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Chapter 4: Problem 15

Iron metal reacts with chlorine gas. How many grams of \(\mathrm{FeCl}_{3}\) are obtained when \(515 \mathrm{g} \mathrm{Cl}_{2}\) reacts with excess Fe? $$ 2 \mathrm{Fe}(\mathrm{s})+3 \mathrm{Cl}_{2}(\mathrm{g}) \longrightarrow 2 \mathrm{FeCl}_{3}(\mathrm{s}) $$

Short Answer

Expert verified
783.6 grams of FeCl3 can be produced.

Step by step solution

01

Determine the molar mass of Cl2 and FeCl3

From the periodic table, the molar mass of Cl is approximately 35.5 g/mol. Therefore the molar mass of Cl2 is \(35.5 g/mol * 2 = 71 g/mol\). The molar mass of FeCl3 is \(55.85 g/mol + 35.5 g/mol * 3 = 162.2 g/mol\). This is the number of grams per mole for each substance.
02

Convert mass of Cl2 to moles

The mass of Cl2 given is 515g. To convert this mass to moles, we use the molar mass from step 1. Therefore, the number of moles of Cl2 is \(515g ÷ 71g/mol = 7.25 mol\). This is how many moles of Cl2 we have.
03

Use stoichiometry to find moles of FeCl3

By looking at the balanced chemical equation, we can see that for every 3 moles of Cl2, 2 moles of FeCl3 are produced. Therefore, we can create a stoichiometric ratio and multiply it by the moles of Cl2 to find the moles of FeCl3. The number of moles of FeCl3 is \(7.25 mol * 2/3 = 4.83 mol\). This is how many moles of FeCl3 we can produce.
04

Convert moles of FeCl3 to grams

To find the mass of FeCl3 produced, we use the molar mass of FeCl3 from step 1 and multiply it by the moles of FeCl3. The mass of FeCl3 is \(4.83 mol * 162.2 g/mol = 783.6 g\). This is how many grams of FeCl3 we can produce.

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Most popular questions from this chapter

High-purity silicon is obtained using a three-step process. The first step involves heating solid silicon dioxide, \(\mathrm{SiO}_{2^{\prime}}\) with solid carbon to give solid silicon and carbon monoxide gas. In the second step, solid silicon is converted into liquid silicon tetrachloride, \(\mathrm{SiCl}_{4}\) by treating it with chlorine gas. In the last step, \(\mathrm{SiCl}_{4}\) is treated with hydrogen gas to give ultrapure solid silicon and hydrogen chloride gas. (a) Write balanced chemical equations for the steps involved in this three- step process. (b) Calculate the masses of carbon, chlorine, and hydrogen required per kilogram of silicon.

Baking soda, \(\mathrm{NaHCO}_{3}\), is made from soda ash, a common name for sodium carbonate. The soda ash is obtained in two ways. It can be manufactured in a process in which carbon dioxide, ammonia, sodium chloride, and water are the starting materials. Alternatively, it is mined as a mineral called trona (left photo). Whether the soda ash is mined or manufactured, it is dissolved in water and carbon dioxide is bubbled through the solution. Sodium bicarbonate precipitates from the solution. As a chemical analyst you are presented with two samples of sodium bicarbonate-one from the manufacturing process and the other derived from trona. You are asked to determine which is purer and are told that the impurity is sodium carbonate. You decide to treat the samples with just sufficient hydrochloric acid to convert all the sodium carbonate and bicarbonate to sodium chloride, carbon dioxide, and water. You then precipitate silver chloride in the reaction of sodium chloride with silver nitrate. A \(6.93 \mathrm{g}\) sample of baking soda derived from trona gave \(11.89 \mathrm{g}\) of silver chloride. A \(6.78 \mathrm{g}\) sample from manufactured sodium carbonate gave \(11.77 \mathrm{g}\) of silver chloride. Which sample is purer, that is, which has the greater mass percent \(\mathrm{NaHCO}_{3} ?\)

Chromium(II) sulfate, \(\mathrm{CrSO}_{4^{\prime}}\) is a reagent that has been used in certain applications to help reduce carbon-carbon double bonds \((\mathrm{C}=\mathrm{C})\) in molecules to single bonds ( \(\mathrm{C}-\mathrm{C}\) ). The reagent can be prepared via the following reaction. $$\begin{array}{c} 4 \mathrm{Zn}(\mathrm{s})+\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}(\mathrm{aq})+7 \mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{aq}) \longrightarrow 4 \mathrm{ZnSO}_{4}(\mathrm{aq})+2 \mathrm{CrSO}_{4}(\mathrm{aq})+\mathrm{K}_{2} \mathrm{SO}_{4}(\mathrm{aq})+7 \mathrm{H}_{2} \mathrm{O}(\mathrm{l}) \end{array}$$

Under appropriate conditions, copper sulfate, potassium chromate, and water react to form a product containing \(\mathrm{Cu}^{2+},\) \(\mathrm{CrO}_{4}{^2}{^-},\) and \(\mathrm{OH}^{-}\) ions. Analysis of the compound yields \(48.7 \% \mathrm{Cu}^{2+}, 35.6 \% \mathrm{CrO}_{4}{^2}{-},\) and \(15.7 \% \mathrm{OH}^{-}\). (a) Determine the empirical formula of the compound. (b) Write a plausible equation for the reaction.

Iron ore is impure \(\mathrm{Fe}_{2} \mathrm{O}_{3} .\) When \(\mathrm{Fe}_{2} \mathrm{O}_{3}\) is heated with an excess of carbon (coke), metallic iron and carbon monoxide gas are produced. From a sample of ore weighing \(938 \mathrm{kg}, 523 \mathrm{kg}\) of pure iron is obtained. What is the mass percent \(\mathrm{Fe}_{2} \mathrm{O}_{3},\) by mass, in the ore sample, assuming that none of the impurities contain Fe?
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