BOND TYPES: THE CLASSIFICATION OF SUBSTANCES

Transcription

1 BOND TYPES: THE CLASSIFICATION OF SUBSTANCES Every (pure) substance has a unique set of intrinsic properties which distinguishes it from all other substances. What inferences, if any can be made from a substance s properties? Since the interactions between the atoms, ions, or molecules in a sample of matter largely determine its behavior, the types of interactions, generally called bonding, can be correlated with its properties. In this experiment, you will examine some of the properties of a number of substances and classify each substance on the basis of your observations. The properties of compounds do not resemble the properties of the component elements. Consider sodium chloride, NaCl. Sodium, Na, is a soft, lustrous, low-melting metal which is so reactive that it tarnishes immediately when exposed to air and reacts explosively with water. Chlorine, Cl 2, is a greenish yellow gas which is very corrosive and poisonous. Yet sodium chloride or table salt, a compound of sodium and chlorine, is a hard, brittle, high-melting, nonvolatile, transparent crystalline solid that is nontoxic, dissolves quietly in water, and is unreactive toward most other substances. But even though the compound does not inherit the properties of its parent elements, nevertheless their properties do determine the type of chemical bonding involved, which in turn largely determines the properties of the compound. Although there are no sharply defined boundaries, chemical bonds can be classified into 3 main types: ionic bonds, covalent bonds, and metallic bonds. Following is a brief discussion of each type of bond and the general properties found in typical substances in which the bond type occurs. Ionic Bonds and Ionic Compounds A metal atom and a nonmetal atom interact by electron transfer. The metal atom loses one or more electrons forming a positively charged ion called a cation. The nonmetal atom gains one or more electrons forming a negatively charged ion called an anion. The compound is held together by the electrost atic attractions between the positive and negative ions, called the ionic bond. The cations and anions are arranged in a rigid structure called a crystal lattice in which each positive ion is surrounded by a number of negative ions and each negative ion is surrounded by a number of positive ions. The pattern is determined by the ions sizes and charges. A compound made up of ions - a compound involving any ionic bonds - is called an ionic compound. Ionic compounds are usually hard, brittle, high-melting, nonvolatile, transparent crystalline solids. They usually melt at temperatures of about C. (Some ionic compounds or salts which contain water of crystallization may lose it at temperatures lower than 300 C.) Many ionic compounds dissolve in water because of the strong attractions between the ions and the highly polar water molecules. The water solutions are good conductors of electricity unless they are very dilute. In some ionic compounds, the attractions between the ions are so much larger than the attractions between the ions and water molecules, that the compound does not dissolve in water sufficiently to form a highly conducting solution. With nonpolar solvents such as hexane, the attractions between the non-polar molecules of the liquid and the ions are too weak to bring about solution. The polar organic solvent alcohol is a better solvent Schwartz, R./PC/1342B/I 1/00

2 for ionic compounds than hexane but not as good as water. When an ionic compound is melted, the ions are freed from their positions in the lattice and can conduct electricity by moving. (A moving charge is a current of electricity.) Covalent Bonds and Covalent Compounds 2 A nonmetal atom and another nonmetal atom interact by sharing a pair, or sometimes, several pairs of electrons. The shared pairs of electrons are called covalent bonds. A compound involving only covalent bonds is called a covalent compound. In most covalent compounds, the atoms which share electrons form discrete, uncharged particles called molecules. The unequal sharing of electron pairs results in an unbalanced or polar covalent bond. An unsymmetrical arrangement of polar bonds in a molecule makes the molecule polar, with an unbalanced distribution of electrical charge. Altho ugh the covalent bonds between the atoms in a molecule are strong bonds, the attractions between the molecules, especially nonpolar molecules, are usually much weaker than the attractions between positive and negative ions. Covalent compounds of low molecular weight (less than about 100 amu) are usually gases or volatile liquids. Those of higher molecular weight are soft, crumbly, volatile, low-melting solids. These usually melt below 300 C. (For substances with similar molecules, melting points and boiling points increase with increasing molecular weight.) A substance with polar molecules melts and boils higher than a nonpolar substance of similar molecular weight. A covalent substance is more likely than an ionic substance to be soluble in an organic solvent. Polar covalent substances are soluble in polar solvents such as alcohol or water and nonpolar substances are soluble in nonpolar solvents such as hexane. The attractions between water molecules and nonpolar molecules that the two kinds of molecules do not mingle and nonpolar covalent substances are not appreciably soluble in water. Solutions of covalent substances are usually poor electrical conductors or nonconductors. But a few highly polar covalent substances form water solutions which are very good conductors. These substances or, more properly, their water solutions, are called strong acids. Some examples are nitric acid, HNO 3, hydrochloric acid, HCl, and sulfuric acid, H 2 SO 4. Melted covalent compounds are usually poor conductors of electricity or nonconductors. A few covalent compounds such as silicon dioxide, SiO 2 (quartz, sand) and silicon carbide, SiC (Carborundum), are called macromolecular substances or covalent network solids because a network of covalent bonds joins together all the atoms in a crystal into a giant molecule. Melting such a crystal involves breaking of covalent bonds. Melting a molecular solid does not involve the breaking of covalent bonds. As the molecular solid is heated, thermal energy increases the motions of individual molecules. When their energies are large enough to overcome the relatively weak attractions between molecules (Van der Waals forces), the molecular crystal melts. In contrast, macromolecular solids are very hard, brittle, and very high-melting (above 1000 C). They are insoluble in all solvents and, because they have no loosely bound electrons, are nonconductors of electricity. (Some of them find uses as insulators.) Some substances have a bond system that must be considered intermediate between ionic and covalent. A principally ionic compound may involve a very small, highly charged cation and a large anion. The cation so polarizes the electron cloud of the anion that the bond is described as partially covalent. A principally co valent compound may be so highly polar that the bonding is described as partially ionic. Some authors use the term ionic-covalent to describe a bond that is principally ionic but partially covalent and covalent-ionic to describe a bond that is principally

