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Chapter 9: Problem 176

Consider the following decomposition reaction in which \(47.20 \mathrm{~g}\) of some compound is decomposed into its elements.

Short Answer

Expert verified
A general outline for solving decomposition reactions involves: understanding the decomposition reaction, writing the balanced chemical equation, finding the molar mass of the compound and its constituent elements, converting grams of the compound to moles, using stoichiometry to determine the moles of each element produced, converting moles of elements to grams, and verifying the conservation of mass. In this specific case, we don't have enough information to provide a complete solution, but knowing these steps will help you solve similar problems in the future.

Step by step solution

01

Understand the decomposition reaction

In a decomposition reaction, a single compound breaks down into two or more simpler substances (usually the constituent elements). These reactions typically require an input of energy in the form of heat, light, or electricity, e.g., 2H₂O → 2H₂ + O₂.
02

Write the balanced chemical equation

Write the balanced chemical equation for the decomposition reaction. Make sure the number of atoms on both sides of the equation is equal. For example, if the compound AB decomposes into A and B, make sure both sides of the equation have an equal number of A and B atoms: AB → A + B.
03

Find the molar mass of the compound and its constituent elements

Use a periodic table to find the molar mass of each element in the compound and the compound itself. Molar mass is the mass of one mole of the substance in grams, and it is equal to the sum of the atomic masses of the constituent elements.
04

Convert grams of the compound to moles

Divide the given mass of the compound (47.20 grams) by its molar mass to find the moles of the compound. Moles = Mass / Molar Mass.
05

Use stoichiometry to determine the moles of each element produced

Refer to the balanced chemical equation and use stoichiometry to determine the number of moles of each constituent element that would be produced when the compound decomposes. This will involve determining the mole ratio between the compound and each element in the reaction.
06

Convert moles of elements to grams

Multiply the moles of each element by its molar mass to determine the mass of each element produced in grams. Mass = Moles × Molar Mass.
07

Verify conservation of mass

The law of conservation of mass states that mass cannot be created nor destroyed in a chemical reaction. Ensure the sum of the masses of the constituent elements produced is equal to the initial mass of the compound. In this case, the sum of the masses of the elements should be equal to 47.20 grams. Remember that this is a general outline of the steps to solve a decomposition reaction problem. In this particular case, we do not have the specific compound or balanced chemical equation needed for a complete solution but understanding these steps prepares you to solve similar problems in the future.

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Most popular questions from this chapter

When \(490.0 \mathrm{mg}\) of iron reacts with excess bromine, a mixture of \(\mathrm{FeBr}_{2}\) and \(\mathrm{FeBr}_{3}\) is produced. It is determined that \(35.5 \%\) of the mixture is iron(II) (bromide). What is the total mass of \(\mathrm{FeBr}_{2} / \mathrm{FeBr}_{3}\) mixture produced?

A \(1.000-\mathrm{g}\) sample of a liquid is subjected to combustion analysis, yielding \(92.3 \% \mathrm{C}\) and \(7.7 \% \mathrm{H}\). It may or may not also contain oxygen. (a) What is the empirical formula for this compound? (b) The molar mass of this compound is determined to be about \(78 \mathrm{~g} / \mathrm{mol}\). What is the molecular formula for this compound?

For the reaction \(\mathrm{Ca}_{3} \mathrm{P}_{2}+\mathrm{H}_{2} \mathrm{O} \rightarrow \mathrm{Ca}(\mathrm{OH})_{2}+\mathrm{PH}_{3}\), (a) Balance the equation. (b) How many grams of \(\mathrm{H}_{2} \mathrm{O}\) would you need to react with \(60.0 \mathrm{~g}\) of \(\mathrm{Ca}_{3} \mathrm{P}_{2} ?\) (c) How many grams of \(\mathrm{PH}_{3}\) could theoretically be produced from the reactant amounts calculated in part (b)?

(a) Write a balanced chemical equation for the combustion of \(\mathrm{CH}_{4}(\mathrm{~g})\) with \(\mathrm{O}_{2}(g)\) to form \(\mathrm{CO}_{2}(g)\) and \(\mathrm{H}_{2} \mathrm{O}(g)\) (b) Consider the reaction of three molecules of \(\mathrm{CH}_{4}(g)\) with four molecules of \(\mathrm{O}_{2}(g)\). The following boxes represent before and after pictures of the tiny, sealed container in which this reaction takes place. Draw pictures to show the contents of this container before and after the reaction happens:

Consider the unbalanced chemical equation \(\mathrm{As}_{4} \mathrm{~S}_{6}+\mathrm{O}_{2} \rightarrow \mathrm{As}_{4} \mathrm{O}_{6}+\mathrm{SO}_{2}\) (a) How many grams of \(\mathrm{O}_{2}\) are needed to react completely with \(58.9 \mathrm{~g}\) of \(\mathrm{As}_{4} \mathrm{~S}_{6}\) ? (b) If \(41.2 \mathrm{~g}\) of \(\mathrm{SO}_{2}\) is produced, what is the percent yield for the reaction?
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