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Chapter 3: Problem 38

If we know the empirical formula of a compound, what additional information do we need to determine its molecular formula?

Short Answer

Expert verified
To find the molecular formula from the empirical formula, additional information needed is the molar mass of the compound.

Step by step solution

01

Understanding Empirical Formula

The empirical formula of a compound represents the simplest ratio of the elements present in it. For example, the empirical formula of Benzene is CH, which means the elements Carbon (C) and Hydrogen (H) are present in the simplest ratio 1:1.
02

Need for Molecular Formula

The molecular formula of a compound represents the actual number of atoms of each element present in one molecule of the compound. In the example of Benzene, the molecular formula is C6H6 which represents that in one molecule of Benzene, there are 6 Carbon atoms and 6 Hydrogen atoms.
03

Identifying the Additional Information Needed

To determine the molecular formula from the empirical formula, it's necessary to know the molar mass of the compound. If we divide the molar mass of the compound by the molar mass implied by the empirical formula, we'll get the factor to multiply the empirical formula to determine the molecular formula.

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Most popular questions from this chapter

What are the empirical formulas of the compounds with the following compositions? (a) 2.1 percent \(\mathrm{H}\), 65.3 percent \(\mathrm{O}, 32.6\) percent \(\mathrm{S}\) (b) 20.2 percent \(\mathrm{Al}\), 79.8 percent \(\mathrm{Cl}\)

Describe the steps involved in the mole method.

Fermentation is a complex chemical process of wine making in which glucose is converted into ethanol and carbon dioxide: $$ \mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6} \longrightarrow 2 \mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}+2 \mathrm{CO}_{2} $$ glucose ethanol Starting with \(500.4 \mathrm{~g}\) of glucose, what is the maximum amount of ethanol in grams and in liters that can be obtained by this process? (Density of ethanol = \(0.789 \mathrm{~g} / \mathrm{mL} .)\)

How many atoms are there in 5.10 moles of sulfur (S)?

An iron bar weighed 664 g. After the bar had been standing in moist air for a month, exactly one-eighth of the iron turned to rust \(\left(\mathrm{Fe}_{2} \mathrm{O}_{3}\right)\). Calculate the final mass of the iron bar and rust.
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