Chemistry

I INTRODUCTION TO CHEMISTRY

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Chemistry

INTRODUCTION

Everyday Applications of Chemistry

Chemical reactions occur continuously in the atmosphere, in factories, in vehicles, in the environment, and in our bodies. In a chemical reaction, one or more kinds of matter is changed into a new kind—or several new kinds—of matter. A few common chemical reactions are shown here. Life as we know it could not exist without these processes: plants could not photosynthesize, cars could not move, pudding could not thicken, muscles could not burn energy, glue could not stick, and fire could not burn.

Chemistry, study of the composition, structure, properties, and interactions of matter. Chemistry arose from attempts by people to transform metals into gold beginning about ad 100, an effort that became known as alchemy (see Chemistry, History of). Modern chemistry was established in the late 18th century, as scientists began identifying and verifying through scientific experimentation the elemental processes and interactions that create the gases, liquids, and solids that compose our physical world. As the field of chemistry developed in the 19th and 20th centuries, chemists learned how to create new substances that have many important applications in our lives.

Chemists, scientists who study chemistry, are more interested in the materials of which an object is made than in its size, shape, or motion. Chemists ask questions such as what happens when iron rusts, why iron rusts but tin does not, what happens when food is digested, why a solution of salt conducts electricity but a solution of sugar does not, and why some chemical changes proceed rapidly while others are slow. Chemists have learned to duplicate and produce large quantities of many useful substances that occur in nature, and they have created substances whose properties are unique.

Much of chemistry can be described as taking substances apart and putting the parts together again in different ways. Using this approach, the chemical industry produces materials that are vital to the industrialized world. Resources such as coal, petroleum, ores, plants, the sea, and the air yield raw materials that are turned into metal alloys; detergents and dyes; paints, plastics, and polymers; medicines and artificial implants; perfumes and flavors; fertilizers, herbicides, and insecticides. Today, more synthetic detergent is used than soap; cotton and wool have been displaced from many uses by artificial fibers; and wood, metal, and glass are often replaced by plastics.

Chemistry is often called the central science, because its interests lie between those of physics (which focuses on single substances) and biology (which focuses on complicated life processes). A living organism is a complex chemical factory in which precisely regulated reactions occur between thousands of substances. Increased understanding of the chemical behavior of these substances has led to new ways to treat disease and has even made it possible to change the genetic makeup of an organism. For example, chemists have produced strains of food plants that are hardier than the parent strain.

Because the field of chemistry covers such a broad range of topics, chemists usually specialize. Thus, chemistry is divided into a number of branches, some of which are discussed at the end of this article. Nevertheless, the process of learning the properties of a substance and of taking it apart is fundamental to nearly all of chemistry.

The first step in investigating a complex material is to try to break it down into simpler substances. Sometimes this is easy. A mixture of brass and iron tacks, for instance, could be sorted with a magnet or even by hand. Getting the salt out of brine or seawater is a little harder, but the water can be evaporated, leaving the salt. Changes of this sort, which do not alter the fundamental nature of the components of the mixture but do modify their physical condition, are called physical changes. Grinding a rock, hammering a metal, or compressing a gas causes physical changes. Another example of physical change is the melting of ice, in which water changes from the solid to the liquid state.

Salt and water may not only be separated when in solution, but each may be broken down into other substances. This, however, involves a different kind of change—one that usually requires more energy than a physical change and that alters the fundamental nature of the material. This type of change is called a chemical change. By applying electrical energy, water can be broken down into two gases, hydrogen and oxygen. Hydrogen is a light gas that burns; oxygen is a gas that is necessary to sustain animal life. Salt can be broken down by melting it, then passing an electric current through it. This produces a pungent yellow-green gas called chlorine and a soft, silvery metal called sodium, which burns readily in air.

Some materials can be broken down simply by heating them. Other materials yield to attack by another substance; for example, iron oxide ore heated with coke yields metallic iron.

II ELEMENTS AND COMPOUNDS

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ELEMENTS AND COMPOUNDS

Periodic Table of Elements

The periodic table of elements groups elements in columns and rows by shared chemical properties. Elements appear in sequence according to their atomic number. Clicking on an element in the table provides basic information about the element, including its name, history, electron configuration, and atomic weight. Atomic weights in parentheses indicate the atomic weight of the most stable isotope.

More than 100 chemical elements—substances that cannot be decomposed or broken into more elementary substances by ordinary chemical means—are known to exist in the universe. However, several of these elements, such as the so-called transuranium elements, have not been found in nature and can only be produced artificially.

