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
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 Dmitri 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).
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.
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.
C 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.
C1 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.
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.
C2 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 px, py, and pz, 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).
D 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.
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 Dmitri I. Mendeleyev was a major step in organizing information about the known elements and in predicting the properties of unknown ones.
A Ionic Bonds
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.
B Covalent Bonds
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.
C Polarity and Electronegativity
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.
D Oxidation State
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.
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, SO42-, is treated as a single symbol (included in molecules such as Al2(SO4)3, aluminum sulfate), as is the nitrate group, NO3- (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), SO42- the sulfate ion (addition of two electrons), Fe2+ the ferrous ion (loss of two electrons), and Fe3+ the ferric ion (loss of three electrons).
F 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:
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.
G1 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) + Na2SO4(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.
G2 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+ + SO42- → 2H2O + 2Na+ + SO42-.
The sodium ions (Na+) and the sulfate ion (SO42-) 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.
H 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:
I 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) + 3O2(g) → 2Fe2O3(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.
J 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 CH4(gas) + 2O2(gas) → CO2(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
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.
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.
B 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).
C 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.
D 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
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.
A 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.
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.
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.
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 × 1023 molecules (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 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, 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 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 solution of sulfuric acid is a two normal (2 N) solution. A 0.1 N solution (containing 0.1 moles of H+) of sulfuric acid contains 4.90397 grams of H2SO4 per liter of solution ([98.0795/2] × 0.1).
VI THE GAS LAWS
Because all gases behave slightly differently when exposed to a range of pressures, chemists describe gas behavior using a hypothetical gas known as an ideal gas. An ideal gas is a theoretical gas that adheres completely to the following gas laws: (1) one mole of a gas at standard conditions—760 mm pressure and 0° C (32° F)—occupies a volume of 22.4 liters; (2) if the temperature and amount (number of moles) of a gas are held constant, the volume of the gas varies inversely with changes in pressure; (3) if the pressure and amount of a gas are held constant, the volume of the gas will change in direct proportion to a change in the absolute temperature (measured in degrees Kelvin; 0 K is equal to –273.15° C or –459.67° F); and (4) if temperature and pressure are held constant, the volume changes in direct proportion to a change in the number of moles of gas.
These laws can be expressed in the formula PV = nRT, where P is the pressure of the gas, V is the volume, T is its absolute temperature, n is the number of moles of the gas, and R is a constant. When volume is measured in liters and pressure is measured in atmospheres, the constant R is equal to 0.0821 liter atmospheres per degree per mole.
Gases that exactly obey the gas laws are called ideal gases. The noble gases, and others with a very low boiling point (for example, hydrogen, oxygen, and nitrogen), come closest to being ideal gases. Gases with relatively high boiling points, such as carbon dioxide gas, obey these laws only approximately.
The ideal gas laws make the assumptions that gas molecules (or atoms) do not attract each other and that gas particles occupy no volume. The first assumption is reasonable, because the kinetic energy (energy of motion) possessed by gas molecules vastly exceeds their mutual attraction, preventing the particles from condensing to liquid. The second assumption is approximately correct, since a gas occupies a much larger volume (most of which is empty space) than the same number of molecules (or atoms) would occupy as a liquid.
If a gas is enclosed in a container, the pressure exerted by the gas is the energy of the molecules bombarding the container walls. At the same temperature, all molecules and atoms (regardless of type) in a gas possess the same kinetic energy. As a result, a mole of hydrogen gas (H2) will exert the same pressure against the walls of a container as a mole of propane (C3H8) gas, or any other gas, at the same temperature and volume. Similarly, at the same temperature and pressure, a mole of hydrogen gas will occupy the same volume as a mole of propane gas, or any other gas.
VII CHEMICAL REACTIONS
Chemical reactions occur when atoms or molecules combine to form products with new properties. Plants and animals are literal chemical factories driven by chemical reactions. Our modern quality of life depends on producing and transforming substances through chemical reactions.
