The Mole Concept

2 The mole concept - the mole makes it possible to correlate the 

Chemists are interested in the ratios in which chemical elements and compounds react together during chemical reactions. This is important when preparing a pure substance in the laboratory, and even more so in the chemical industry. Using excess reactant, unless necessary, will result in additional costs in order to remove it from the product. 

Atoms are very small with very small masses, for example a hydrogen atom ( H) weighs only 1.67355 X 10 kg. However, the masses of atoms of different elements are different, for example a carbon-12 atom is twelve times more massive than an atom of hydrogen-L 

For this reason, weighing out the same mass of different elements results in different numbers of atoms being present in the samples. It is very difficult for chemists to count large numbers of atoms directly so instead a chemist counts atoms indirectly by weighing samples of elements. 

The amount of substance (symbol n) is a quantity that is directly proportional to the number of particles in a sample of substance. It is one of the seven base quantities of the SI unit system. The unit of amount is the mole (mol). 

Note that the atomic mass of chlorine, Cl, is 35.45 amu, but the molecular mass of chlorine, CI2, is two times 35.45 amu. This illustrates an important point Always calculate formula or molecular masses exactly as the formula is written. The formula, of course, must be written correctly  

GOAL 4 Define the term mole. Identify the number of objects that corresponds to one mole. 

5 Given the number of moles (or units) in any sample, calculate the number of units (or moles) in the sample. 

In Section 2.1 we said that you would study chemistry at the particulate level. In other words, chemists are interested in the individual particles that make up a sample of matter. In Section 2.6 we identified atoms and molecules as two of these particles. To understand the amounts of substances in a chemical change, we must know the number of particles of the different substances in the reaction. That s not an easy number to find. Literally counting atoms and molecules is not practical. The number is extremely large. 

To describe the number of particles, chemists use a quantity called the mole One mole is the amount of any substance that contains the same number of units as the number of atoms in exactly 12 grams of carbon-12. In calculation setups the mole is abbreviated mol. 

For most of the questions we will be called upon to answer, working with the actual numbers of molecules is awkward because there are so many of them in even the most minute portions of matter. This awkwardness can be removed, however, by using 6.023 x 10 molecules Avogadro s number, the number of C atoms in exactly 12g of C), the chemist s dozen, as a basic unit. The amount of matter containing an Avogadro number of molecules of any substance is called a mole. This amount will have a mass in grams equal to its gram molecular mass (GMW). To calculate the 

For instance, if 10 g H3PO4 is to be neutralized with NaOH, what weight of NaOH is required We see from Equation 1 -2 that 10/ GMW moles 

This reaction also produces 10 (GMWjj pQ /GMWjj po ) g of Na3P04 and 10(3 GMWh o GMW po ) g of H2O.  

With this approach, we see that we can write expressions for the amounts of substances related by participation in the same (set of) reaction(s). Let us consider the question of how much BaS04, xg, can be produced from 1.00 g of NaCl by the following reactions  

Launching out bravely on our discussion of the gravimetric factor, we can anticipate that the answer will have the form  

By the way, the major contribution of Avogadro to chemistry was his hypothesis that atmospheric gases are made of molecules, which is the name he used for the 

FIGURE 4.R.1 Amedeo Avogadro s caricature (drawn by the historian W. B. Jensen and reproduced here with his permission) and Avogadro s number. 

Out of the atomic theory developed by John Dalton and other chemistry pioneers in the nineteenth century grew a number of important concepts essential to an understanding of all areas of chemistry, including pyrotechnics and explosives. The basic features of the atomic theory are  

The atom is the fundamental building block of matter, and consists of positive, negative, and neutral subatomic particles. Approximately ninety naturally occurring elements are known to exist (additional elements have been synthesized in the twentieth century in the laboratory using high-energy nuclear reactions, but these unstable species are not found in nature). 

Elements may combine to form more complex species called compounds. The molecule is the fundamental unit of a compound and consists of two or more atoms joined together by chemical bonds. 

All atoms of the same element are identical in terms of the number of protons and electrons contained in the neutral species. Atoms of the same element may vary in the number of neutrons, and therefore may vary in mass. 

The chemical reactivity of an atom depends on the number of electrons therefore, the reactivity of all atoms of a given elanent should be the same, and reproducible, anywhere in the world. 

---

If these assumptions are valid, however, stoichiometric calculations provide a reliable basis for quantitative predictions. It is important to be able to make these calculations with ease. Fortunately, they all can be made with a single pattern based upon the mole concept. 

The initial set of experiments and the first few textbook chapters lay down a foundation for the course. The elements of scientific activity are immediately displayed, including the role of uncertainty. The atomic theory, the nature of matter in its various phases, and the mole concept are developed. Then an extended section of the course is devoted to the extraction of important chemical principles from relevant laboratory experience. The principles considered include energy, rate and equilibrium characteristics of chemical reactions, chemical periodicity, and chemical bonding in gases, liquids, and solids. The course concludes with several chapters of descriptive chemistry in which the applicability and worth of the chemical principles developed earlier are seen again and again. 