Russian chemist Dmitry Ivanovich Mendeleyev and German physicist Julius Lothar Meyer independently developed the periodic law of the chemical elements at about the same time in the late 19th century. Mendeleyev is generally credited with the findings, because he established the periodic law in 1869, and Meyer established this chemical law in 1870. Both discovered that arranging the elements in order of increasing atomic mass produced a table of chemical properties and reactivity patterns that were regularly repeated. This phenomenon—known as the periodic law—is most often represented in the periodic table of the elements (see Atom).

Elements

Hydrogen, oxygen, chlorine, sodium, and iron are examples of elements. Elements cannot be resolved into simpler substances by ordinary heat, light, electricity, or attack by other substances. To say that elements can never be broken down would not be accurate, but breaking them down takes millions of times more energy than can be applied by ordinary means. It requires either special equipment, such as a particle accelerator, or temperatures like those in the interior of the sun. An element can therefore be defined as a substance that cannot be broken down into simpler substances by ordinary means.

Ninety elements are known to occur in nature, and 22 more have been made artificially. Out of this limited number of elements, all the millions of known substances are made.

Abbreviating the names of the elements is often convenient. For each element, a symbol has been chosen that consists of one or two letters. The symbols are derived from the names of the elements; for example, H stands for hydrogen, He for helium, C for carbon, and so on. The abbreviations are not always derived from the English names, however. The symbol Fe for iron comes from the Latin ferrum, and W for tungsten comes from the German wolfram. These symbols are internationally recognized and are used even by people whose native languages do not use the Roman alphabet, such as Russian and Japanese.

Compounds

Water Molecule

Water is an example of a compound. A water molecule consists of an oxygen atom and two hydrogen atoms.

Salt, water, iron rust, and rubber are examples of compounds. A compound is made up of elements, but it looks and behaves quite differently, as a rule, from any of its component elements. Iron rust, for example, does not look and feel like its components: oxygen gas and iron metal. Some synthetic fabrics, with fibers made from coal, air, and water, do not feel at all like any of the components that make them up. This individuality of properties, as well as other qualities, distinguishes a compound from a simple mixture of the elements it contains. Another important characteristic of a compound is that the weight of each element in the compound always has a fixed, definite ratio to the weight of the other elements in the compound. For example, water always breaks down into 2.016 parts of hydrogen by weight to 16.000 parts of oxygen by weight, which is a ratio of about 1 to 8, regardless of whether the water came from the Mississippi River or the ice of Antarctica. In other words, a compound has a definite, invariable composition, always containing the same elements in the same proportions by weight; this is the law of definite proportions.

Many elements combine in more than one ratio, giving different compounds. In addition to forming water, hydrogen and oxygen also form hydrogen peroxide. Hydrogen peroxide has 2.016 parts of hydrogen to 32 parts of oxygen; that is, 1.008 parts of hydrogen to 16 parts of oxygen. Water, as stated above, has 2.016 parts of hydrogen to 16 parts of oxygen. The figure 2.016 is twice 1.008. This example illustrates the law of multiple proportions: When two elements combine to form more than one compound, the element whose mass varies combines with a fixed mass of the second element weights in a simple whole-number ratio such as 2:1, 3:1, or 3:2.

Atoms and Molecules

Atoms Made Visible

Individual atoms of the element silicon can be seen in this image obtained through the use of a scanning transmission electron microscope. The atoms in each pair are less than a ten-millionth of a millimeter (less than a hundred-millionth of an inch) apart.

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.

The Structure of Atoms

Models of the Atom

Experimental data has been the impetus behind the creation and dismissal of physical models of the atom. Rutherford's model, in which electrons move around a tightly packed, positively charged nucleus, successfully explained the results of scattering experiments, but was unable to explain discrete atomic emission—that is, why atoms emit only certain wavelengths of light. Bohr began with Rutherford’s model, but then postulated further that electrons can only move in certain quantized orbits; this model was able to explain certain qualities of discrete emission for hydrogen, but failed completely for other elements. Schrödinger’s model, in which electrons are described not by the paths they take but by the regions where they are most likely to be found, can explain certain qualities of emission spectra for all elements; however, further refinements of the model, made throughout the 20th century, have been needed to explain all observable spectral phenomenon.

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.

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.

Elementary Particle Tracks

These tracks were formed by elementary particles in a bubble chamber at the CERN facility located outside of Geneva, Switzerland. By examining these tracks, physicists can determine certain properties of particles that traveled through the bubble chamber. For example, a particle's charge can be determined by noting the type of path the particle followed. The bubble chamber is placed within a magnetic field, which causes a positively charged particle's track to curve in one direction, and a negatively charged particle's track to curve the opposite way; neutral particles, unaffected by the magnetic field, move in a straight line.

