Chemistry

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Atoms and Molecules

The concepts of atoms and of the groups of linked atoms called molecules are the foundation of all chemistry (see Atom). An atom is the smallest unit of an element that has the properties of the element; a molecule is the smallest unit of a compound or the form of an element in which atoms bind together that has the properties of the compound or element.

The idea of atoms is an old one. Greek philosopher Leucippus and his student Democritus appear to have originated the idea during the 4th and 5th centuries bc. According to them, matter consisted of small, indivisible particles called atoms. All atoms were made of the same basic material, but neither philosopher stated what this material was. The atomic theory was developed further by another Greek philosopher, Epicurus, who added the property of weight to the atoms and attributed a horizontal, as well as a vertical, motion to them in order to explain how atoms combine to form matter. These ideas were restated by Roman poet Lucretius in the 1st century bc.

In the 18th century ad, English schoolmaster John Dalton developed his well-known atomic theory, which explained the laws of definite and multiple proportions. Convincing proof that atoms exist, however, has only been generated since 1900. Much, but not all, of this proof came from the study of radioactivity and of energetic particles. When Lucretius watched dust particles dancing in a sunbeam and said that they were being battered by the invisible blows of restless atoms, he was basically right. True, most of the dancing was caused by air currents, yet even in still air, specks of dust or smoke are in constant motion, as are minute particles suspended in water. This constant random movement of particles is the so-called Brownian motion. Two thousand years after Lucretius, French scientist Jean-Baptiste Perrin, armed with a microscope and, more importantly, a mathematical theory, measured the random motions of suspended dye particles and calculated the number of the invisible molecules whose collisions were causing the visible dye particles to move. This way of counting molecules helped substantiate the existence of atoms and molecules.

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The Structure of Atoms

The present picture, or model, of atoms is as follows. An atom has a central nucleus, which is very small compared with the rest of the atom and contains most of the atomic mass (or weight). The nucleus carries a positive electric charge and is surrounded by a diffuse shell, or cloud, of negatively charged particles called electrons. The diameter of the atom is determined by the size of this electron cloud and is about 10^-8 cm (3.94 x 10^-9 in), whereas the nucleus is about 10^-12 cm (3.94 x 10^-13 in) in diameter. The size ratio of the atom to the nucleus is 10,000 to 1.

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The simplest atom of all, hydrogen, has one particle—called a proton—in its nucleus. The mass of a proton is 1836 times the mass of an electron. A proton carries a positive electric charge with an assigned value of +1, and an electron carries a negative electric charge with an assigned value of -1. The atoms of other elements have more than one proton in their nucleus, and, in addition, other elements have another kind of nuclear particle called a neutron. The neutron has nearly the same mass as the proton, but the neutron has no electric charge.

The number of protons in the nucleus of an atom of a certain element determines the element’s atomic number. The number of protons in the nucleus can be determined by measuring the positive charge on the nucleus. For example, an atom with a nuclear charge of +26 has 26 protons in its nucleus and must be iron. To be electrically neutral, an atom of iron must have 26 electrons surrounding the nucleus.

The total number of protons and neutrons in a nucleus is called the mass number, since these particles account for almost the entire mass of an atom. Generally, the number of neutrons in the nucleus is equal to the number of protons (See also Chemical Elements). However, atoms of the same element can have the same number of protons but different numbers of neutrons, thus giving rise to varieties, or isotopes, of the same chemical element. The word isotope (Greek iso and topos) means “same place.” Different isotopes of the same element occupy the same place in the periodic table of the chemical elements and have very nearly identical chemical properties. Thus hydrogen, which has a mass number of 1, has an isotope, deuterium, which has one proton and one neutron in its nucleus and a mass number of 2. Deuterium accounts for 1.5 percent of naturally occurring hydrogen. Hydrogen and deuterium undergo the same chemical reactions, although not necessarily with equal ease.

The term atomic weight means the average weight (more correctly, the mass) of an atom of an element, taking into account the masses of all its isotopes and the percentage of their occurrence in nature. The atomic weight of an element was originally expressed relative to oxygen by assigning a value of 16.0000 to the mixture of oxygen isotopes found in nature. In 1961 this standard was changed by international agreement, and atomic weights are now determined relative to the weight of an atom of the most abundant isotope of carbon, carbon-12 (written ¹²C) which contains six neutrons and six protons. The weight of ¹²C is arbitrarily set equal to its mass number of 12.0000.