3 3 covalent but partially ionic. Metallic Bonds and Metallic Solids Metal atoms interact with other metal atoms to form metallic bonds, the bond type found in pure metals, alloys, and certain intermetallic compounds. The metal cations from a lattice and the valence electrons form a surrounding sea or gas. These valence electrons are mobile and delocalized; there are apparently no associations of particular electrons with the particular cations. Because of the mobility of the valence electrons, metals are good conductors of electricity in both the solid and liquid state. Because of the delocalization of the negative charge, metal crystals are easily deformed and reshaped making metals malleable and ductile. The energy states of the electrons in the gas are virtually continuous so that all the visible light frequencies are both absorbed and reradiated, giving metals their characteristic shiny or lustrous appearance. Certain metals may react with liquid solvents but metals do not simply dissolve in solvents the way, for example, salts do. Their melting points range from quite low (-39 C for mercury, Hg) to very high (3415 C for tungsten, W). Their hardness varies from that of potassium, K, and cesium, Cs, which are about as hard as a firm cheese to that of special purpose aloys so hard that special tools are required to cut them. On the next page is a summary of properties typical of members of each of the 4 main classes of solids.

4 4 COMPARISON OF THE FOUR MAIN TYPES OF CRYSTALLINE SOLIDS Ionic Molecular Macromolecular or Covalent Network Metallic Examples NaCl, MgSO 4, Ca(NO 3 ) 2, AlF 3 solid CO 2, S 8, P 4, naphthalene (C 10 H 8 ), paradichlorobenzene (C 6 H 4 Cl 2 ) diamond, gemstones, ceramics, Carborundum pure metals, alloys What occupies the lattice points in the crystal? cations and anions individual molecules atoms covalently bonded to one another metal cations (The valence electrons are delocalized.) What is the strongest force binding them in the lattice? the ionic bond Van der Waals forces (intermolecular attractions) the covalent bond the metallic bond Hard or soft? hard soft very hard variable Brittle or malleable? brittle crumbly very brittle malleable High or low melting point? high (usually C) low (usually under 300 C) very high (usually over 1000 C) variable (-39 C Hg; 3415 C W) Good conductor? no (unless melted) no no (insulators) excellent often soluble in water; usually insoluble in nonpolar solvents polar substances soluble in polar solvents, nonpolar in nonpolar solvents insoluble insoluble

5 5 Procedure You are to estimate the melting points, solubilities in 3 different solvents, and conductivity as a solid and a water solution of various solids, both known and unknown. On the basis of your observations, you are to classify each one as a member of one of the 4 main types of solids (ionic, molecular, macromolecular, or metallic) and cite supporting evidence for your choice. Melting Point Place a small amount of the solid to be tested in a test tube and heat the test tube in a beaker containing boiling water. If the solid melts in the boiling water bath, its melting point is at or below 100 C, the boiling point of water. If the substance does not melt in the boiling water bath, place a sample of it in a crucible supported on a triangle and heat it gently heat with the Tirrill burner. If the solid melts, its melting point can be considered to be at or below 300 C. If the substance does not melt when heated gently, heat the crucible at maximum heat with the Tirrill burner. (This will require repositioning the crucible, or better yet, a separate ring stand assembly.) If the substance melts, its melting point can be considered to be at or below 600 C. If the substance does not melt at maximum heat with the Tirrill burner, its melting point is above 600 C. After these tests, you will be able to place the substance s melting point in one of the following categories: Record the results. room temperature above 600 C C C C You are to test the solubility of each substance in each of 3 solvents: 1. water, 2. A polar organic solvent such as alcohol, and 3. A nonpolar organic solvent such as hexane For each solvent, test the solid as follows. Half fill a medium sized test tube with the solvent to be tested. If the solid is not finely ground, crush it to hasten dissolving. Add a very small portion of the solid (half the size of a match head or less) to the liquid and shake and stir vigorously at intervals for a period of several minutes (be sure to stopper the test tube with a cork before shaking). If the sample dissolves, add another small portion and repeat the shaking and stirring. Continue until no more solid dissolves. Repeat the test with each substance with each solvent. Record the results. Conductivity

6 6 Use the meters provided to test the conductivity of each solid sample. Record the results. Before you begin to test the conductivity of the water solutions, test the conductivity of distilled water for comparison. In estimating the conductivity of the water solution, you will have to take into account the solubility of the dissolved substance. A very dilute solution of an ionic compound is not as good a conductor of electricity as a more concentrated solution of the same substance would be. Record the results. Classification On the basis of the properties you have been able to observe, classify each known and each unknown as a member of one of the 4 main groups of solids and list the evidence which supports your choice of classification. Results and Notes: Melting Points Conductivity

7 7 Substance Appearance MP Range Conductivity CaCl 2 H 2 O Alc Hex Classification Urea, H 2 NCONH 2 Silicon dioxide, SiO 2 Naphthalene, C 10 H 8 Sucrose, C 12 H 22 O 11 KBr Zinc, Zn Camphor NaC 2 H 3 O 2 Silicon carbide, SiC Cholesterol Sulfur, S 8 NaHCO 3