One challenge in using hydrogen as a fuel, rather than fossil fuels for instance, is how to safely and reliably produce and store the highly flammable gas. Chemists have proposed using metal hydrides for this application. Magnesium hydride, for instance, has shown some promise. Magnesium hydride reacts with water to form magnesium hydroxide and hydrogen gas. Determine the mass of calcium hydride required to fill a 15.4 gallon fuel tank at a pressure of 20.0 atm at 23.4 °C.
Added by Michael J.
Close
Step 1
First, we need to calculate the volume of the fuel tank in liters, since the pressure and temperature are given in SI units. 15.4 gallons = 58.3 liters Show more…
Show all steps
Your feedback will help us improve your experience
Adi S and 54 other Chemistry 101 educators are ready to help you.
Ask a new question
Labs
Want to see this concept in action?
Explore this concept interactively to see how it behaves as you change inputs.
Key Concepts
Recommended Videos
The so-called hydrogen economy is based on hydrogen produced from water using solar energy. The gas is then burned as a fuel: 2H2(g) + O2(g) → 2H2O(l). A primary advantage of hydrogen as a fuel is that it is nonpolluting. A major disadvantage is that it is a gas and therefore is harder to store than liquids or solids. Calculate the volume of hydrogen gas at 28°C and 1.00 atm required to produce an amount of energy equivalent to that produced by the combustion of a gallon of octane (C8H18). The density of octane is 2.66 kg/gal, and its standard enthalpy of formation is -249.9 kJ/mol. Assume that the products of the combustion of octane are CO2(g) and H2O(l).
Adi S.
The so-called hydrogen economy is based on hydrogen produced from water using solar energy. The gas is then burned as a fuel: $$2 \mathrm{H}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}(l)$$ A primary advantage of hydrogen as a fuel is that it is nonpolluting. A major disadvantage is that it is a gas and therefore is harder to store than liquids or solids. Calculate the number of moles of $\mathrm{H}_{2}$ required to produce an amount of energy equivalent to that produced by the combustion of a gallon of octane $\left(\mathrm{C}_{8} \mathrm{H}_{18}\right) .$ The density of octane is $2.66 \mathrm{kg} / \mathrm{gal},$ and its standard enthalpy of formation is $-249.9 \mathrm{kJ} / \mathrm{mol}$.
Hydrogen has been suggested as the fuel of the future. One way to store it is to convert it to a compound that can be heated to release the hydrogen. One such compound is calcium hydride, $\mathrm{CaH}_{2}$. This compound has a heat of formation of $-186.2 \mathrm{~kJ} / \mathrm{mol}$ and a standard entropy of $42.0 \mathrm{~J} / \mathrm{mol} \cdot \mathrm{K}$. What is the minimum temperature to which calcium hydride would have to be heated to produce hydrogen at one atmosphere pressure?
Recommended Textbooks
Chemistry: Structure and Properties
Chemistry The Central Science
Chemistry
Transcript
Watch the video solution with this free unlock.
EMAIL
PASSWORD