An organic compound was found to contain only \(\mathrm{C}, \mathrm{H},\) and \(\mathrm{Cl}\). When a \(1.50-\mathrm{g}\) sample of the compound was completely combusted in air, \(3.52 \mathrm{~g}\) of \(\mathrm{CO}_{2}\) was formed. In a separate experiment, the chlorine in a \(1.00-g\) sample of the compound was converted to \(1.27 \mathrm{~g}\) of AgCl. Determine the empirical formula of the compound.

Short Answer

Expert verified
Based on the given data, the empirical formula of the organic compound containing C, H, and Cl is \(C_{2}H_{3}Cl\).

Step by step solution

01

Calculate moles of Carbon

First, convert the mass of CO2 produced to moles of Carbon using the molecular weight of CO2. Since the molecular weight of CO2 is 44.01 g/mol (12.01 g/mol for C and 32.00 g/mol for O2), and each CO2 molecule contains one atom of carbon, we can find the moles of carbon as follows: \(\text{moles of Carbon} = \frac{3.52\,\text{g CO}2}{44.01\,\text{g/mol}}\)
02

Calculate moles of Chlorine

Next, convert the mass of AgCl produced to moles of Chlorine using the molecular weight of AgCl. Since the molecular weight of AgCl is 143.32 g/mol (107.87 g/mol for Ag and 35.45 g/mol for Cl), and each AgCl molecule contains one atom of Chlorine, we can find the moles of Chlorine as follows: \(\text{moles of Chlorine} = \frac{1.27\,\text{g AgCl}}{143.32\,\text{g/mol}}\)
03

Calculate the mass of Hydrogen

Now, we will calculate the mass of Hydrogen present in the 1.50-g sample by subtracting the masses of Carbon and Chlorine from the sample mass (1.50 g). \(\text{mass of Hydrogen} = \text{mass of sample} - \text{mass of Carbon} - \text{mass of Chlorine}\)
04

Calculate moles of Hydrogen

Next, convert the mass of Hydrogen calculated in Step 3 to moles of Hydrogen using the molecular weight of H (1.01 g/mol). \(\text{moles of Hydrogen} = \frac{\text{mass of Hydrogen}}{1.01\,\text{g/mol}}\)
05

Determine Mole Ratios of C, H, and Cl

Now that we have the moles of each element, we can find the simplest whole number ratio (mole ratio) of the elements by dividing each value by the smallest number of moles among C, H, and Cl.
06

Determine the Empirical Formula

Using the mole ratios obtained in Step 5, we can determine the empirical formula of the organic compound. The compound consists of C, H, and Cl in the simplest whole number ratios found in Step 5. This will give us the empirical formula of the compound.

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Most popular questions from this chapter

A sample of the male sex hormone testosterone, $\mathrm{C}_{19} \mathrm{H}_{28} \mathrm{O}_{2}\(, contains \)3.88 \times 10^{21}$ hydrogen atoms. (a) How many atoms of carbon does it contain? (b) How many molecules of testosterone does it contain? (c) How many moles of testosterone does it contain? (d) What is the mass of this sample in grams?

Valproic acid, used to treat seizures and bipolar disorder, is composed of \(\mathrm{C}, \mathrm{H},\) and \(\mathrm{O} .\) A \(0.165-\mathrm{g}\) sample is combusted to produce \(0.166 \mathrm{~g}\) of water and \(0.403 \mathrm{~g}\) of carbon dioxide. What is the empirical formula for valproic acid? If the molar mass is \(144 \mathrm{~g} / \mathrm{mol}\), what is the molecular formula?

(a) Ibuprofen is a common over-the-counter analgesic with the formula \(\mathrm{C}_{13} \mathrm{H}_{18} \mathrm{O}_{2} .\) How many moles of \(\mathrm{C}_{13} \mathrm{H}_{18} \mathrm{O}_{2}\) are in a 500-mg tablet of ibuprofen? Assume the tablet is composed entirely of ibuprofen. (b) How many molecules of $\mathrm{C}_{13} \mathrm{H}_{18} \mathrm{O}_{2}$ are in this tablet? (c) How many oxygen atoms are in the tablet?

If $2.0 \mathrm{~mol} \mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{CH}_{2} \mathrm{COOH}, 2.0 \mathrm{~mol} \mathrm{C}_{4} \mathrm{H}_{10},\( and \)2.0 \mathrm{~mol}\( \)\mathrm{C}_{6} \mathrm{H}_{6}$ are completely combusted in oxygen, which one produces the largest number of moles of $\mathrm{H}_{2} \mathrm{O}$ ? Which one produces the least? Explain.

Aluminum sulfide reacts with water to form aluminum hydroxide and hydrogen sulfide. (a) Write the balanced chemical equation for this reaction. (b) How many grams of aluminum hydroxide are obtained from \(14.2 \mathrm{~g}\) of aluminum sulfide?

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