will remain in the system after the reaction is complete?

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Textbook Question

Ammonia and hydrogen chloride react to form solid ammonium chloride: NH₃(g) + HCl(g) → NH₄Cl(s)

Two 2.00-L flasks at 25 °C are connected by a valve, as shown in the drawing. One flask contains 5.00 g of NH₃(g), and the other contains 5.00 g of HCl(g). When the valve is opened, the gases react until one is completely consumed. (b) What will be the final pressure of the system after the reaction is complete? (Neglect the volume of the ammonium chloride formed.)

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Textbook Question

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Textbook Question

Gaseous iodine pentafluoride, IF5, can be prepared by the reaction of solid iodine and gaseous fluorine: I21s2 + 5 F21g2¡2 IF51g2 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. (c) Draw the Lewis structure of IF5.

A 6.53-g sample of a mixture of magnesium carbonate and calcium carbonate is treated with excess hydrochloric acid. The resulting reaction produces 1.72 L of carbon dioxide gas at 28°C and 743 torr pressure. c. Assuming that the reactions are complete, calculate the percentage by mass of magnesium carbonate in the mixture.