Molecular compounds converting empirical formulas

Defining the Mole 72 Determining Molar Mass 73 Converting Between Amount, Mass, and Number of Chemical Entities 74 The Importance of Mass Percent 77 Determining the Formula of an Unknown Compound 80 Empirical Formulas 80 Molecular Formulas 81 Isomers 84... 

A 1.50 2.99 1.00 ratio converts to 3 6 2 by multiplying by 2. Therefore, the simplest formula for this compound is C3H602 (FW = 74.1 g/mol). The true molecular formula for the compound is the same, C3H602, because the true molecular formula is the same as the empirical formula. 

In the problem above, we determined the percentage data from the chemical formula. We can determine the empirical formula if we know the percent compositions of the various elements. The empirical formula tells us what elements are present in the compound and the simplest whole-number ratio of elements. The data may be in terms of percentage, or mass or even moles. However, the procedure is still the same—convert each element to moles, divide each by the smallest, and then use an appropriate multiplier if necessary. We can then determine the empirical formula mass. If we know the actual molecular mass, dividing the molecular formula mass by the empirical formula mass, gives an integer (rounded if needed) that we can multiply each of the subscripts in the empirical formula. This gives the molecular (actual) formula, which tells what elements are in the compound and the actual number of each. 

The percent composition of glucose can be calculated either from the molecular formula (CgH Og) or from the empirical formula (CH20). Using the molecular formula, for instance, the C H 0 mole ratio of 6 12 6 can be converted into a mass ratio by assuming that we have 1 mol of compound and carrying out mole-to-gram conversions ... 

The empirical formula is the formula for a compound that is expressed in the lowest ratio that can be calculated (refer to Chapter 2). Often, a substance must be analyzed to gather information leading to its identity. Various processes can be used to determine the composition of a sample, and an effective way of expressing these data is in the form of weight. Weights can be converted to moles and expressing a formula is the next logical step. The empirical formula is not necessarily the actual molecular formula however, the empirical formula does contain important information. 

Calculation of the Empirical Formula Molecular formulas can be determined by a two-step process. The first step is the determination of an empirical formula, simply the relative ratios of the elements present. Suppose, for example, that an unknown compound was found by quantitative elemental analysis to contain 40.0% carbon and 6.67% hydrogen. The remainder of the weight (53.3%) is assumed to be oxygen. To convert these numbers to an empirical formula, we can follow a simple procedure. 

Whenever we know the molecular formula of a compound, we can determine its empirical formula. The converse is not true, however. If we know the empirical formula of a substance, we cannot determine its molecular formula unless we have more information. So why do chemists bother with empirical formulas As we will see in Chapter 3, certain common methods of analyzing substances lead to the empirical formula only. Once the empirical formula is known, additional experiments can give the information needed to convert the empirical formula to the molecular one. In addition, there are substances that do not exist as isolated molecules. For these substances, we must rely on empirical formulas. 

Calculation of the Molecular Formula How do we know the correct molecular formula We can choose the right multiple of the empirical formula if we know the molecular weight. Molecular weights can be determined by methods that relate the freezing-point depression or boiling-point elevation of a solvent to the molal concentration of the unknown. If the compound is volatile, we can convert it to a gas and use its volume to determine the number of moles according to the gas law. Newer methods include mass spectrometry, which we will cover in Chapter 11. 

Making particle numbers manageable with Avogadro s number Converting between masses, mole counts, and volumes Dissecting compounds with percent composition Moving from percent composition to empirical and molecular formulas... 

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