Chemical Principles in the Laboratory
11th Edition
ISBN: 9781305264434
Author: Emil Slowinski, Wayne C. Wolsey, Robert Rossi
Publisher: Brooks Cole
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When limestone (solid CaCO3) is heated, it decomposes into lime (solid CaO) and carbon dioxide gas. This is an extremely useful industrial process of great antiquity, because powdered lime mixed with water is the basis for mortar and concrete — the lime absorbs CO2 from the air and turns back into hard, durable limestone.
Suppose a limekiln of volume 300.L is pressurized with carbon dioxide gas to 15.5atm, and heated to 950.0°C. When the amount of CO2 has stopped changing, it is found that 1.79kg of CaCO3 have appeared. Calculate the pressure equilibrium constant Kp this experiment suggests for the equilibrium between CaCO3 and CaO at 950.0°C. Round your answer to 2 significant digits.
Note for advanced students: it's possible there was some error in this experiment, and the value it suggests for Kp does not match the accepted value.
Butane gas, C4H10, is sold to campers as bottled fuel. Its density at 258C and 1.00 atm is 2.38 g/L. What volume of butane gas at 258C and 1.00 atm is required to heat one gallon of water (d 5 1.00 g/mL) from 258C to 988C? The reaction for the combustion of butane (DHf 8 5 2125.6 kJ/mol) is C H (g)113O (g)9:4CO (g)15H O(g)
4a
Assume that worldwide combustion of fossil fuel accounts for an energy use of 500 x 1015 kJ/year. Also, assume that all fossil fuels can be represented by the formula C3H5 with energy content of 32,500 kJ/kg.Finally, assume that air is 79% N2 and 21% O2 , and has a molecular weight of 29 kg/kmol.
i. Estimate the annual release of CO2 to the atmosphere by burning fossil fuels, in kg/yr
ii. If all that CO2 entered the atmosphere (and none was removed), estimate the increase in atmospheric concentration, in ppm by volume. Assume the atmosphere contains 6.7 x 1018 kg of air.
Chapter 8 Solutions
Chemical Principles in the Laboratory
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- At the critical point for carbon dioxide, the substance is very far from being an ideal gas. Prove this statement by calculating the density of an ideal gas in g/cm3 at the conditions of the critical point and comparing it with the experimental value. Compute the experimental value from the fact that a mole of CO2 at its critical point occupies 94 cm3.arrow_forwardWhat is the value of FP for a sample of gas whose temperature is -33.0 C and volume is 0.0250 L? What temperature is required to change the volume to 66.9 cm3?arrow_forwardButane gas, C4H10, is sold to campers as bottled fuel. Its density at 25C and 1.00 atm is 2.38 g/L. What volume of butane gas at 25C and 1.00 atm is required to heat one gallon of water (d=1.00g/mL) from 25C to 98C ? The reaction for the combustion of butane (H f =125.6kJ/mol) is C4H10(g)+132 O2(g)4CO2(g)+5H2O(g)arrow_forward
- In the equation w = P V, why is there a negative sign?arrow_forwardYou heat 1.000 quart of water from 25.0C to its normal boiling point by burning a quantity of methane gas, CH4. What volume of methane at 23.0C and 745 mmHg would you require to heat this quantity of water, assuming that the methane is completely burned? The products are liquid water and gaseous carbon dioxide.arrow_forwardIn the 1880s, Frederick Trouton noted that the enthalpy of vaporization of 1 mol pure liquid is approximately 88 times the boiling point, Tb, of the liquid on the Kelvin scale. This relationship is called Troutons rule and is represented by the thermochemical equation liquid gas H = 88 Tb, joules Combined with an empirical formula from chemical analysis, Troutons rule can be used to find the molecular formula of a compound, as illustrated here. A compound that contains only carbon and hydrogen is 85.6% C and 14.4% H. Its enthalpy of vaporization is 389 J/g, and it boils at a temperature of 322 K. (a) What is the empirical formula of this compound? (b) Use Troutons rule to calculate the approximate enthalpy or vaporization or one mole of the compound. Combine the enthalpy of vaporization per mole with that same quantity per gram to obtain an approximate molar mass of the compound. (c) Use the results of parts (a) and (b) to find the molecular formula of this compound. Remember that the molecular mass must be exactly a whole-number multiple of the empirical formula mass, so considerable rounding may be needed.arrow_forward
- 5-107 If 60.0 g of NH3 occupies 35.1 L under a pressure of 77.2 in. Hg, what is the temperature of the gas, in °C?arrow_forwardThe emission of NO2 by fossil fuel combustion can be prevented by injecting gaseous urea into the combustion mixture. The urea reduces NO (which oxidizes in air to form NO2) according to the reaction:2 CO(NH2)2( g) + 4 NO( g) + O2( g)------------->4 N2( g) + 2 CO2( g) + 4 H2O( g)Suppose that the exhaust stream of an automobile has a flow rate of 2.55 L/s at 655 K and contains a partial pressure of NO of 12.4 torr. What total mass of urea is necessary to react completely with the NO formed during 8.0 hours of driving?arrow_forwardIn the reaction 2 H2O ==> 2 H2 + O2 the DH = 286 kJ/mole. How much energy (in kJ) must you put in to generate 17.0 Liters of Hydrogen gas (at STP) ?arrow_forward
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