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

1. Understand the law of conservation of mass as it applies to chemical reactions.
2. Use the law of definite proportions and see how this law leads to the assignment of unique formula masses to compounds.
3. See the ramifications resulting from the postulates of Dalton's atomic theory.
4. Describe the processes of ionic, covalent, and hydrogen bonding and be able to write the formulas for compounds formed by the first two processes.
5. Name the compounds made from metals that can bond in more than one ionic form.

Discussion

Chemical bonds are an important part of the structure of nearly everything we see around us. Electromagnetic forces among the electrons and nuclei of the atoms that make up molecules are the keys to understanding the various combinations of atoms that can be created. Understanding of these forces was not possible until modern mathematical and computer techniques were developed to model these often complex processes.

Careful study shows that the total mass of the materials involved in any chemical reaction is the same after the reaction is complete as it was before the reaction began, even though the structures of the resulting substances are often quite different from those of the interacting materials. Such observations have been formulated into the law of conservation of mass, which states that there is no detectable change in the total mass during a chemical reaction. This results from the fact that atoms are not created or lost during chemical reactions, but are simply rearranged into different configurations.

The law of definite proportions requires that each compound formed from specific types of atoms must always combine in a fixed and predictable way. This means that different samples of a pure compound always contain the same elements in the same proportion by mass. Using this idea, John Dalton was able to postulate an atomic theory based on the idea that each element is composed of small, indivisible particles called atoms, which are identical for that element but are different from the atoms of other elements. He further stated that chemical combinations are the bonding of a definite number of atoms of each of the combining elements, and that no atoms are gained or lost during a chemical reaction.

The formation of compounds involves the combination of individual atoms using several different types of chemical bonding. Ionic bonds result from the complete transfer of one or more valence electrons from one atom to another, thereby forming charged ions that are then held together in a crystal lattice by attractive electric forces. Covalent bonding involves the sharing of valence electrons, and the hydrogen bond is a special kind of dipole-dipole interaction involving small, highly electronegative atoms that are already covalently bonded to hydrogen atoms.


Section  12.1Law of Conservation of Mass

The total mass of the reacting materials in any chemical process is always measured to be exactly equal to the mass of the products of the reaction (plus any unreacted reactants), even when our most sensitive balances are used. This equality if known as the law of conservation of mass. It is very useful in predicting the amounts of substances necessary to make a certain quantity of a compound, or in determining how much of such a compound will be formed using a given amount of reactant.

Because a specific chemical formula exists for each compound, it is possible to calculate a formula mass for the formula units of that compound. This is done by taking the sum of all of the atomic masses that appear in the chemical formula for that compound. We must multiply the atomic mass of any atom having a subscript in the chemical formula by that subscript before that mass is added to the total. The formula mass can be either for an element, in the case of a single atom or diatomic molecule, or for a very complex compound such as some of the organic molecules that contain hundreds or even thousands of individual atoms. For example, the formula mass of the diatomic element O₂ is simply 2(16.0 u) = 32.0 u, whereas the formula mass for the more complex compound H₂SO₄ is 2(1.0 u) + (32.1 u) + 4(16.0 u) = 98.1 u. This means that, on the average, one molecule of diatomic oxygen gas has a mass of 32.0 u, and a single molecule of sulfuric acid has a mass of 98.1 u. (The atomic mass for each element can be found in the periodic table in the front of your textbook.)


Section  12.2Law of Definite Proportions

Careful examination of the masses of interacting substances has led chemists to another important basic law. No matter what size sample of a pure compound is analyzed, it will always contain the same elements in the same proportion by mass. This is called the law of definite proportions. To explain the law, consider that a molecule of a specific compound has a fixed proportion of atoms of the elements of which it is composed. Also, each atom of an element has a specific atomic mass. Therefore, it is really quite easy to see that a molecule of a compound must have a fixed proportion by mass (as well as by atoms), and this proportion will remain the same no matter how many molecules of that compound make up the sample.

Sometimes when chemical reactions occur, the elements (or compounds) that are combining are not present in exactly the correct proportions for the compounds being formed. In this case one of the elements will be completely used up, whereas some amounts of the other elements may be left in their original forms. The element completely used up is called the limiting reactant, and the ones not completely used are called excess reactants. The sum of the masses of the initial reactants present must be equal to the sum of the masses of the resulting compounds plus the masses of any excess reactants remaining. This is true because both the law of conservation of mass and the law of definite proportions must be satisfied in every reaction.


Section  12.3Dalton's Atomic Theory

According to what we have just seen, the basic structure of matter must be quite complex. To explain this structure better, John Dalton developed an atomic theory that accounted for both the law of definite proportions and the law of conservation of mass. The three postulates of this theory are:

1. Elements are composed of small, indivisible particles called atoms that are identical for any given element but are not alike when different elements are compared.
2. Chemical bonding is the combining of a definite, small number of atoms of each of the combining elements to make up molecules.
3. In a chemical reaction no individual atoms are lost or gained. This means that all atoms retain their individual identities but are sometimes rearranged into new, well-defined structures.

Dalton's postulates also can be applied to compounds composed of ions, in which case the ions occur in definite proportions by number and the resultant "molecules" can be represented by a formula unit. A formula unit is defined as the smallest combination of ions that gives the correct formula for that compound. The term formula unit can be applied to ion combinations, entire molecules, or even single atoms, depending on the makeup of the matter under study.

