1.9: Chapter Summary

An atom consists of a positively charged nucleus surrounded by enough negatively charged electrons to yield zero net charge. The nucleus is constructed from positively charged protons and neutral neutrons, each of mass approximately 1 amu. The mass of an electron is approximately 1/ 1836 the mass of a proton; the charge on an electron is equal but opposite in sign to the charge on a proton. The total number of protons in the nucleus (and electrons in a neutral atom) is the atomic number, Z. The total number of both protons and neutrons is the mass number, and the mass of the atom, in atomic mass units, is its atomic weight. The atomic weight is always slightly less than the sum of masses of the particles that go into making an atom, because mass is converted to energy and lost when the atom is formed.

All atoms with the same number of protons, and therefore the same atomic number, are classified as the same element and represented by a one- or two-letter symbol. Atoms of the same element with varying numbers of neutrons are called isotopes of the element. Isotopes are identified by placing the mass number as a superscript to the left of the symbol of the element (e.g., ³⁷Cl). The atomic number is sometimes added as a subscript (e.g., Cl), although it is actually not necessary since the element's name and atomic number are known from the symbol. Each isotope of an element has its own atomic weight, and the natural atomic weight is the weighted average of these isotopic values, the weighting being according to the natural abundance of each isotope.

A collection of atoms held together by chemical bonds is a molecule Usually, but not always, the bonding in a molecule can be explained in terms of electron pairs, each holding two atoms together. Such an electronpair bond is a covalent bond. The sum of the atomic weights of all the atoms in a molecule is its molecular weight. Although atoms in different molecules are not directly bonded to one another, all molecules are slightly "sticky," and are attracted to other molecules. These van der Waals attractions will make the molecules of a gas adhere to one another to form a liquid if the temperature falls low enough, and make the molecules of a liquid fit together in a regular crystalline array in a solid if the temperature falls lower still. The temperatures at which these two transitions occur are the boiling point, T_b , and the melting point, T_m' respectively.

If two atoms differ in their intrinsic electron-pulling power or electronegativity, then the electron pair of the bond between them will be shifted toward the atom with the greater attraction, giving it a negative charge and the other atom a positive charge. The bond, and molecules that contain such bonds, are said to be polar. Polar molecules can attract one another, and they can also attract positively and negatively charged ions. Melting and boiling points of polar molecules are higher than would be expected from van der Waals attractions alone, because their polarity provides a second type of intermolecular attraction.

Atomic and molecular weights are measured on a scale of atomic mass units (amu), where 1 amu is defined as exactly one-twelfth of the mass of a ^I2C atom. A quantity of a chemical substance (atoms, molecules, or ions) equal to the atomic weight expressed in grams is defined as 1 mole of that substance. One mole of any substance-atoms, molecules or ions-contains the same number of particles of that substance. This property makes the mole a useful means of counting out particles merely by weighing them. The units of atomic and molecular weights are either grams per mole or amu per molecule (or atom).

Some atoms, those of metals in particular, have a weak hold on their electrons and can lose one, two, or more electrons to become positively charged ions, or cations. Many nonmetals or groups of atoms can acquire one or more negative charges to become negatively charged ions, or anions. A salt is a compound of the relative number of cations and anions that will produce zero overall charge. Common table salt, NaCl, contains equal numbers of Na⁺ and Cl^- ions. The pulling away or outright removal of electrons is termed oxidation, and the addition to or shifting of electrons toward an atom is reduction. Since electrons are never created or destroyed in chemical reactions, whenever one substance is oxidized, some other substance must be reduced.

Simple anions made by adding electrons to single atoms have names ending in -ide, as chloride, Cl^-, and sulfide, S^2-, ions. For complex ions of a nonmetal atom with oxygen, the higher and lower oxidation state ions are differentiated by the suffixes -ate and -ite. The oxidation state of a metal cation (see Chapter 10) is indicated by a Roman numeral after the name of the metal, as in Fe³⁺, iron(III), or by the suffixes -ic and -ous.

Although salts do not have separate molecules and, strictly speaking, cannot have molecular weights, they do have chemical formulas that express their overall composition in the simplest possibJe way. The weight of 1 mole of these atoms is the formula weight of the salt, but it is customary to refer to this as the salt's "molecular weight." Thus magnesium chloride has one Mg²⁺ ion for every two Cl- ions, a net charge of zero, a chemical formula of MgCl², and a molecular weight of 95.211 g mole^-1.

The coordination number in a complex ion or molecule is the number of atoms or chemical groups bonded directly to the central atom. These bonding groups can be simple ions such as O^2- and Cl^- or molecules such as ammonia (NH₃) and water (H₂O). The maximum coordination number for a given central atom depends on the size of the atom and the size of its surrounding groups. The most common coordination numbers are 2, 3, 4, and 6.

Salts have higher melting and boiling points than molecular substances, because heat energy must be supplied to break apart the stable crystal lattice, and even more heat energy is required to force positive and negative ions to pair off and share electrons in neutral molecules that can go into a gas phase. Salts often dissolve readily in water, however, because polar attractions by the water molecules can compensate for the attractions of other ions in the crystal. Ions surrounded by polar water molecules in solution are said to be hydrated. Gasoline and similar nonpolar liquids cannot dissolve salts because they cannot hydrate (or solvate, if the solvent is other than water) the ions.

If a current of electricity is passed through molten salt or a salt solution, the current is carried by ions migrating in opposite directions. At the cathode, where electrons enter the salt medium, metal cations can be reduced to pure metal. At the anode, where electrons flow out of the salt and back into the external circuit, anions can be oxidized to liberate pure nonmetallic elements. This is the process of electrolysis. Faraday found a quantitative relationship between the amount of charge passed through a cell and the amount of chemical change produced: 96,485 coulombs of charge will bring about 1 mole of a change that involves one electron per ion. The quantity, 96,485 coulombs, is simply the charge on 1 mole of electrons and is called 1 faraday () of charge.

Electrons as separate particles were studied by physicists interested in low-pressure gas discharges under high voltages. Cathode rays consist of a beam of electrons stripped away from the gas atoms. J. J. Thomson showed, by means of deflecting magnetic and electrostatic fields, that the cathode rays were made of negatively charged particles, and he measured the charge-to-mass ratio of the particles. R. A. Millikan completed the process, in his oil-drop experiment, by successfully measuring the charge on the electron. This, combined with Faraday's results, led to the calculation of Avogadro's number, the number of electrons in a faraday of charge, or the number of particles in a mole of any substance. The mass spectrometer, a descendant of Thomson's gas-discharge tubes, is a modern analytical tool and a means of finding the charge-to-mass ratio for any atomic or molecular species that can be given a charge.

Contributors and Attributions

• R. E. Dickerson, H. B. Gray, and G. P. Haight, Jr. Content was used from "Chemical Principles", an introductory college-level text for General Chemistry with permission of the Caltech library and Harry B. Gray, on behalf of the authors.