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Ch.10 - Gases: Their Properties & Behavior
All textbooksMcMurry 8th EditionCh.10 - Gases: Their Properties & BehaviorProblem 139
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Chapter 10, Problem 139

Consider the combustion reaction of 0.148 g of a hydrocarbon with the formula CnH2n+2 in an excess of O2 within a 400.0-mL steel container. Before the reaction, the gaseous mixture had a temperature of 25.0 °C and a pressure of 2.000 atm. After complete combustion and heat loss, the products and excess O2 had a temperature of 125.0 °C and a pressure of 2.983 atm. (a) What is the formula and molar mass of the hydrocarbon? (b) What are the partial pressures in atmospheres of the reactants? (c) What are the partial pressures in atmospheres of the products and the excess O2?

Verified step by step guidance
1
Step 1: Use the ideal gas law, PV = nRT, to calculate the initial moles of gas in the container before the reaction. Convert the temperature from Celsius to Kelvin by adding 273.15 to the Celsius temperature.
Step 2: Write the balanced chemical equation for the combustion of the hydrocarbon CnH2n+2 with oxygen, which produces CO2 and H2O. Use the stoichiometry of the reaction to relate the moles of hydrocarbon to the moles of CO2 and H2O produced.
Step 3: Calculate the change in moles of gas using the ideal gas law again for the conditions after the reaction. Use the final temperature and pressure to find the total moles of gas present after the reaction.
Step 4: Determine the moles of hydrocarbon combusted by comparing the initial and final moles of gas. Use the stoichiometry of the balanced equation to find the moles of CO2 and H2O produced, and thus the value of n in CnH2n+2.
Step 5: Calculate the molar mass of the hydrocarbon using the mass given (0.148 g) and the moles of hydrocarbon determined. Use this information to identify the molecular formula of the hydrocarbon. Then, use Dalton's Law of Partial Pressures to find the partial pressures of the reactants and products.
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Textbook Question

The apparatus shown consists of three temperature-jacketed 1.000-L bulbs connected by stopcocks. Bulb A contains a mixture of H2O(g), CO2(g), and N2(g) at 25 °C and a total pressure of 564 mm Hg. Bulb B is empty and is held at a temperature of -70 °C. Bulb C is also empty and is held at a temperature of -190 °C. The stopcocks are closed, and the volume of the lines connecting the bulbs is zero. CO2 sublimes at -78 °C, and N2 boils at -196 °C.

(a) The stopcock between A and B is opened, and the system is allowed to come to equilibrium. The pressure in A and B is now 219 mm Hg. What do bulbs A and B contain?

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Textbook Question

The apparatus shown consists of three temperature-jacketed 1.000-L bulbs connected by stopcocks. Bulb A contains a mixture of H2O(g), CO2(g), and N2(g) at 25 °C and a total pressure of 564 mm Hg. Bulb B is empty and is held at a temperature of -70 °C. Bulb C is also empty and is held at a temperature of -190 °C. The stopcocks are closed, and the volume of the lines connecting the bulbs is zero. CO2 sublimes at -78 °C, and N2 boils at -196 °C.

(b) How many moles of H2O are in the system?

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Textbook Question
When solid mercury(I) carbonate, Hg2CO3, is added to nitric acid, HNO3, a reaction occurs to give mercury(II) nitrate, Hg1NO322, water, and two gases A and B: Hg2CO31s2 + HNO31aq2¡ Hg1NO3221aq2 + H2O1l 2 + A1g2 + B1g2 (a) When the gases are placed in a 500.0-mL bulb at 20 °C, the pressure is 258 mm Hg. How many moles of gas are present?
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Textbook Question
A mixture of CS21g2 and excess O21g2 is placed in a 10.0-L reaction vessel at 100.0 °C and a pressure of 3.00 atm. A spark causes the CS2 to ignite, burning it completely, according to the equation CS21g2 + 3 O21g2¡CO21g2 + 2 SO21g2 After reaction, the temperature returns to 100.0 °C, and the mixture of product gases (CO2, SO2, and unreacted O2) is found to have a pressure of 2.40 atm. What is the partial pressure of each gas in the product mixture?
416
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Textbook Question
When 10.0 g of a mixture of Ca1ClO322 and Ca1ClO22 is heated to 700 °C in a 10.0-L vessel, both compounds decompose, forming O21g2 and CaCl21s2. The final pressure inside the vessel is 1.00 atm. (b) What is the mass of each compound in the original mixture?
596
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Textbook Question

When 10.0 g of a mixture of Ca(ClO3)2 and Ca(ClO)2 is heated to 700 °C in a 10.0-L vessel, both compounds decompose, forming O2(g) and CaCl2(s). The final pressure inside the vessel is 1.00 atm. (a) Write balanced equations for the decomposition reactions.

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