Chemistry of the Representative Elements

Submitted by ChemPRIME Staff on Thu, 12/16/2010 - 14:32


So far we have been devoted to explaining fundamental concepts and principles such as the atomic theory, electronic structure and chemical bonding, intermolecular forces and their effects on solids, liquids, or gases, and classes of reactions such as redox or acidIn Arrhenius theory, a substance that produces hydrogen ions (hydronium ions) in aqueous solution. In Bronsted-Lowry theory, a hydrogen-ion (proton) donor. In Lewis theory, a species that accepts a pair of electrons to form a covalent bond.-baseIn Arrhenius theory, a substance that increases the concentration of hydroxide ions in an aqueous solution. In Bronsted-Lowry theory, a hydrogen-ion (proton) acceptor. In Lewis theory, a species that donates a pair of electrons to form a covalent bond.. It is well to remember, though, that all these concepts and principles have been developed and used by chemists in order to better understand, recall, and systematize macroscopic laboratory observations. In other words, although the concepts we have described so far have their own inherent beauty as great ideas, they are primarily important because they reduce the quantity of memorization which is necessary to master descriptive chemistry, allowing people to recall facts that otherwise might be forgotten.

Such a reductionThat part of a chemical reaction in which a reactant gains electrons; simultaneous oxidation of a reactant must occur. in memory workA mechanical process in which energy is transferred to or from an object, changing the state of motion of the object. is only possible if you know how to apply principles to specific elements, their compounds, and the reactions they undergo. This is not as easy as it may seem, but neither is it impossibly difficult. We will describe some of the chemistry of the representative elements, showing as we do so how their properties may be rationalized on the basis of concepts and principles. Explanation of these properties is organized according to the concept of periodicity, with each subsequent section corresponding to one of the eight groups of representative elements. As you read, you should try to see why a certain reaction occurs, as well as what actually happens, and instead of memorizing each specific equation, you should try to organize the chemistry of these elements according to the generalizations you have already learned.

At this point it is useful to look and consolidate some of the general trends observed for the representative elements. First, metals on the far left of the periodic tableA chart showing the symbols of the elements arranged in order by atomic number and having chemically related elements appearing in columns. are good reducing agents, while nonmetals on the far right (excluding noble gases) are strong oxidizing agents. Thus these elements are quite reactive, especially when one from the left combines with one from the right. Hydrogen compounds (hydrides) of the alkali and alkaline-earth metals contain strongly basic H ions and produce basic solutions. Toward the middle of the periodic table acid-base properties of hydrogen compounds are harder to predict. Some, like CH4 in Group IVA, are neither acids nor bases, but others, like NH3, have lone pairs of electrons and can accept protons. Protons can be easily donated and are acidic only when they are bound to halogens or oxygen.

The acidic behavior of oxides also increases from left to right across the periodic table and decreases from top to bottom. The situation is complicated by the fact that the higher the oxidation stateA formally defined charge that an atom in a compound or ion would have if the compound or ion consisted entirely of monatomic ions. Based on a Lewis diagram, the charge that an atom would have if all electrons in bonds were assigned to the more electronegative atom or divided equally between atoms of the same electronegativity. of an atom, the more covalent its oxide and the more acidic it will be. Thus SO3 dissolves in water to give a strong acidAn acid that ionizes completely in a particular solvent., while SO2 gives a weak one. Taking account of both of these trends, one can fairly well predict which oxides are likely to be basic, which amphotericA substance that can behave as either an acid or a base. Examples are aluminum metal, aluminum hydroxide, and zinc hydroxide., and which acidic.

General rules can also be used to predict which oxidationThat part of a chemical reaction in which a reactant loses electrons; simultaneous reduction of a reactant must occur. states will be most common. On the left of the periodic table the group number gives the most common oxidation state. From group IIIA on, the group number minus 2 (for the ns2 electrons) is also common, especially for elements near the bottom of the table. .The group number is a good choice when an element combines with a highly electronegative element, but the group number minus 2 is more common when one element is bonded to another element of intermediateIn chemical kinetics, a species that is formed in an early step in a reaction mechanism and then consumed in a later step; evidence of existence of an intermediate may be important for the interpretation of a rate law. electron-withdrawing power. For example from the Chalcogens, SF4 and SF6 are both stable, but SF4 is the most stable sulfur chloride.

From group VA on to the right of the table, the group number minus 8 is an important oxidation state, especially for the first member of a group. Oxidation numbers other than those already mentioned usually differ by increments of 2. For example, chlorine exhibits –1, +1, +3, +5 and +7 oxidation states in stable compounds, and sulfur is found in –2, +2, +4, and +6 states.

In conclusion, do not overlook the forest by concentrating too much on individual trees. Look for and try to understand and use the generalizations and correlations that have been developed in this chapter. If you do this, you will retain the facts presented here much more efficiently.