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.

The Electron Cloud

Electron Density and Orbital Shapes

Atomic orbitals are mathematical descriptions of where the electrons in an atom (or molecule) are most likely to be found. These descriptions are obtained by solving an equation known as the Schrödinger equation, which expresses our knowledge of the atomic world. As the angular momentum and energy of an electron increases, it tends to reside in differently shaped orbitals. The orbitals corresponding to the three lowest energy states are s, p, and d, respectively. The illustration shows the spatial distribution of electrons within these orbitals. The fundamental nature of electrons prevents more than two from ever being in the same orbital. The overall distribution of electrons in an atom is the sum of many such pictures. This description has been confirmed by many experiments in chemistry and physics, including an actual picture of a p-orbital made by a Scanning Tunneling Microscope.

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).

Metals and Nonmetals

Metallic Bonding

Silver, a typical metal, consists of a regular array of silver atoms that have each lost an electron to form a silver ion. The negativly charged electrons distribute themselves throughout the entire piece of metal and form nondirectional bonds between the positive silver ions. This arrangement, known as metallic bonding, accounts for the characteristic properties of metals: they are good electrical conductors because the electrons are free to move from one place to another, and they are malleable (as shown here) because the positive ions are held together by nondirectional forces. A force applied to a malleable substance shifts the positions of the atoms without breaking the bonds that hold them together.

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.

III CHEMICAL BONDS, FORMULAS, AND EQUATIONS

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III

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.

Ionic Bonds

Ionic Bonding: Salt

The bond (left) between the atoms in ordinary table salt (sodium chloride) is a typical ionic bond. In forming the bond, sodium becomes a cation (a positively charged ion) by “giving up” its valence electron to chlorine, which then becomes an anion (a negatively charged ion). This electron exchange is reflected in the size difference between the atoms before and after bonding. Attracted by electrostatic forces (right), the ions arrange themselves in a crystalline structure in which each is strongly attracted to a set of oppositely charged “nearest neighbors” and, to a lesser extent, all the other oppositely charged ions throughout the entire crystal.

Oppositely charged ions have a strong mutual electrostatic attraction when brought together, but, if brought too close, the electron clouds repel each other. Thus, a pair of mutually attracted ions will maintain a certain distance from each other. This distance is called the bond length, and the electrostatic attraction of the ions constitutes an ionic (or electrovalent) bond. Ionic bonds are very common and are exemplified by table salt, in which a sodium ion attracts a chloride ion to form Na⁺Cl^-or, as usually written, NaCl. Calcium ions (Ca²⁺) and chloride ions (Cl^-) combine in a one-to-two ratio to form calcium chloride, CaCl₂. The total charge on each combination of ions, NaCl and CaCl₂, is neutral, or zero.

Covalent Bonds

Predicting a Molecule's Shape

Molecules with a central atom form predictable shapes that depend on the number of electron pairs residing in the outer shell of the central atom. The repulsive force between these outer electron pairs causes them to take positions as far from each other as possible. The position of the electron pairs in the outer shell then determines the angles at which the central atom bonds with the surrounding atoms in the molecule.

Another common type of bond, the covalent bond, results when two atoms share one or more pairs of electrons in an attempt to fill their outer shells and become more energetically stable. The atoms are held together by the mutual electrostatic attraction between the protons in their nuclei and these electrons. The bonded atoms form a stable unit called a molecule.

For example, because a chlorine atom is one electron short of completing its outer shell (and attaining a noble-gas configuration), two chlorine atoms combine to form a chlorine molecule by sharing two electrons. The atoms thereby complete each other’s outer shell: Cl + Cl → Cl₂. Electron sharing distinguishes a covalent bond from an ionic bond (in which the electrostatic attraction of the ions results in bonding). In an ionic compound there are no molecules—only charged particles composing an extensive three-dimensional array.

Covalent bonds tend to form when the bonded atoms have nearly the same attraction for electrons; ionic bonds form when the electron-attracting power of the atoms differs markedly. If the valence electrons are represented by dots, the difference between bond types becomes more obvious:

Sometimes two or three pairs of electrons will be shared between two atoms, producing double or triple bonds:

Compounds such as LiF, BeO, and BeF₂ are ionic in character, whereas molecules formed between neighbors or near neighbors in the periodic table are more often covalent (such as CO₂, CF₄, NO₂, N₂, O₂, and F₂). Some metals, however, form both ionic and covalent bonds. A convenient rule for remembering whether the bond between two elements is likely to be ionic or covalent is that if one element appears on the left side of the periodic table and the other on the right, the bond is ionic.