Isotopes are generally written as ¹²C or carbon-12, with the number denoting the total number of protons and neutrons in the atom. Four out of every five elements occur in nature as mixtures of isotopes (see Atomic Weight). For example, chlorine occurs in nature as a combination of two isotopes. Samples of chlorine contain 75.77 percent ³⁵Cl (with an atomic mass of 34.9689), and 24.23 percent ³⁷Cl (with an atomic mass of 36.9659). The average atomic mass of chlorine is (0.7577 × 34.9689) + (0.2423 × 36.9659) = 35.4527.

The molecular weight of a molecule is the sum of the atomic weights of the atoms making up that molecule (see Molecule). Thus the molecular weight of water (H₂O) is 2 × 1.00794 (for two hydrogen atoms) + 15.9994 (one oxygen atom), or 18.01528.

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2

The Electron Cloud

Most of the physical and chemical properties of atoms, and hence of all matter, are determined by the nature of the electron cloud enclosing the nucleus.

The nucleus of an atom, with its positive electric charge, attracts negatively charged electrons. This attraction is largely responsible for holding the atom together. The revolution of electrons about a nucleus is determined by the force with which they are attracted to the nucleus. The electrons move very rapidly, and determination of exactly where any particular one is at a given time is theoretically impossible (see Uncertainty Principle). If the atom were visible, the electrons might appear as a cloud, or fog, that is dense in some spots, thin in others. The shape of this cloud and the probability of finding an electron at any point in the cloud can be calculated from the equations of wave mechanics (see Quantum Theory). The solutions of these equations are called orbitals. Each orbital is associated with a definite energy, and each may be occupied by no more than two electrons. If an orbital contains two electrons, the electrons must have opposite spins, a property related to the angular momentum of the electrons. The electrons occupy the orbitals of lowest energy first, then the orbitals next in energy, and so on, building out until the atom is complete (see Atom).

The orbitals tend to form groups known as shells (so-called because they are analogous to the layers, or shells, around an onion). Each shell is associated with a different level of energy. Starting from the nucleus and counting outward, the shells, or principal energy levels, are numbered 1, 2, 3, … , n. The outer shells have more space than the inner ones and can accommodate more orbitals and therefore more electrons. The nth shell consists of 2n-1 orbitals, and each orbital can hold a maximum of 2 electrons. For example, the third shell contains five orbitals and holds a maximum of 10 electrons; the fourth shell contains seven orbitals and holds a maximum of 14 electrons. Among the known elements, only the first seven shells of an atom contain electrons, and only the first four shells are ever filled.

Each shell (designated as n) contains different types of orbitals, numbered from 0 to n-1. The first four types of orbitals are known by their letter designations as s, p, d, and f. There is one s-orbital in each shell, and this orbital contains the most firmly bound electrons of the shell. The s-orbital is followed by the p-orbitals (which always occur in groups of three), the d-orbitals (which always occur in groups of five), and finally the f-orbitals (which always occur in groups of seven). The s-orbitals are always spherically shaped around the nucleus; each p-orbital has two lobes resembling two balls touching; each d-orbital has four lobes; and each f-orbital has eight lobes. The p-, d-, and f-orbitals have a directional orientation in space, but the spherical s-orbitals do not. The three p-orbitals are oriented perpendicular to one another along the axis of an imaginary three-dimensional Cartesian (x, y, z) coordinate system. The three p-orbitals are designated p_x, p_y, and p_z, respectively. The d- and f-orbitals are similarly arranged about the nucleus at fixed angles to one another.

When elements are listed in order of increasing atomic number, an atom of one element contains one more electron than an atom of the preceding element (see Chemical Elements). The added electrons fill orbitals in order of the increasing energy of the orbitals. The first shell contains the 1s orbital; the second shell contains the 2s orbital and the 2p orbitals; the third shell contains the 3s orbital, the 3p orbitals, and the 3d orbitals; the fourth shell contains the 4s orbital, the 4p orbitals, the 4d orbitals, and the 4f orbitals.

