16 in areaction to produce ammonia, the theoretical yield is 420g. what is the p
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Question
16 in areaction to produce ammonia, the theoretical yield is 420g. what is the percent yield ifthe actual yield is 350.g? A, 83.3% B, 20.0% C, 16.7% D, 120% Chapter Eleven: Gases 1. Which of the following gases would have the greatest kinetic energy at 300 K? 2. Which gas would have the lowest velocity at 280 K? 3. Under which set of conditions will a real gas be least likely to act as an ideal gas? A. N2 B. NH3 C. Ar D. All of them would have the same kinetic energy A. Ar B. Ne C. H2 D. CO A. High temperature and high pressure C. Low temperature and high pressure B. High temperature and low pressure D. Low temperature and low pressure Under which set of conditions will a real gas be most likely to act as an ideal gas? A. High temperature and high pressure C. Low temperature and high pressure 4. B. High temperature and low pressure D. Low temperature and low pressure 5. Which gas will diffuse most rapidly? C. He A. CO2 B. N2 D. Cl2 6. A sample of a gas is held at constant temperature. If the number of moles of gas in the sample is doubled while the pressure is halved, what will happen to the volume of the gas sample? A. It will increase. B. It will decrease. C. It will not change. As the number of molecules in a gas sample increases, temperature and volume remaining constant, the pressure exerted by the gas A. increases. 7. B. decreases. C. remains the same.Explanation / Answer
16) The percent yield of a reaction is given as
Percent yield = (actual yield)/(theoretical yield)*100
The actual yield is 350. g while the theoretical yield is 420. g; therefore,
percent yield = (350. g)/(420. g)*100 = 83.33
The answer is closest to (A) and hence, (A) is the correct answer.
Gases
1) Assume all the gases to behave ideally. It is know that the kinetic energy of an ideal gas is only dependent on the temperature of the gas. All the gases are at 300 K; hence, all the gases will have the same kinetic energy. Therefore, (D) is the correct answer.
2) The kinetic energy of an ideal gas is given as
KE = 3/2*k*T
where k = Boltzmann’s constant. Again,
KE = ½*m*vrms2 = 3/2*k*T
====> vrms2 = 3*k*T/m = 3*k*(NA)*T/(m*NA)
where NA = Avogadro’s constant. Therefore,
vrms2= 3*R*T/M
====> vrms = ?3RT/M
where M is the molar mass of the gas.
Since R is a constant and all the gases are at the same temperature T, we have,
vrms ? 1/?M
The velocity of a gas is inversely proportional to the molar mass of the gas. The higher the molar mass of the gas, the lower is the velocity and viceversa.
The molar masses of the gases are in the order H2 < Ne < CO < Ar, i.e, Ar has the highest molar mass and hence, the lowest velocity. The correct answer is therefore, (A).
3) The kinetic molecular theory of gases assume that the gas particles are point masses and do not interact with each other, i.e, there are no attractive or repulsive forces between the gas particles. The only interactions of the gas particles are with the walls of the container and these interactions are responsible for the pressure of the gas. However, in reality, the gas particles interact with each other. However, at very high temperatures, the gas molecules are in random motion with large velocities and the forces of interaction between gas molecules is negligible. Moreover, at low pressures, the volume of the gas is large compared to the volume of individual particles and the gas molecules can be treated as massless particles. Thus, a real gas tends to behave as a real gas at low temperature and high pressures (when the volume cannot be neglected). Therefore, the correct answer is (C).
4) A real gas behaves like an ideal gas only when the temperatures are high and pressures are low. At high temperatures, the gas molecules move about randomly with large velocities and the interactions between molecules can be considered negligible. At low pressures, the volume of the gas molecules are negligible (Boyle’s law) and hence, the molecules can be considered as massless particles. Therefore, the correct answer is (B).
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