Polarity and Electronegativity

Alignment of Polar Molecules

Polar molecules form when a strongly electronegative atom (an atom that attracts electrons) such as chlorine bonds with a less electronegative atom such as hydrogen. Because the electronegative atom pulls the electron(s) away from the other atom, the molecule formed by the atoms is polar, meaning one end is negatively charged and the other end is positively charged. The molecules align themselves because the negative end of each molecule is attracted to the positive end of other molecules, and vice versa. Polar molecules containing hydrogen are so strongly polar (and therefore attracted to each other), they form a type of bond called a hydrogen bond.

Between a pure covalent bond (as in Cl₂) and a strongly ionic bond (as in LiF), there is a gradual shift from one bond type to the other that is related to the differences in electron attraction between the bonded atoms. The ability of an atom to attract electrons in a bond is called its electronegativity—the stronger an atom pulls electrons, the higher its electronegativity. Elements on the right side of the periodic table (except for the noble gases) are the most electronegative, because they need fewer electrons than elements on the left side to fill their outer shell and attain the stability of a noble gas. For example, fluorine is much more electronegative than potassium.

Bonds between atoms of widely different electronegativity are highly ionic, because strongly electronegative atoms (such as oxygen) tend to pull electrons away from less electronegative atoms. Bonds between atoms of more similar electronegativity take on a more covalent character and eventually become completely covalent. Since the chlorine atoms in Cl₂ have identical electronegativity, this bond is purely covalent, and the atoms exactly share their electrons. The bond between hydrogen and oxygen atoms in water is mostly covalent and has some ionic character: the oxygen is more electronegative than the hydrogen and has a greater share of the bonding electrons. Such bonds are said to be polar, because the shared electrons are held more tightly by the oxygen atom (giving the oxygen atom a slightly negative charge) and pulled away from the hydrogen atoms (giving the hydrogen atoms a slightly positive charge). Molecules with more ionic bonds are more polar than molecules with less ionic, more covalent bonds.

Oxidation State

Oxidation: A Chemical Reaction

Oxidation, in its original sense, refers to the combination of oxygen with another substance to produce a compound called an oxide. Iron, in the presence of water, combines with atmospheric oxygen to form a hydrated iron oxide, commonly called rust.

The polarity of the bond between hydrogen and oxygen, the HO bond, illustrates the concept of oxidation, which is a useful way to visualize how electrons are shared between atoms and to explain certain chemical reactions. Since the oxygen atom has the greater attraction for the shared pair of electrons, this atom effectively gains one electron at the expense of the hydrogen atom. The number of electrons an atom gains or loses in a bond determines the atom’s oxidation state, or its effective charge. Except in special cases, an oxygen atom will form two covalent bonds (to complete its outer shell), as in H₂O. An oxygen atom normally accepts two electrons when bonding, giving it an oxidation state of -2. Since the hydrogen atom is assumed to have lost one electron, this atom has an oxidation state of +1. In a neutral compound, the sum of the oxidation states of all atoms is zero. In a charged compound, the sum of the oxidation states of all atoms equals the net charge of the compound.

These rules allow calculation of the oxidation state of an atom in a particular compound. In SO₂, each oxygen atom takes two electrons (-2 each; total of -4) from the sulfur atom, and, since the total charge must be zero, the oxidation state of S is +4 (meaning sulfur loses four electrons to oxygen). In SO₄^-2, each oxygen atom has an oxidation number of –2, and since there are four oxygen atoms in SO₄^-2, the total oxidation number for the oxygen atoms is (4 × –2) = -8. Because the SO₄^-2 ion has an overall charge of –2, sulfur must have an oxidation number of +6 (calculated from: –8 + x = -2). In the compound H₂S, sulfur is more electronegative than hydrogen, so the sulfur causes each hydrogen atom to effectively lose an electron, giving S an oxidation number of –2.

Formulas

To show the atomic constituents of a compound, two or more chemical symbols are put together with subscripts, giving formulas. Since molecules are made up of atoms in simple whole-number ratios, molecular composition can be expressed by means of element symbols, with subscript numbers used to indicate the relative number of atoms of each element. Thus water, with molecules composed of two atoms of hydrogen and one atom of oxygen, is written as H₂O; hydrogen peroxide, with two atoms of each of these elements, is H₂O₂. The formula CO stands for carbon monoxide (one atom of carbon and one atom of oxygen), and CO₂ stands for carbon dioxide (one atom of carbon and two atoms of oxygen). The formula FeSO₄ (one atom each of iron and sulfur and four oxygen atoms) stands for ferrous sulfate. The formula Fe₂(SO₄)₃ represents another sulfate of iron, ferric sulfate, which contains three sulfate ions (SO₄) and two ferric ions (Fe).

