Compounds and Chemical Bonds - Observations in Science

Observations in Science

2.4 Compounds and Chemical Bonds

A basic picture of atoms is a good starting point for understanding the properties of polymers. But to begin to see how the observable properties of a polymer might be related to its atomic and molecular makeup, we will need to consider the connections between atoms. Which atoms are actually attached to one another? And what different types of connections—or chemical bonds—are involved? Once again we’ll begin by trying to establish some vocabulary that will help us understand compounds and the chemical bonds that hold them together.

Chemical Formulas

A chemical compound is a pure substance made up of atoms of two or more elements joined together by chemical bonds. In any compound, the atoms combine in fi xed whole number ratios. In any such combination, the resultant substance behaves differently from the atoms alone. In many compounds, atoms combine to form discrete particles called molecules. Molecules can be broken down into their constituent atoms, but the resulting collection of atoms no longer behaves like the original molecule. Other materials are composed of vast arrays or extended structures of atoms or ions but do not form discrete molecules. Alloys, metals, and ionic solids (composed of paired ions) fall into this category of chemical compounds. We’ve seen how we can use atomic symbols as shorthand notation to designate atoms. That same idea can be extended to describe the composition of either molecules or extended compounds in a simple symbolic representation.

A chemical formula describes a compound in terms of its constituent elements.

We will actually encounter two distinct types of chemical formulas: molecular formulas and empirical formulas. The molecular formula of a compound is a kind of parts list that describes the atomic composition of a molecule effi ciently. The molecular formula of the ethylene monomer from which polyethylene is produced is C₂H₄; this tells us that there are two carbon atoms and four hydrogen atoms per molecule.

The empirical formula tells us only the relative ratio between the numbers of atoms of the different elements present. Let’s consider ethylene again. The ratio of carbon atoms to hydrogen is 1:2. So the empirical formula is CH₂. When dealing with an empirical formula, it is important to realize that it does not tell how large or small an individual molecule of the compound might be; only the relative numbers of atoms of each element are given. We often emphasize this fact by writing a subscript ‘n’ on the entire formula. For ethylene, this would give us (CH₂)_n, which means that each molecule must contain some integral number of CH₂ units.

Empirical formulas, or a minor variation on them, are especially common when dealing with polymer molecules. Because polymer molecules are so large, the exact number of monomer units in a molecule is generally not very important. And in fact, the exact length of the polymer chains is often not the same for all molecules in a given sample. Instead, there is usually some range of chain lengths that will exist, Chlorine gas released in a January 2005

train accident in South Carolina led to eight deaths and forced many residents from their homes for days.

Chlorine gas released in a January 2005 train accident in South Carolina led to eight deaths and forced many residents from their homes for days.

In Section 3.5, we will learn how to fi nd the empirical formula for a compound from experimental data.

Ethylene, C₂H₄ Ethylene, C₂H₄

depending on how the polymer was actually produced. As long as the chains are all within some reasonable range of lengths, the macroscopic properties of the polymer are not affected substantially. So polymer formulas are most often written like empirical formulas. The repeating unit contributed by each monomer molecule is written in parentheses or brackets, and a subscript ‘n’ is used to emphasize that a large number of these units will be found in any individual molecule. For polyethylene, we would write the formula as ![CH₂CH₂]_n!. Here the dashes are added to stress that these units are attached end to end to build up the long chain of the polymer. For the most common forms of polyethylene, the number of monomer units (i.e., the value of ‘n’) is in the tens of thousands. We could write similar formulas for the other polymers mentioned in the opening section on pages 32 and 33.

Poly(vinyl chloride): ![CH₂CHCl]_n! Poly(vinylidene chloride): ![CH₂CCl₂]_n!

There are four rules that allow us to write most formulas that we will need in this textbook.

1. Indicate the types of atoms in the substance by their atomic symbols.

2. The number of each atom in the compound is indicated by a subscript to the right of the atomic symbol. For example, the chemical formula of ethylene, C₂H₄, tells us that each molecule contains two carbon atoms and four hydrogen atoms.

3. Groups of atoms can be designated by using parentheses. Subscripts outside these parentheses mean that all atoms enclosed in the parentheses are multiplied by the value indicated in the subscript.

4. Water molecules associated with certain compounds called hydrates are indicated separately from the rest of the compound.

Example Problem 2.2 shows how to interpret chemical formulas by inverting some of these rules.

E X A M P L E P RO B L E M 2 . 2

We cannot generally produce a polymer by simply mixing a large sample of the desired monomers. Instead, additional compounds called initiators or catalysts are almost always needed to start a polymerization. One polymerization catalyst is diethylaluminum chloride, Al(C₂H₅)₂Cl. How many of each type of atom are in a molecule of this compound?

Strategy The subscripts in a formula indicate how many atoms of each type are in the molecule. The parentheses designate a group of atoms, and the subscript associated with the parentheses multiplies each atom in the group.

Solution In each molecule of Al(C₂H₅)₂Cl, there is one aluminum atom, one chlorine atom, and two groups of C₂H₅. Each of the C₂H₅ groups contains two carbon atoms and fi ve hydrogen atoms. We multiply those numbers by two because there are two C₂H₅ groups present; so we have four carbon atoms and ten hydrogen atoms.

