Working through this chapter of the study guide will enable you to:

1. Distinguish between the physical and the chemical properties of substances, and tell how chemical reactions can change these properties.
2. Appreciate the size and meaning of Avogadro's number.
3. Balance simple chemical equations.
4. Understand the role of energy in chemical reactions and explain how various factors can affect reaction rate.
5. Differentiate between acids and bases and tell the properties of each.
6. Explain the differences between oxidation and reduction.
7. Describe basic types of chemical reactions such as combination, decomposition, single-replacement, and double-replacement.

Discussion

Chemical bonds can be made and broken in many different ways. This chapter looks into these processes using the concept of chemical reactions. The equations that can be written to describe chemical reactions tell us a great deal about the reactants involved in these processes and about the products produced when they occur. Combined with an understanding of energy, reaction rate, and the ideas of equilibrium and reversible chemical processes, many of these seemingly complex reactions can be quite easily understood.

Several types of chemical reactions will be discussed, including acid-base interactions, oxidation, reduction, and decomposition. Understanding the basic conditions under which each of these reactions can occur, and how to describe them using balanced chemical equations, will give us a great deal of insight into the formation of various types of compounds.

Equations describing chemical reactions are most useful when they are balanced, which involves making sure that the number of atoms of each element on both sides of the reaction arrow is exactly the same. This requires that coefficients be placed in front of the various elements or compounds on each side of the equation, but we must remember to never change the actual chemical formula for any compound, because these formulas are determined by the bonding requirements between the atoms in the compound and this bonding is not arbitrary. Balancing complex chemical equations can sometimes be a little difficult, but you will be able to balance most chemical equations by following the tips given in Section 13.1 of the textbook.


Section  13.1Balancing Chemical Equations

Changes occur around us all the time. Some of these changes involve the physical properties of materials, such as density, hardness, color, and conductivity, and are referred to as physical changes. Physical changes do not alter the chemical composition of the substances involved and do not require further attention here. Chemical properties describe the ability of one substance to be transformed into another. Such transformations are referred to as chemical changes or, more often, chemical reactions. Chemical reactions involve forming and breaking chemical bonds as discussed in the last chapter. Studying the general types of such reactions gives us a better overall understanding of chemical processes.

When two or more substances interact by chemical processes to form new substances, the interaction can be described by the use of a chemical equation. The initial substances, called reactants, "react to give" or "yield" new substances called products. During the process, heat and other forms of energy are either released or absorbed. If the process can proceed in either direction, we say that the reaction is reversible. Reversible reactions progress until the rate at which the reactants are combining to form the products is the same as the rate at which the products are combining to form the reactants. This stage of a reversible reaction is known as equilibrium and is the point beyond which no net change in the amount of each reactant and product can occur unless reaction conditions change.

Chemical equations are shorthand notations that show the reactants in the chemical interaction on the left side and the resultant products on the right. An arrow drawn from left to right indicates the direction in which the reaction normally progresses. A second arrow in the opposite direction is sometimes used to show that the reaction is reversible. We should remember that no matter how rapidly a reaction appears to occur in one direction, some reverse reaction is also occurring. The rate at which the reaction progresses in either direction can be greatly affected by changing the physical conditions under which the reaction takes place.

Chemical equations are the key to understanding and analyzing many processes in chemistry. They do not express the true ratio in which substances react and form until they are balanced. Adjusting the coefficients assigned to each chemical substance in the equation is necessary to achieve an exact balance of the number of atoms on one side of the equation to a number of like atoms on the opposite side. The techniques used in balancing chemical equations, as outlined in the textbook, must be mastered fully before anyone can claim to have a basic understanding of chemistry. Chemical reactions occur in many forms, but the simplest ones, such as combination reactions, decomposition reactions, and single-replacement reactions, can be used to show the basic procedures that must be followed when balancing all types of chemical equations.


Section  13.2Energy and Rate of Reaction

Since all chemical reactions produce a change in energy, we can expect heat, light, or other types of energy to be either released or absorbed. Reactions that absorb energy are called endothermic. Reactions that release energy are known as exothermic. Chemical reactions involve the formation and breaking of chemical bonds between atoms. An endothermic reaction is one in which more energy is used in breaking the original bonds than is released when the new bonds are formed. An exothermic reaction results when the energy released on formation of the new bonds is greater than that required for breaking the original bonds.