Sub-micro representations are used extensively in teaching the mole concept, stoichiometiy, solubihty and chemical equilibrium at UCT. Having students draw and annotate chemical diagrams representing chemical phenomena at the sub-micro level can provide some insight into their understanding of chemistiy at the macro level. The following examples are typical of the questions used to probe links between the sub-micro and symbohc levels of representations as part of the assessment practice for this course. For example, students were asked to balance the equation shown in Fig. 8.7. 

We developed the concept of the mole In terms of pure chemical substances, but many chemical reactions take place In solution. To treat solution reactions quantitatively, we need ways to apply the mole concept to solutions. A substance used to dissolve solutes Is a solvent, and a pure substance dissolved In solution Is a solute. Most of the time, the solvent Is a liquid and Is present In much larger quantities than any solutes. 

How can the mole concept be used to predict the limiting reactant in a chemical reaction ... 

Mass and volume relations with emphasis on the mole concept, including empirical formulas and limiting reactants... 

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. 

It is possible to expand these examples to any titration problem, acid-base, redox, precipitation, and so on. Just remember that the key is the mole concept. 

Our goal in this chapter is to assist you in learning the concepts of gases and gas laws. Be sure that you know how to properly use your calculator and, if you need to, refer to Chapter 3 on the mole concept. It s especially true with gas law problems that the only way to master them is to Practice, Practice, Practice. 

In this chapter, we will help you learn about the energy changes, especially heat, which occurs during both physical and chemical changes. You might need to review the Unit Conversion Method in Chapter 1 and the sections in Chapter 3 on balancing chemical reactions and the mole concept if you are not comfortable with them already. And remember to Practice, Practice, Practice. 

The major goal of this chapter is to help you master the concepts associated with solutions—concentration units, solubility, and especially colligative properties. We will also examine the properties of colloids. If you are still unsure about calculations and the mole concept, review Chapters 1,3, and 4. And again, the only way to master these concepts is to Practice, Practice, Practice. 

Our goal in this chapter is to help you understand how to balance redox equations, know the different types of electrochemical cells, and how to solve electrolysis problems. Have your textbook handy—you may need to find some information in electrochemical tables. We will be using the mole concept, so if you need some review refer to Chapter 3, especially the mass/mole relationships. You might also need to review the section concerning net-ionic equations in Chapter 4. And don t forget to Practice, Practice, Practice. 

A production chemist is interested primarily in the macroscopic world, not the microscopic one of atoms and molecules. Even a chemistry student working in the laboratory will not be weighing out individual atoms and molecules, but large numbers of them in grams. There must be a way to bridge the gap between the microscopic world of individual atoms and molecules, and the macroscopic world of grams and kilograms. There is—it is called the mole concept, and it is one of the central concepts in the world of chemistry. 

The actual number of atoms in one mole of an element has been determined by several elegant experimental procedures to be 6.02 X 10 This quantity is known as Avogadro s number, in honor of one of the pioneers of the atomic theory. One can then see that one mole of carbon atoms (12.01 grams) will contain exactly the same number of atoms as one mole (55.85 grams) of iron. Using the mole concept, the chemist can now go into the laboratory and weigh out equal quantities of atoms of the various elements. 

Students will explain the mole concept and use this concept to prepare chemical solutions of particular molarities. 

During this procedure, your teacher will introduce the mole concept. Use a periodic table to find the relative masses of all the elements in the molecule CuS04 and K2Cr04, respectively Cu (copper), S (sulfur) and O (oxygen) and K (potassium), Cr (chromium), O (oxygen). The relative mass in grams for any element contains the same number of atoms. This number of atoms, 6.02 x 1023, is called a mole. In the preparation of any 0.1 M solution, 0.1 mole of molecules is needed. A 0.1 M solution, by definition, contains 0.1 mole of a substance dissolved in 1.0 liter of a solvent. 

Several topics are suggested here that could be de-emphasized or eliminated, affording instructors time to explore biochemical topics more fully electron configuration, quantum numbers, atomic orbitals, the mole concept, limiting reactant and stoichiometry, organic nomenclature, and organic reactions by functional group. 

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... 

Assume that your friend has missed several chemistry classes and that she has asked you to help her prepare for a stoichiometry test. Unfortunately, because of other commitments, you do not have time to meet face to face. You agree to email your friend a set of point-form instructions on how to solve stoichiometry problems, including those that involve a limiting reactant. She also needs to understand the concept of percentage yield. Write the text of this email. Assume that your friend has a good understanding of the mole concept. 

On this day, schools have different activities related to the mole concept or chemistry. 

In chemistry, what is meant by the term mole What is the importance of the mole concept ... 

You can read more about the use of the mole concept in Unit 4.3 Taking the earlier example of burning methane gas in oxygen ... 

The students ultimately find that the second reaction scheme utilizes fewer moles of sodium hydroxide and sulfuric acid. Exercises such as this not only help teach about green chemistry principles, but also they teach about fundamental chemical concepts such as balancing chemical equations, stoichiometry, and the mole concept. 

Concept Mapping Design a concept map that illustrates the mole concept. Include moles, Avogadro s number, molar mass, number of particles, percent composition, empirical formula, and molecular formula. 

The mole concept that connects weighing and counting molecules and atoms... 

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 ... 

---