Dalton's atomic theory was proved correct by the experimental verification of the fact that proportions in which the same elements combine into two or more different compounds also follow simple ratios. This idea was proposed by Dalton in his law of multiple proportions, which states that if two elements combine to form more than one compound, the various masses of one element, combined with a constant mass of the other element in these compounds, will be related by a ratio of small whole numbers. To understand this law better, study the multiple proportions of the gases carbon monoxide and carbon dioxide as discussed in Section 12.3 in the textbook. When the predictions made by this law were experimentally verified, Dalton's atomic theory was placed on extremely solid ground. His simple atomic model of matter has become the cornerstone of modern chemistry.


Section  12.4Ionic Bonding

Atoms are held together in molecules by electrical forces. These forces form bonds that can arise in several different ways. We will study some of these types of chemical bonds in this and subsequent sections.

The first type of chemical bonding we will consider is ionic bonding. Ions are formed in chemical reactions by the complete transfer of one or more electrons from one atom to another. The number of electrons gained or lost by an atom can be predicted from its electron configuration and the fact that an especially stable valence configuration is attained when the electron shell structure is similar to that of the noble gases. Once the electrons have been transferred, the resulting ions are held together by the mutual attraction between unlike electrostatic charges.

In the noble gases, the valence electron shell is completely filled with eight electrons. The only exception is helium, whose valence electron shell is completely filled with two electrons. Noble gases generally do not readily form chemical bonds with other elements. Combined with the similarity in chemical properties among all elements that have the same number of valence electrons, the inert nature of the noble gases leads us to what is commonly called the octet rule. The octet rule states that when forming compounds, atoms tend to gain, lose, or share electrons in such a way that they achieve an electron configuration similar to that of the noble gases; that is, they tend to transfer or share electrons until eight electrons (an octet) make up their outermost electron shells.

Even though there are some exceptions, the octet rule is an effective tool in predicting and explaining most of the combinations of atoms that occur during chemical reactions. When the electrons are completely transferred (lost or gained by individual atoms), the process is called ionic bonding. Positive ions are called cations, and negative ions are called anions. The number of valence electrons in an atom of an element generally determines how many electrons will be transferred when an ion is formed and, therefore, what the charge on each ion will be in units of the fundamental electron charge.

Some metals are capable of forming ions with more than one electrical charge. Such metals (Fe is one example) can therefore form more than one compound with a given nonmetal. The chemical formulas for such compounds show this clearly (FeCl₂ and FeCl₃, for example), but both of these compounds would be called iron chloride using the basic conventional system for naming compounds. To avoid such confusion, the Stock system of using a Roman numeral to indicate the metal's ionic charge in a given compound was developed. For the above examples, Fe has a 2+ charge in the first formula and a 3+ in the second, so these two compounds are called iron(II) chloride and iron(III) chloride, respectively.

Ionic compounds always form in such a way that the resulting formula units have an overall neutral charge. This idea will be helpful when we learn to write and balance equations for chemical reactions. The atoms of each reacting element that make up a compound must combine in such a way that the total electrical charge of the formula unit formed is zero and all atoms must be isoelectronic with a noble gas. (Isoelectronic means having the same electron configuration.)

Lewis symbols are a convenient way to represent the valence shell electrons of atoms. Table 12.1 in the textbook shows how Lewis symbols are drawn for the representative elements. Lewis structures are used to show the changes in the valence shell when electrons are transferred (or shared). Understanding this notation is a prerequisite to dealing with ideas presented in later chapters, so please study it carefully.


Section  12.5Covalent Bonding

So far we have only considered the total transfer of electrons for one atom to another, that is, ion formation. It is possible for atoms simply to share electrons and thereby attain a noble gas configuration. This leads to a second type of bonding called covalent bonding. Electron sharing can be quite flexible and results in the formation of single, double, and even triple covalent bonds, depending on how many electrons are shared by the participating atoms. Again, the Lewis structure notation is useful and can be written either by using dots to show individual valence electrons or by using a long dash to indicate each shared pair of electrons. We should remember that the total number of electrons associated with a given atom, including those shared with adjacent atoms, must conform to the octet rule. In general, compounds formed only of nonmetals are covalently bonded, whereas those formed from active metals and nonmetals (or metals and polyatomic ions) are ionically bonded.

Sometimes both covalent and ionic bonds are present in a single compound, as in the case of polyatomic ions that combine with other ions by way of the ionic bond but are themselves held together by electron sharing, which involves covalent bonds. In certain molecules the sharing of electrons is unequal, so that the shared electrons spend more time associated with one atom than another. Such bonds are known as polar covalent bonds, and the amount of sharing and thus the degree of polarity associated with the molecule depend on the difference in the electronegativities of the atoms involved in the covalent bonding process. If little or no electronegativity difference (D EN) exists, the bonds will be nonpolar.


Section  12.6Hydrogen Bonding

Hydrogen bonding is a special kind of dipole-dipole interaction that occurs between adjacent molecules and not within individual molecules. It is possible because a small hydrogen atom has such a highly concentrated positive charge that it forms a strong electrical attraction with small, very electronegative atoms of oxygen, fluorine, and nitrogen. This bond is responsible for some of the unique and interesting properties of water, without which life as we know it on Earth would not exist. This bond is also responsible for the special properties of some other compounds, such as HF and NH₃.

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