Discussion The number of atoms present might be easier to see if we wrote this formula as AlC₄H₁₀Cl. Right now you might feel that would be simpler. But when we write it as Al(C₂H₅)₂Cl, we are actually conveying some additional information about the way the atoms are connected to one another. Specifi cally, we are showing that the carbon and hydrogen atoms are arranged as two C₂H₅ groups and that each of these groups is attached to the aluminum atom. Later we will learn that such a C₂H₅ unit is called an ethyl group.

In the closing Insight section for this chapter, we will look more closely at the way large changes in chain length infl uence the properties of polyethylene.

Check Your Understanding A compound with the rather imposing name of 2,29-azo-bis-isobutyrylnitrile is used to initiate the growth of some polymers, includ- ing poly(vinyl chloride). If the molecular formula is C₈H₁₂N₄, how many of each type of atom are in a molecule of the compound? What is the empirical formula of this compound?

Chemical Bonding

Atoms combine to make compounds by forming chemical bonds. Several different types of chemical bonds are possible, and once we learn to recognize them, these types of bonds will help us to understand some of the chemical properties of many substances.

All chemical bonds share two characteristics. First, all bonds involve exchange or sharing of electrons. We will return to this concept often in this text as we investigate chemical reactions and properties of molecules. Second, this exchange or sharing of electrons results in lower energy for the compound relative to the separate atoms.

A chemical bond will not form, or will have only a fl eeting existence, unless it lowers the overall energy of the collection of atoms involved.

Chemical bonds can be divided into three broad categories: ionic, covalent, and metallic. Some compounds are composed of collections of oppositely charged ions that form an extended array called a lattice. The bonding in these compounds is called ionic bonding. To form the ions that make up the compound, one substance loses an electron to become a cation, while another gains an electron to become an anion. We can view this as the transfer of an electron from one species to another.

Figure 2.8 shows this concept for one ionic compound, NaCl.

Ionic compounds form extended systems or lattices of alternating positive and negative charges, such as that shown in Figure 2.9. Although the formula NaCl cor- rectly indicates that sodium and chlorine are present in a 1:1 ratio, we cannot re- ally identify an individual “molecule” of NaCl. To emphasize this distinction, we sometimes refer to a formula unit, rather than a molecule, when talking about ionic Lattices are often depicted as having

shapes that are essentially cubic, but there are actually 17 different geometric shapes, not all of which are cubic.

Lattices are often depicted as having shapes that are essentially cubic, but there are actually 17 different geometric shapes, not all of which are cubic.

Na Na

Na ClCl

Cl Cl

Na Cl

Cl Transfer of an

electron produces a pair of oppositely

charged ions.

The dots represent outer electrons.

Oppositely charged ions attract one another.

⫹ ⫺

Na Cl

⫹ ⫺ Figure 2.8 ❚ A conceptual

illustration showing the transfer of one electron from a sodium atom to a chlorine atom, forming a pair of ions (Na⁺ and Cl^–). Once electron transfer takes place, coulombic force draws the ions together.

compounds. The formula unit is the smallest whole number ratio of atoms in an ionic compound.

Metals represent another type of extended system, but here the chemical bonding is totally different. In metals, the atoms are once again arranged in a lattice, but positively and negatively charged species do not alternate. Instead, the nuclei and some fraction of their electrons comprise a positively charged “core” localized at these lattice points, and other electrons move more or less freely throughout the whole array. This is called metallic bonding. Metallic bonding leads to electrical conductiv- ity because electrons can move easily through the bulk material. Figure 2.10 shows a schematic illustration of the concept of metallic bonding.

When electrons are shared between pairs of atoms rather than donated from one atom to another or mobile across an entire lattice, we have covalent bonds. In covalent bonds, electrons are usually shared in pairs. Two electrons (and sometimes four or six) are located between two nuclei and the sharing leads to an attraction between the nuclei. The long chains in all polymers are formed by covalent bonds in which electrons are shared between adjacent carbon atoms. Smaller, more familiar molecules such as water, carbon dioxide, and propane are simpler examples. All three types of chemical bonds will be discussed in much greater detail in Chapters 7 and 8.

Polymer molecules are built up by the successive addition of monomers to form characteristic long-chain backbones. The bonds that hold the monomers to one another, as well as the bonds between atoms within each monomer unit, are covalent bonds. But ionic bonding is important in many compounds that are used to initiate or sustain the reactions needed to grow a polymer.

Figure 2.9 ❚ Two different representations of the NaCl crystal structure are shown. In each case, the green spheres represent chloride anions, and the gray spheres denote sodium cations. The view on the left emphasizes the positions of the ions, and that on the right better illustrates their relative sizes. In a macroscopic salt crystal, additional ions would simply extend this structure, repeating the same alternating pattern.

Na+

Cl–

Figure 2.10 ❚ In this simple conceptual picture of metallic bonding, each metal atom contributes one or more electrons to a mobile “electron sea.” The ability of the electrons to move freely through this “sea” allows the metal to conduct electricity. Here the blue area depicts those mobile (or “delocalized”) electrons, and the red circles represent the positively charged “cores” of the individual atoms.

+ + + + + + + +

The nucleus and inner electrons provide a positively charged “core.”

The outer electrons form a “sea” of negative charge surrounding the positive cores.

Water, H₂O

Propane, C₃H₈ Carbon dioxide, CO₂

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