Calculations Using Moles and Molar Masses

Looking at these three compounds, we see that nitroglycerin has the smallest molar mass even though it clearly contains the most atoms per molecule. Does this make sense The key factor is that both lead azide and mercury fulminate contain elements with large molar masses, so it is not surprising that nitroglycerin is the least massive of these three molecules. 

The number of significant figures to which the molar masses of the elements are known varies from one element to another, as you can see if you examine the periodic table inside the front cover of this book. So the appropriate number of significant figures in the molar masses here varies according to the elements in each molecule. 

What are the molar masses of the following compounds, which are used in the preparation of various explosives (a) H2SO4, (b) HNO3, (c) (NH2)2HN03 

We ve seen the importance of the concept of the mole when dealing with chemical reactions involving macroscopic quantities of material. But generally it is not possible to measure the number of moles in a sample directly because that would imply that we can count molecules. Molar masses provide a crucial connection and allow us to convert from masses, which can be measured easily, to numbers of moles. The mass and the number of moles are really just two different ways of expressing the same information—the amount of a substance present. The molar mass functions much like a unit conversion between them, allowing us to go from 

A sample of the explosive TNT (C7H5N3O6) has a mass of 650.5 g. How many moles of TNT are in this sample How many molecules is this  

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Calculations Using Moles and Molar Masses Elemental Analysis ... 

In Chapter 13, you learned how to use moles and molar mass along with a balanced chemical equation to calculate the masses of reactants and products in a chemical reaction. Now that you know how to relate volumes, masses, and moles for a gas, you can do stoichiometric calculations for reactions involving gases. 

In calculations using equations, the molar masses of the substances and their mole-mole factors are used to change the number of grams of one substance to the corresponding grams of a different substance. 

Chapter 7, Chemical Quantities and Reactions, introduces moles and molar masses of compounds, which are used in calculations to determine the mass or number of particles in a given quantity. Students leam to balance chemical equations and to recognize the types of chemical reactions combination, decomposition, single replacement, double replacement, and combustion reactions. Section 7.5 discusses Oxidation-Reduction Reactions using real-life examples, including biological reactions. Section... 

Changing grams to moles and moles to grams is perhaps the most important calculation you will have to make all year (Fig. 4-2), Some authors use the term molar mass for the mass of 1 mol of any substance. The units are typically grams per mole. 

A variation of this experiment uses the mass of the anhydrous material (calculated from the difference between masses 1 and 3). The moles of the anhydrous material and water are then calculated from their respective masses and molar masses. The simplest ratio of the moles gives the empirical formula. 

STRATEGY Once we know the mole fraction of the solvent (water) in the solution, the calculation is a straightforward application of Raoult s law. To calculate the mole fractions, convert masses to moles by using the molar masses of sucrose and water, then divide each value by the total number of moles. To use Raoult s law, we need the vapor pressure of the pure solvent, Table 8.2 or 8.3. Expect a lower vapor pressure when the solute is present. 

The problem gives the number of moles of NaHC03 and asks for a mole-to-mass conversion. First, calculate the formula mass and molar mass of NaHC03. Then use molar mass as a conversion factor, and set up an equation so that the unwanted unit cancels. 

Scientists also use the simplest formula to represent one mole of an ionic compound. They often use the term formula unit when referring to ionic compounds, because they are not found as single molecules. A formula unit of an ionic compound represents the simplest ratio of cations to anions. A formula unit of KBr is made up of one ion and one Br ion. One mole of an ionic compound has 6.022 x 10 of these formula units. As with molecular compounds, the molar mass of an ionic compound is the sum of the masses of all the atoms in the formula expressed in g/mol. Table 1 compares the formula units and molar masses of three ionic compounds. Sample Problem F shows how to calculate the molar mass of barium nitrate. 

Using atomic mass and molar mass, we can calculate the mass in grams of a single carbon-12 atom. From our discussion we know that 1 mole of carbon-12 atoms weighs exactly 12 grams. This allows us to write the equality... 

Plan We find the masses of CO2 and H2O by subtracting the masses of the absorbers before the reaction from the masses after. From the mass of CO2, we use the mass fraction of C in CO2 to find the mass of C (see Comment in Sample Problem 3.3). Similarly, we find the mass of H from the mass of H2O. The mass of vitamin C (I.OOO g) minus the sum of the C and H masses gives the mass of O, the third element present. Then, we proceed as in Sample Problem 3.5 calculate numbers of moles using the elements molar masses, construct the empirical formula, determine the whole-number multiple from the given molar mass, and construct the molecular formula. 

Realize the usefulness of the mole concept, and use the relation between molecular (or formula) mass and molar mass to calculate the molar mass of any substance ( 3.1) (EPs 3.1-3.5, 3.7-3.10)... 

One very important use of the ideal gas law is in the calculation of the molar mass (molecular weight) of a gas from its measured density. To see the relationship between gas density and molar mass, consider that the number of moles of gas n can be expressed as... 

The equation relating the amount of substance in moles, mass and molar mass can be used to calculate any quantity given the values of the other two ... 

Throughout the main text of this book standard solutions and quantities have all been expressed in terms of molarities, moles and relative molecular masses. However, there are still many chemists who have traditionally used what are known as normal solutions and equivalents as the basis for calculations, especially in titrimetry. Because of this it has been considered desirable to include this appendix defining the terms used and illustrating how they are employed in the various types of determinations. 

STRATEGY We convert from the given volume of gas into moles of molecules (by using the molar volume), then into moles of reactant molecules or formula units (by using a mole ratio), and then into the mass of reactant (by using its molar mass). If the molar volume at the stated conditions is not available, then use the ideal gas law to calculate the amount of gas molecules. 

To determine the mass of NiCl2 6 H2 O required to prepare the solution, first calculate the number of moles of the salt required, and then use the molar mass to determine the mass in grams ... 

The chemicai formuia identifies the ions that are present in the finai soiution. The formula also tells us how many moles of each ion are present in one moie of the sait. Use mass, molar mass, and volume to calculate molarity. 

The question asks for the mass of oxygen. We can use the ideal gas equation to calculate the number of moles of oxygen, and then molar mass leads us from moles to grams. 

The ideal gas equation can be combined with the mole-mass relation to find the molar mass of an unknown gas PV = nRT (ideal gas equation) and n — (mole-mass relation) if we know the pressure, volume, and temperature of a gas sample, we can use this information to calculate how many moles are... 

Measuring and Using Numbers Use the moles of acetic acid and the molar mass of acetic acid to calculate the mass of acetic acid in the vinegar sample. 

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