The main force in bringing about chemical reactions is the drive in any system to attain the lowest energy state, that is, to attain a more stable electron arrangement in the products than in the reactants. A second force is the tendency of a system to become as disorganized as possible, that is, to achieve the highest entropy (see Thermodynamics; Inorganic Chemistry).
When there is a large energy difference (a large stability difference) between the reactants and the products, a reaction tends to consume the reacting materials completely to generate more energetically stable products. For example, if hydrogen and oxygen are mixed in the proper proportions and the mixture is ignited, these gases react completely to form water, with none of the reactants remaining: 2H2 + O2 → 2H2O.
Bond strength is usually measured by bond energy, which is the energy necessary to break apart certain bonds in molecules. Reactions tend to lead toward greater bond strengths (or bond energies). For example, in the above reaction, the bond strength in the two OH bonds in H2O is greater than the sum of the strengths of the HH bond in H2 and the OO bond in O2. Because the O-H bond strengths in H2O exceed the combined strengths of the H-H bond and the O-O bond, water is more energetically stable, which explains why this reaction occurs.
Often, the additional stability gained by a chemical reaction is not large. In these cases, the reacting materials are not completely converted to the products, and appreciable amounts of reactants and products exist together. In such cases, the system is said to be in equilibrium (indicated by ⇄). For example, acetic acid (CH3COOH), which is very soluble in water, reacts with water to form ions: H2O + CH3COOH ⇄ CH3COO- + H3O+. This reaction stops short of complete ionization (complete dissociation of CH3COOH into CH3COO- and H3O+). After about 1 percent of the acetic acid molecules have dissociated into ions, the rate of recombination of ions to molecules (CH3COO- + H3O+ → CH3COOH + H20) equals the rate of ionization. As a result, the system reaches a state of dynamic equilibrium, and there is no further net change in the concentrations of the products or reactants.
Adding or removing any of the products or reactants involved in a reaction can change the position of equilibrium. Changing the concentration of any of the reactants or products when they are in equilibrium with each other will shift the position of the equilibrium to minimize this change in concentration. This principle is known as Le Châtelier's principle after its formulator, French chemist Henri Louis Le Châtelier. In the above reaction between water and acetic acid, adding more water to the system causes more CH3COO- and H3O+ to form as the system seeks to minimize the higher H2O concentration by producing more ions (CH3COO- + H3O+). As a result, adding water shifts the equilibrium to the right. Conversely, adding more H3O+ to the system will cause more H2O and CH3COOH to form as the system seeks to minimize the higher H3O+ concentration. Thus, adding more H3O+ shifts the equilibrium to the left.
The rate, or speed, at which a chemical reaction proceeds varies tremendously and cannot be predicted from the position of equilibrium. For instance, a mixture of hydrogen and oxygen can be stored indefinitely unless something, such as a spark or flame, triggers a reaction. When hydrogen and oxygen do react, the reaction proceeds with explosive violence. Conversely, diamond transforms into another crystalline form of carbon, known as graphite, but this reaction requires millions of years to occur.
Whether or not a reaction will proceed is analogous to a coin standing on its edge—the coin would be more stable on its side, but it will not tip unless pushed. In order for a reaction to occur, a certain threshold energy must be reached. This threshold is called the activation energy. As a rule, the lower the activation energy, the faster the reaction.
The amount of decrease in molecular orderliness (measured in entropy) also plays an important role in causing reactions to occur. If other factors are held equal, a chemical reaction will proceed spontaneously if the products are more disordered (have higher entropy) than the reactants.
Many reactions are not as simple as their net equations suggest (a net equation shows the overall reaction, omitting intermediate steps). Possible mechanisms for a reaction are deduced from measurements of the rate of that reaction. For the decomposition of ozone (O3) into oxygen, 2O3(g) → 3O2(g), the evidence suggests a two-step process.
