Using your balanced chemical equation and stoichiometry, calculate liters of CO2 from grams of octane. Once you have that value, calculate how many 2 litre bottles would be needed to contain all the CO2 your dream vehicle produces. (You’ll need the following conversion factor since CO2 is a gas (22.4 L/ mol)
Added by Taylor P.
Step 1
The balanced equation is: \[ 2 \, \text{C}_8\text{H}_{18} + 25 \, \text{O}_2 \rightarrow 16 \, \text{CO}_2 + 18 \, \text{H}_2\text{O} \] Now, let's proceed with the calculations step by step. Show more…
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Calculate the mass of carbon dioxide (in grams) produced from the combustion of 2.50 L of pure liquid octane (C8H18, MW = 114 g/1-mole, D = 0.703 g/1-mL & M = 6.17 moles/1-L). 2 C8H18 (liq) + 25 O2 (g) ⟶ 16 CO2 (g) + 18 H2O (g) Mass of CO2 produced from 2.50 L of octane = [CO2] grams. Be sure to round your answer off to 3 significant figures.
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Calculate the theoretical amount of CO2 produced from 1.000 gallon of octane. Your final units should be lb CO2/gal fuel. The density of Octane is 0.750 g/mL; MW Octane = 114.2 g/mol; MW CO2 = 44.01 g/mol; 1.000 gal = 3.7854 L; 453.59 g = 1.000 lb. Then, calculate the theoretical amount of CO2 produced from 1.000 gallon of ethanol (lb CO2/gal). The density of ethanol is 0.789 g/mL, the molecular weight of ethanol is 46.07 g/mol, and the molecular weight of carbon dioxide is 44.01 g/mol. Why does ethanol emit less CO2 per gallon than gasoline?
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I Assume that a car burns pure octane, $$ \mathrm{C}_{8} \mathrm{H}_{18}\left(d=0.703 \mathrm{~g} / \mathrm{cm}^{3}\right) $$ (a) Write the balanced cquation for burning octanc in air, forming $\mathrm{CO}_{2}$ and $\mathrm{H}_{2} \mathrm{O}$. (b) If the car has a fuel efficiency of 32 miles per gallon of octane, what volume of $\mathrm{CO}_{2}$ at $25^{\circ} \mathrm{C}$ and $1.0 \mathrm{~atm}$ is generated when the car goes on a $10.0$ -mile trip? (The volume of $1 \mathrm{~mol} \mathrm{CO}_{2}(\mathrm{~g})$ at $25^{\circ} \mathrm{C}$ and $1 \mathrm{~atm}$ is $24.5 \mathrm{~L}$.. $)$
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