When groups of atoms combine as a subunit, the subunit is usually treated as a single symbol. For example, the sulfate group, SO₄^2-, is treated as a single symbol (included in molecules such as Al₂(SO₄)₃, aluminum sulfate), as is the nitrate group, NO_3- (included in molecules such as NH₄NO₃, ammonium nitrate).

Ions are represented by adding the electric charge as a superscript. Thus Na⁺ represents the sodium ion (loss of one electron, or one less electron than Na), Cl^- the chloride ion (addition of one electron, or one more electron than Cl), SO₄^2- the sulfate ion (addition of two electrons), Fe²⁺ the ferrous ion (loss of two electrons), and Fe³⁺ the ferric ion (loss of three electrons).

Structural Formulas

Structural formulas show how atoms are connected together, so that different compounds containing equal numbers of the same elements—known as isomers—can be distinguished from each other. Examples of isomers are ethyl alcohol and dimethyl ether. Both share the simple chemical formula C₂H₆O. Although ethyl alcohol and dimethyl ether both contain equal numbers of atoms of the same elements, the atoms within the respective molecules are arranged differently, producing two different compounds. These differences can be shown by structural formulas, which show how atoms within a molecule are arranged by using lines to represent chemical bonds:

Simpler representations, such as CH₃CH₂OH (ethyl alcohol) and CH₃-OCH₃(dimethyl ether) can also be used to distinguish these isomers. In ring, or circular, compounds of carbon, chemical symbols are sometimes omitted, and only the outline of the ring bonds is shown. Below (left) is a structural formula for benzene and (right) is an abbreviation:

Equations

Chemical reactions can be expressed through equations that resemble mathematical equations. The reactants (the substances that are combined to react with one another) appear on the left side of the equation, and the products (substances produced by the reaction) are written on the right side of the equation. The reactants and products are typically connected by an arrow (→) or various types of double arrows. The single arrow shows that a reaction only proceeds in the direction indicated, while the double arrow indicates that a reaction can proceed in either direction (that products are also reacting with each other to reform reactants). Thus, the equation 2Cl ⇄ Cl₂ indicates that two chlorine atoms react to form a molecule of chlorine and that a reaction can also take place in the reverse direction. This reaction, like all reactions, is affected by conditions such as temperature. In the case of 2Cl ⇄ Cl₂, the reaction goes to the right at room temperature (25° C; 77° F) and to the left only at much higher temperatures (about 500° C; about 932° F). Sometimes the condition under which a reaction takes place is written over or below the arrow(s).

The Greek letter delta, Δ, is sometimes used to indicate that the reactants must be heated for a reaction to take place. An equation may also show the physical state of the reactants and products by using the letters g, l, s, and aq to stand for gas, liquid, solid, and aqueous solution, respectively. Thus, the equation

indicates that when a solution of calcium bicarbonate is heated, a precipitate of calcium carbonate (CaCO₃) forms, carbon dioxide is evolved as a gas, and water is liberated (as a gas or a liquid, depending on the temperature). On the other hand, the reaction can go from right to left if gaseous carbon dioxide is reacted with solid calcium carbonate in the presence of water and without heat.

Balanced Equations

The reaction between sulfuric acid (H₂SO₄) and sodium hydroxide (NaOH) to form sodium sulfate (Na₂SO₄) and water can be written: NaOH + H₂SO₄ → H₂O + Na₂SO₄. This is an incomplete equation, since the same number of atoms does not appear on both sides of the reaction (for example, 1 sodium atom appears in the reactants, but 2 sodium atoms appear in the products). Such a reaction could not actually occur. Correcting this fault is called balancing the equation. Placing a 2 in front of the NaOH balances the number of sodium atoms, as well as the hydrogen and oxygen atoms. The balanced equation is: 2NaOH_(aq) + H₂SO₄ → 2H₂O_(l) + Na₂SO_4(aq); this equation implies that two molecules of NaOH are necessary to react with each molecule of H₂SO₄. An equation must be balanced before a chemist can make calculations based on it.

Ionic Equations

In solution, many substances dissociate, or break apart, into ions, which may then enter into chemical reactions. For example, the reactants in the above equation, sodium hydroxide and sulfuric acid, dissolve and react in water. This reaction can be written in terms of the ions that the reactants form: 2Na⁺ + 2OH^- + 2H⁺ + SO₄^2- → 2H₂O + 2Na⁺ + SO₄^2-.

