Compounds mole concept

This is a critical chapter in your study of chemistry. Our goal is to help you master the mole concept. You will learn about balancing equations and the mole/mass relationships (stoichiometry) inherent in these balanced equations. You will learn, given amounts of reactants, how to determine which one limits the amount of product formed. You will also learn how to determine the empirical and molecular formulas of compounds. All of these will depend on the mole concept. Make sure that you can use your calculator correctly. If you are unsure about setting up problems, refer back to Chapter 1 of this book and go through Section 1-4, on using the Unit Conversion Method. Review how to find atomic masses on the periodic table. Practice, Practice, Practice. 

Using the mole concept and the periodic table, you can determine the mass of one mole of a compound. You know, however, that one mole represents 6.02 x 1023 particles. Therefore you can use a balance to count atoms, molecules, or formula units ... 

In this chapter, you have learned about the relationships among the number of particles in a substance, the amount of a substance in moles, and the mass of a substance. Given the mass of any substance, you can now determine how many moles and particles make it up. In the next chapter, you will explore the mole concept further. You will learn how the mass proportions of elements in compounds relate to their formulas... 

The empirical formula of a compound can be simply related to the mass percentage of its constituent elements using the mole concept. For example, the empirical formula for ethylene (molecular formula C2H4) is CH2. Its composition by mass is calculated from the masses of carbon and hydrogen in 1 mol of CH2 formula units ... 

We apply the mole concept to determine chemical formulas from the masses of each element in a given quantity of a compound. 

As we learned in Section 2.6, the empirical formula for a substance tells us the relative number of atoms of each element in the substance. The empirical formula H2O shows that water contains two H atoms for each O atom. This ratio also applies on the molar level 1 mol of H2O contains 2 mol of H atoms and 1 mol of O atoms. Conversely, the ratio of the numbers of moles of all elements in a compound gives the subscripts in the compound s empirical formula. Thus, the mole concept provides a way of calculating empirical formulas. 

You can also use the mole concept to calculate the empirical formula of a compound using the percentage composition data for that compound — the percentage by weight of each element in the compound. (The empirical formula indicates the different types of elements in a molecule and the lowest whole-number ratio of each kind of atom in the molecule. See Chapter 7 for details.)... 

Use the mole concept and molecular formulas to obtain relationships between number of moles, number of grams, and number of atoms or molecules for compounds, and use those relationships to obtain factors for use in factor-unit calculations. (Section 2.7)... 

In a strict sense, 1 mol is a specific number of particles. However, in chemistry it is customary to follow the useful practice of also letting 1 mol stand for the mass of a sample of element or compound that contains Avogadro s number of particles. Thus, the application of the mole concept to sulfur (at. wt. = 32.1 u) gives the following relationships ... 

The mole concept can also be applied to particles that are molecules instead of atoms. The compound carbon dioxide consists of molecules that contain one carbon atom, C, and two oxygen atoms, O. The formula for the molecule is CO2. The molecular weight of the molecule is calculated as shown earlier by adding together the atomic weight of one carbon atom and the atomic weight of two oxygen atoms ... 

According to Section 2.1 and as demonstrated in Example 2.8, the formula for a compound is made up of the symbols for each element present. Subscripts following the elemental symbols indicate the number of each type of atom in the molecule represented. Thus, chemical formulas represent the numerical relationships that exist among the atoms in a compound. AppUcation of the mole concept to the atoms making up the formulas provides additional useful information. 

The mole is often referred to as a chemist s unit of quantity. Counting atoms is a difficult process and beyond the scope of most calculators, but measuring the mass of a sample is easy when we can relate the number of atoms in a sample to its mass. This is the unique purpose of the mole. A mole of any substance is its molecular formula weight expressed in grams. Avogadro s number s a universal constant that states the number of molecules in a mole Nq = 6.023 x 10 molecules/mole. One mole (abbreviated mol) of any element (chemical compound) has the same number of chemical particles as one mole of another element (chemical compound). In other words, 1 mole of any compound contains 6.02 x 10 molecules. Review the following problem using the mole concept. 

EXAMPLE 6.4 The Mole Concept—Converting between Grams and Moles for Compounds... 

EXAMPLE 6.5 The Mole Concept—Converting between Mass of a Compound and Number of Molecules... 

The Mole Concept The mole is a specific number (6.022 X 10 ) that allows us to easily coimt atoms or molecules by weighing them One mole of any element has a mass equivalent to its atomic mass in grams, and a mole of any compoimd has a mass equivalent to its formula mass in grams. The mass of 1 mol of an element or compound is its molar mass. 

When we prepare a compound industrially or even study a reaction in the laboratory, we deal with tremendous numbers of molecules or ions. Suppose you wish to prepare acetic acid, starting from 10.0 g of ethanol. This small sample (less than 3 tea-spoonsful) contains 1.31 X 10 molecules, a truly staggering number. Imagine a device that counts molecules at the rate of one million per second. It wonld take more than four billion years—nearly the age of the earth—for this device to count that many molecules Chemists have adopted the mole concept as a convenient way to deal with the enormons numbers of molecules or ions in the samples they work with. 

The mole concept can be used to calculate the formula of a substance from experimental results. The formula obtained is the simplest possible formula (involving integers) for that compound. 

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