It is sometimes necessary for energy to be supplied to start a chemical reaction. This energy is called the activation energy. In some cases the original kinetic energy of the molecules in the reactants themselves is enough to start the reaction. In other cases, heat or an electric spark must be provided. In diagram form, the activation energy is represented by an energy barrier that must be overcome before the reaction can start. (See Figs. 13.6 and 13.7 in the textbook.)

Once they have begun, exothermic reactions not only continue, but they release more energy than was originally supplied. Endothermic reactions must have the activation energy supplied and also require additional energy throughout the process. Chemical reaction rates depend on temperature, concentration, surface area, and other physical conditions. As a general rule, the higher the temperature, the faster the reaction will proceed. This is because of the increased kinetic energy of the individual molecules in higher-temperature samples of any material.

Combustion, the combination of a substance with oxygen to produce heat and light, is an example of an exothermic reaction. A match may be required to ignite a piece of paper; but once started, the combustion reaction continues to produce heat and light as long as both fuel and oxygen are available. In some cases, the activation energy for a given reaction can be lowered by the addition of a catalyst. A catalyst is defined as any substance that increases the rate of a chemical reaction but is not itself consumed or permanently changed during the reaction.


Section  13.3Acids and Bases

An acid is a substance that releases hydrogen ions when dissolved in water. A base is a substance that produces hydroxide ions in water. Although both can react with a large number of other chemicals, we are most concerned here with the acid-base neutralization reaction. In this reaction an acid combines with a base to produce water plus a salt. A salt is, therefore, defined as an ionic compound that is made up of the cation of a base combined with the anion of an acid. Some salts occur in our environment as hydrates, which means that they contain molecules of water bonded into their crystal lattices. When hydrate salts have their water molecules removed by heating, the salts that result are called anhydrous salts. Acid-base reactions are examples of double-replacement reactions in which the positive and negative ions in the reactants switch "partners" to form the reaction products. Such reactions often result in the production of a precipitate, an insoluble solid that appears when two liquids are mixed.


Section  13.4Single-Replacement Reactions

An important aspect of single-replacement reactions is the chemical activity of the various elements involved. Such reactions can only take place if the chemical activity of the newly introduced substance is higher than that of the element it is to replace in the compound that already exists. It is, therefore, necessary to know the relative chemical activity of the various elements. This information is available for some of the more common metals in Table 13.4 in the textbook, which lists these elements in increasing order of chemical activity. Such a list is known as an activity series. Notice from this table that both lithium and magnesium will replace zinc from its compounds, but copper will not. The precious metals platinum and gold are very low in the activity series, indicating that they will not react readily with other elements and therefore do not tarnish easily. This makes then attractive choices for jewelry and for the plating of critical metal contacts in electrical circuits.


Section  13.5Avogadro's Number

Experiments show that the number of elementary entities in one mole of any substance is always the same no matter what compound or element is being considered. This number, called Avogadro's number (N_A), is equal to 6.02 x 10²³ entities. It was first determined by measuring the electrical charge necessary to reduce one mole of ions to neutral atoms. Since the unit charge on a single electron was known, the number of individual electrons involved in the reduction reaction could be determined, and therefore the actual number of atoms that were in the original sample could be calculated.

The word mole (mol) is a general term that defines the quantity of any substance that contains Avogadro's number of identical components. These individual components can be atoms, molecules, electrons, ions, or any other basic constituent of matter. This means that one mole of any substance must have a molar mass that is equal to its formula mass in grams. Therefore, 1 mol of water (H₂O) has a molar mass of 18.0 g, and 1 mol of sulfuric acid (H₂SO₄) has a molar mass of 98.1 g. See problem 7 in the Paired Exercises for Section 13.5 in this study guide to see how these values were obtained. The textbook also illustrates how to convert from moles to grams or from grams to moles, and how to determine the number of entities in a given number of moles, and vice versa.

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