The sodium ions (Na⁺) and the sulfate ion (SO₄^2-) are not changed and appear on both sides of the equation. Only the hydroxide (OH^-) ions and the hydrogen (H⁺) ions actually react, and the net reaction is therefore: 2OH^- + 2H⁺ → 2H₂O or more simply: OH^- + H⁺ → H₂O.

Acids, Bases, and Salts

Acids are compounds (or ions) that react with water to produce hydrogen ions (H⁺) (see Acids and Bases). Hydrogen ions account for the characteristic properties of strong acids, such as a sour taste and the ability to react with bases. Bases are compounds that yield the hydroxide ion (OH^-) in water solutions. Salts are ionic compounds that are generally formed by the reaction of an acid and a base:

Weight Relations

The atomic weight of an element, which is given on the periodic table, is the average mass of atoms of an element (since most elements consist of different isotopes, the differing weights of these isotopes are taken into account). For example, in a naturally occurring sample of nitrogen (N), the average weight of the nitrogen atoms is 14.00674 atomic mass units. The proportions by weight of elements in a compound can be determined from the chemical formula of the compound and the atomic weights of the component elements. For example, the proportions by weight of elements in Fe₂O₃, ferric oxide, which forms when iron rusts, can be determined by finding the atomic weights of iron and oxygen on the periodic table and adding them:

The fraction of iron in pure ferric oxide is 111.6940/159.6922 = 0.6994, or 69.94 percent. The quantity 159.6922 is called the formula weight or molecular weight of the compound.

Extending these weight concepts to chemical equations makes possible the calculation of how much of each reactant will be needed to react (without leaving one in excess) and how much of the various products are formed. Thus, in the oxidation of iron: 4Fe_(s) + 3O_2(g)→ 2Fe₂O_3(s), 223.3880 grams (4 x 55.8470) of iron reacts with 95.9964 grams (6 × 15.9994) of oxygen to produce 319.3844 grams (2 × 159.6922; or 223.3880 + 95.9964) of ferric oxide. In this way, the amount of ferric oxide formed from a given amount of iron can be calculated.

Gas Volumes in Chemical Reactions

At the same temperature and pressure, equal volumes of different gases contain identical numbers of molecules (see Avogadro’s Law). The volumes and number of molecules in a chemical reaction all relate to one another in identical ratios (see The Gas Laws section below).

When methane (CH₄), the main constituent of natural gas, is burned, the balanced equation for the reaction is CH_4(gas) + 2O_2(gas)→ CO_2(gas) + 2H₂O_(gas). This equation shows that one volume of methane needs two volumes of oxygen to burn it, and that one volume of carbon dioxide and two volumes of water vapor are produced. The volumes are compared at the same temperature and pressure.

IV PHYSICAL PROPERTIES

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PHYSICAL PROPERTIES

Physical properties of a material are those properties that do not depend on the chemical behavior of the material. Physical properties include the state of a material (gas, liquid, or solid), melting point, boiling point, crystal structure, and electrical conductivity.

State

The state of a material is determined by the attraction between its atoms or molecules and by the temperature of the material. In the solid state, the attraction between the atoms or molecules is so strong that it holds them rigidly in place. The energy of vibration of the molecules of a material increases with a rise in temperature. As the temperature rises, the molecules eventually acquire enough energy to break away from their fixed positions, and the solid either melts or transforms directly into gas (a process called sublimation). The material melts if the molecular attraction remains great enough to hold the molecules together, and the material sublimes to a gas (in which the molecules are free to move randomly) if the attraction is too small.

Melting Point

The melting point (or freezing point) of a substance is the temperature at which the solid form of the substance changes to a liquid (or from liquid to solid). The melting point of water is 0° on the Celsius (centigrade) temperature scale and 32° on the Fahrenheit scale (see Freezing Point).

Boiling Point

The boiling point of a substance is the temperature at which the liquid form of the substance changes to a gas. The boiling point is sensitive to changes in pressure, because the molecules of a substance will tend to stay in the liquid state if they are under enough pressure. A heated liquid must overcome the atmospheric pressure in order to turn into a gas (if the atmospheric pressure exceeds the vapor pressure of the boiling liquid, the liquid will be unable to turn into vapor). For this reason, water boils at lower temperatures on high mountains (where atmospheric pressure is lower) than at sea level (where atmospheric pressure is higher). The boiling point of water at a pressure of one atmosphere, or 760 mm of mercury (a standard pressure approximating sea-level pressure), is 100° on the Celsius scale and 212° on the Fahrenheit scale.