After the two innermost shells, certain orbitals of outer shells have lower energies than the last orbitals of preceding shells. For this reason, some orbitals of the outer shells fill before the previous shells are complete. For example, the s-orbital of the fourth shell (4s) fills before the d-orbitals of the third shell (3d). Orbitals generally fill in this order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s.

In a notation frequently used to describe the electron configuration of an element, a superscript after the orbital letter gives the number of electrons in that orbital. Thus, 1s²2s²2p⁵ means that the atom has two electrons in the 1s orbital, two electrons in the 2s orbital, and five electrons in the 2p orbitals.

Neutral atoms with exactly eight electrons in the outer shell (meaning the s- and p-orbitals of the outer shell are filled) are exceptionally stable. These neutral atoms are atoms of the noble gases, which are so stable that getting them to chemically react with other elements is very difficult. The unusual stability of the noble-gas electron structures is of great importance in chemical bonding and reactivity. All other elements tend to combine with each other in such a way as to imitate this stable structure. The structure of helium is 1s²; neon adds another stable shell, 2s²2p⁶, to this; argon adds the orbitals 3s²3p⁶; krypton adds the orbitals 4s²3d¹⁰4p⁶; and xenon adds the orbitals 5s²4d¹⁰5p⁶ (the s-orbital fills before the d-orbital of the previous shell).

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Metals and Nonmetals

The structure of the atom, in particular the configuration of the electron cloud, is responsible for the obvious physical differences between metals and nonmetals. Metals have a characteristic luster, are opaque, can be hammered and drawn into various shapes, and conduct electricity. Nonmetal elements, on the other hand, are often gases, and, if solid, nonmetals are generally brittle, sometimes transparent, and do not conduct electricity.

The atoms of metals have outer shells that contain few electrons and are nowhere near filled (and therefore lack the stability of a noble gas). As a result, all metals tend to easily lose some of these outer electrons. This means, chemically, that metals tend to form positively charged ions, or positively charged atoms or molecules, when they enter into chemical combination. Physically, the fact that the outer shells of metal atoms are unfilled means that these “loose” electrons can flow and enable metals to conduct electricity; this fact also accounts for the mechanical properties of metals. Nonmetals, by contrast, have outer shells that are nearly filled (up to the stable grouping of eight electrons); in their chemical reactions they tend to add electrons to achieve the state of a stable noble gas. By adding electrons, nonmetals form negatively charged ions. They can also add electrons by sharing them with another atom and forming a covalent bond. The noble gases, with exactly eight electrons in their outer shells (two in the case of helium), are nonmetals.

There are many more metals than nonmetals, especially among elements of high atomic weight. A partial explanation of this fact is that the added electrons go mostly to fill the incomplete inner shells, leaving only two or three electrons in the outer shell.

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Chemical Bonds, Formulas, and Equations

Elements that do not have a noble-gas configuration (a stable configuration) try to attain such a configuration by entering into chemical reactions. Stable molecules are formed when atoms combine so as to have outer shells holding eight electrons.

If atoms are no more than a few electrons away from a stable configuration, they generally attain it by losing or gaining electrons to form electrically charged particles called ions. Positively charged ions (formed by a loss of electrons) are called cations, and negatively charged ions (formed by an electron gain) are called anions. Ions seldom have a charge greater than three, which means that atoms seldom gain or lose more than three electrons.

Table salt is composed of sodium and chlorine ions. The sodium atom loses its one outer electron to become a positively charged sodium ion. Its outer shell now contains eight electrons. The chlorine atom gains one electron in the outer shell, giving a total of eight electrons, to become a negatively charged chloride ion. The positive and negative ions attract each other and form a solid crystal.

The electrons in the outer shell of an element are called valence electrons. Valence electrons are those electrons that are available to form bonds with other atoms. Groups of elements with similar electron configurations (arrangements of electrons in their orbitals) behave in a similar way in chemical reactions, so these groups have similar chemical and physical properties. These groups of elements are called families (see Periodic Law). The periodic table shows how elements can be grouped into families. Elements with atoms that have one valence electron are in Group I; elements with two valence electrons are in Group II; elements with six are in Group VI; and elements with seven are in Group VII.

The periodic table helps chemists to remember the similarities and gradation of properties within element groups. The discovery of the periodic law and publication of this table in 1869 by Russian chemist Dmitry I. Mendeleyev was a major step in organizing information about the known elements and in predicting the properties of unknown ones.

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