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Chapter 3: Problem 63

Consider the combustion of carbon monoxide (CO) in oxygen gas: $$2 \mathrm{CO}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{CO}_{2}(g)$$Starting with 3.60 moles of \(\mathrm{CO}\), calculate the number of moles of \(\mathrm{CO}_{2}\) produced if there is enough oxygen gas to react with all of the CO.

Short Answer

Expert verified
If there is enough oxygen gas to react with all of the carbon monoxide, 3.60 moles of CO2 would be produced from the reaction given.

Step by step solution

01

Understanding the Balanced Chemical Equation

The balanced chemical equation is \(2 CO(g) + O_{2}(g) \rightarrow 2 CO_{2}(g)\). This indicates that two moles of carbon monoxide (CO) reacts with one mole of oxygen (O2) to produce two moles of carbon dioxide (CO2). Indeed, the coefficients in front of the reactants and products all represent the number of moles involved in the chemical reaction.
02

Applying Stoichiometry

Applying the principle of stoichiometry, the stoichiometric ratio of CO to CO2 is 2:2 or 1:1, meaning that one mole of CO will produce one mole of CO2.
03

Calculating Moles of CO2 Produced

Given that the stoichiometric ratio of CO to CO2 is 1:1 and that there are 3.60 moles of CO available, the number of moles of CO2 produced will also be 3.60 moles, assuming there is enough oxygen present to react with all CO.

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Most popular questions from this chapter

Industrially, nitric acid is produced by the Ostwald process represented by the following equations: $$ \begin{aligned} 4 \mathrm{NH}_{3}(g)+5 \mathrm{O}_{2}(g) & \longrightarrow 4 \mathrm{NO}(g)+6 \mathrm{H}_{2} \mathrm{O}(l) \\ 2 \mathrm{NO}(g)+\mathrm{O}_{2}(g) & \longrightarrow 2 \mathrm{NO}_{2}(g) \\\2 \mathrm{NO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(l) & \longrightarrow \mathrm{HNO}_{3}(a q)+\mathrm{HNO}_{2}(a q) \end{aligned}$$What mass of \(\mathrm{NH}_{3}\) (in grams) must be used to produce 1.00 ton of \(\mathrm{HNO}_{3}\) by the above procedure, assuming an 80 percent yield in each step? ( 1 ton = \(2000 \mathrm{lb} ; 1 \mathrm{lb}=453.6 \mathrm{~g} .)\)

An impure sample of zinc (Zn) is treated with an excess of sulfuric acid \(\left(\mathrm{H}_{2} \mathrm{SO}_{4}\right)\) to form zinc sulfate \(\left(\mathrm{ZnSO}_{4}\right)\) and molecular hydrogen \(\left(\mathrm{H}_{2}\right) .\) (a) Write a balanced equation for the reaction. (b) If \(0.0764 \mathrm{~g}\) of \(\mathrm{H}_{2}\) is obtained from \(3.86 \mathrm{~g}\) of the sample, calculate the percent purity of the sample. (c) What assumptions must you make in (b)?

The molar mass of caffeine is \(194.19 \mathrm{~g}\). Is the molecular formula of caffeine \(\mathrm{C}_{4} \mathrm{H}_{5} \mathrm{~N}_{2} \mathrm{O}\) or \(\mathrm{C}_{8} \mathrm{H}_{10} \mathrm{~N}_{4} \mathrm{O}_{2} ?\)

Aspirin or acetyl salicylic acid is synthesized by reacting salicylic acid with aceticanhydride:$$\mathrm{C}_{7}\mathrm{H}_{6}\mathrm{O}_{3}+\mathrm{C}_{4} \mathrm{H}_{6} \mathrm{O}_{3} \quad \longrightarrow \mathrm{C}_{9} \mathrm{H}_{8}\mathrm{O}_{4}+\mathrm{C}_{2} \mathrm{H}_{4} \mathrm{O}_{2}$$ (a) How much salicylic acid is required to produce \(0.400 \mathrm{~g}\) of aspirin (about the content in a tablet), assuming acetic anhydride is present in excess? (b) Calculate the amount of salicylic acid needed if only 74.9 percent of salicylic acid is converted to aspirin. (c) In one experiment, \(9.26 \mathrm{~g}\) of salicylic acid is reacted with \(8.54 \mathrm{~g}\) of acetic anhydride. Calculate the theoretical yield of aspirin and the percent yield if only \(10.9 \mathrm{~g}\) of aspirin is produced.

When \(0.273 \mathrm{~g}\) of \(\mathrm{Mg}\) is heated strongly in a nitrogen \(\left(\mathrm{N}_{2}\right)\) atmosphere, a chemical reaction occurs. The product of the reaction weighs 0.378 g. Calculate the empirical formula of the compound containing \(\mathrm{Mg}\) and \(\mathrm{N}\). Name the compound.
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