Crystal Structure

Solids may be either amorphous or crystalline in their molecular structure. In amorphous solids, the molecules are arranged haphazardly. Glass is an example of an amorphous material. Like other amorphous materials, glass does not melt at a particular temperature, because the long, randomly intertwined glass molecules cannot easily become disentangled. As a result, glass softens bit by bit as the temperature is raised, eventually becoming liquid. Crystalline materials, on the other hand, have a definite orderly array of atoms, ions, or molecules, as would a pyramid of oranges or cannonballs. The orderly arrangement of particles in a crystal is called a crystal lattice. Sand, salt, sugar, diamond, and graphite are examples of common crystalline materials. Each crystalline material has a unique melting temperature (provided the material is not chemically changed by the heat before it melts, as happens with sugar).

In an ionic crystal, the strength of mutual attraction of the ions in the crystal is reflected in the high melting point of the crystal. Table salt (or sodium chloride, NaCl), for example, melts at 800° C (1472° F). Many ionic compounds, such as sodium chloride, form crystal arrays in which each positive ion is surrounded by negative ions, and each negative ion is surrounded by positive ions. The closely packed arrangement of ions in a crystalline solid, as well as the strong attraction between oppositely charged ions, accounts for the relatively hard, brittle nature of many ionic crystal solids.

Covalent crystal structures are networks of bonded atoms with the atoms occurring at the lattice points of the crystal. In the crystal lattice of diamond, each carbon atom is covalently bonded to four neighboring carbon atoms, forming a giant three-dimensional network. This three-dimensional network that composes diamond forms the hardest-known naturally occurring substance. Most covalent crystal structures are very hard and have very high melting points, because covalent bonds throughout the crystal make it essentially one giant molecule. Other examples of covalent crystals include silicon carbide (SiC), sand, and quartz (SiO₂).

Metallic crystals have unique properties because of the relationship between the positive ions and the electrons of the metal. One of the simpler and more widely used models of metallic crystals shows positive ions arranged at the points of the crystal, with electrons moving freely (as a so-called sea of electrons) among these positive ions. Because electrons in metals do not belong to any single positive ion and can move freely (carrying their electric charge with them), metals are excellent conductors of electricity. If an electric potential is applied to the metal, the electrons will move readily toward the positive electrode, creating an electric current (a stream of electrons). The freely moving electrons also make metals good transmitters of heat (metals are cold to the touch because electrons move heat away from skin).

Molecular crystals are compounds in which the molecules are held together in a crystal lattice by weak intermolecular attractive forces (for more information, see the Solutions and Solubility section of this article). These crystals do not form a complete network. Because of the weak attractions between the molecules, molecular crystals have low melting temperatures (typically well below room temperature) and are relatively soft. Most molecular organic (carbon-containing) and inorganic compounds form molecular crystals. Examples include ice (solid H₂O), solid sulfur dioxide (SO₂), and solid carbon dioxide (CO₂).

V SOLUTIONS AND SOLUBILITY

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SOLUTIONS AND SOLUBILITY

A solution is a homogeneous mixture of two or more substances. Solutions form because even electrically neutral molecules have weak attractions for one another. Much of this attraction comes from the polarity, or slight unevenness of the electrical charge distribution within the molecules—a local region of slight negative charge in one molecule attracting a region of slight positive charge in another. These weak opposite charges hold molecules together in a liquid and also account for the ability of a liquid to dissolve other substances.

When a substance (called the solute) dissolves in a liquid (called the solvent), the molecules of the solvent must force their way between molecules of the solute. This occurs, for example, when water dissolves crystals of sugar. The solvent can only dissolve the solute if the solvent and solute molecules have similar attractive forces, which leads to the rule that like dissolves like. (The word like refers to similar characteristics of polarity.)

In general, polar molecules will be strongly attracted to one another. For example, water and alcohol mix readily. This occurs because the electronegative (electron-attracting) oxygen atom in water has a slight negative charge, giving both hydrogen atoms a slight positive charge. Similarly, the electronegative –OH (hydroxyl) group of alcohol has a slight negative charge, giving the other end of the alcohol molecule a slight positive charge. Thus, when water and alcohol are added together, the oppositely charged regions of the two types of molecules attract each other, allowing the liquids to mix. Other polar molecules that will dissolve in water include sugar, starch, and vitamin C. Ionic compounds that will dissolve in water include baking soda (NaHCO₃) and table salt (NaCl). The following reaction shows table salt dissolving into its constituent ions in water: NaCl (in H₂O_(l))→ Na⁺+ Cl^-_.

Similarly, nonpolar liquids will mix with each other, because such molecules are able to pry apart other nonpolar molecules. For example, gasoline and carbon tetrachloride, which are both composed of nonpolar molecules, mix easily and are good solvents for similarly nonpolar molecules, such as fats, greases, and paraffins.

On the other hand, a liquid composed of polar molecules will not readily mix with a liquid composed of nonpolar molecules, because the nonpolar molecules are not able to pry the polar molecules apart. For instance, water will not readily mix with gasoline or benzene (both of which consist of nonpolar molecules), because the polar water molecules are too strongly held together to allow entry of the nonpolar hydrocarbons.

While some liquids, such as water and alcohol, can dissolve in each other in any proportion, other compounds cannot. For example, salt added to water will dissolve until a threshold is reached, after which new salt added will no longer dissolve. This solution of salt water is then called saturated. A compound’s solubility in a given solvent is measured as the maximum amount of the compound that a solution can dissolve.

Effect of Temperature and Pressure

Raising the temperature usually increases the solubility of liquids and solids. The increase in temperature increases the energy of motion of the molecules (the kinetic energy) and partially overcomes the lack of attraction between polar and nonpolar molecules. Pressure has little effect on the solubility of liquids and solids, because the volumes of these materials change only slightly when they are dissolved.

Pressure has more effect on the solubility of gases in liquid solvents. A gas is more soluble as the pressure increases, because the gas atoms or molecules are crowded together, forcing more of the gas particles into contact with the liquid. Gases, however, become less soluble as the temperature increases, because as the gas molecules gain energy of motion, they are more easily able to escape the solution.

Electrolytes

Solutions of ions conduct an electric current, in much the same manner as a wire does. Ions can move about in a solution and carry a charge, just as electrons moving along a wire conduct a current. Substances that can carry a charge through solution in this way are called electrolytes; those that cannot are called nonelectrolytes.

Concentration

The amount of dissolved material in a solution is called the concentration and can be expressed in units, such as grams per liter or ounces per gallon. Chemists sometimes use percent to indicate concentration, and by convention indicate whether the percent is by weight or volume. A 10 percent solution of alcohol in water would normally be thought of as 10 volumes of alcohol in 90 volumes of water. A 10 percent solution of sodium chloride is thought of as 10 weight units of salt in 90 weight units of water. Scientists often use parts per million (or billion) when the amount of solute is very small. For chemical purposes, expressing concentration in terms of the number of molecules (or ions) in solution is often preferable.

Mole

The mole is one of the seven fundamental units in the International System of Units (SI). The mole is the unit used for measuring the amount of a substance and is defined as the amount of a substance containing the same number of atoms, molecules, or ions as the number of atoms in 12 grams of ¹²C. Because there are 6.022 × 10²³ atoms of carbon in 12 grams of ¹²C, this number (6.022 × 10²³), known as Avogadro's number, is the amount of matter containing 6.022 × 10²³ atoms, molecules, or ions. (see Avogadro’s Law).

The mole concept provides a means of calculating how many atoms, ions, or molecules are in a sample by weighing the substance. From the definition of atomic weight, the amount of any element that has a mass (in grams) equal to its atomic weight (available on the periodic table) will contain 6.022 × 10²³ atoms. Thus, 4.0026 grams of helium, 32.0064 grams of sulfur, and 200.59 grams of mercury each contain 6.022 x 10²³ atoms.

Similarly, a mole of a molecular substance (6.022 × 10²³ molecules) is the amount of the substance whose mass (in grams) is equal to its molecular weight. Molecular weight is derived by summing the atomic weights of the atoms composing a molecule. For example, 70.906 grams (2 × 35.453) of Cl₂ contains 6.022 × 10²³molecules (one mole) of Cl₂.

Chemists use the same principle to measure the number of ions in a compound. For example, one mole of sodium ions (Na⁺) has a mass of 22.9898 grams (atomic weight of Na is 22.9898). One mole of NaCl has a mass of 58.443 grams (22.9898 + 35.453).

Molarity

Molarity is the concentration of a substance in solution and is expressed as moles of solute per liter of the solution. Thus, a 0.1 molar (abbreviated 0.1 M) solution of sodium chloride contains 5.8443 (58.443 × 0.1) grams of NaCl per liter of solution.

Molality

Molality, a term less frequently used than molarity, is the number of moles of solute in 1000 grams of solvent. Thus, a 0.1 molal solution of sodium chloride in water has 5.8443 grams of NaCl in 1000 grams of H₂O.

Normality

Normality is the number of equivalents per liter of solution. For acid-base-salt systems, an equivalent is the amount of the substance that will gain or lose one mole of H⁺ ions. For instance, one mole of sulfuric acid (H₂SO₄), which has a mass of 98.0795 grams, produces two moles of H⁺, or two equivalents. Therefore, a one molar