In the first step, one O3 molecule dissociates, but a state of equilibrium is rapidly attained: (1) O3 ⇄ O2 + O. The second step, which proceeds more slowly, occurs when an atom of oxygen produced in step (1) reacts with another molecule of O3: (2) O + O3 → 2O2. Adding Steps (1) and (2) gives the net equation: 2O3(g) → 3O2(g) (the single oxygen atom cancels, because it occurs on both sides of the equation).
C Types of Reactions
Reactions can be classified according to several schemes, each of which is convenient for certain purposes. The three arbitrary divisions used here are ionic reactions (combining ions to form insoluble products, or products that won’t dissolve); oxidation-reduction reactions (involving the transfer of electrons); and electron-sharing reactions (involving the rearrangement of covalent bonds).
C1 Ionic Reactions
Simple ionic reactions—ionic reactions that do not involve the transfer of electrons—take place when ions are removed from solution to form insoluble (ionic) solids, gases, or covalently bonded molecules. When two soluble compounds each dissociate (in solution) into ions, and these ions subsequently combine to form an insoluble product, the reaction is driven forward to completion. This drive to completion occurs because the insoluble product cannot participate in the reverse reaction (due to its insolubility). For example, silver ions (Ag+) and chloride ions (Cl-) combine in solution to form silver chloride: Ag+(aq) + Cl-(aq) → AgCl(s). Because AgCl(s) is only slightly soluble in water, this reaction does not proceed in the opposite direction.
Double decomposition reactions occur when two reactants are each decomposed, or broken up, into a cation (positive ion) and an anion (negative ion). These ions recombine to form two or more products. The formation of insoluble AgCl(s) drives the following double decomposition reactions to completion: AgNO3(aq) + NaCl(aq) → NaNO3(aq) + AgCl(s) and Ag2SO4(aq) + 2NH4Cl(aq) → (NH4)2SO4(aq) + 2AgCl(s).
Complete reactions between soluble ionic compounds can also be driven by the formation of gases. Once a gas bubbles and escapes from an ion-containing solution, the gas cannot participate in a reverse reaction. In the following reaction, carbonic acid (H2CO3) dissociates into water and carbon dioxide gas. Because carbon dioxide (CO2) is not very soluble in water, the CO2 gas bubbles away, driving the ionic reaction forward: H2CO3(aq) → H2O + CO2(g).
Formation of a covalently bonded product may also drive an ionic reaction to completion. Such a reaction may be generally represented as follows: X+ + :Y- → X:Y (where : represents an electron pair).
When an acid ionizes in solution and produces hydrogen ions (H+), and these hydrogen ions subsequently react with hydroxide ions (OH-), which are provided by a base, water is produced (see Acids and Bases). Because water is a covalently bonded liquid and will not participate in the reverse reaction, the reaction proceeds essentially to completion: H+ + OH- → H2O(l).
Strong acids are those that dissociate completely in water, producing many more H+ ions per mole than a weak acid produces. Similarly, strong bases are those that dissociate almost completely in water, producing many more OH- ions per mole than a weak base. A strong acid, such as hydrochloric acid (HCl), will dissociate almost completely. When HCl is involved in a reaction with a strong base (such as NaOH), the large amounts of H+ ions and OH- ions will combine to produce water. Because water molecules are joined by covalent bonds, water will not participate in the reverse reaction, and the reaction proceeds nearly to completion: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l).
C2 Oxidation-Reduction Reactions
Oxidation-reduction reactions combine a chemical wanting to gain electrons with a chemical willing to give up electrons. Such a reaction may be generally represented as follows: X·+ Y ⇄ XY· (where · represents an electron). The material that loses electrons is said to be oxidized and is called a reducing agent; the material that gains electrons is reduced and is called an oxidizing agent (see Chemical Reaction). The most common examples of oxidation are those reactions involving the combination of materials with the element oxygen, such as the rusting of iron or the burning of any combustible material in air. The equation for the burning of magnesium is: 2Mg(s) + O2(g) → 2MgO(s).
When magnesium reacts with oxygen, each magnesium atom gives two electrons to oxygen. The positive magnesium ions (Mg2+) then combine with negative oxygen ions (O2-) to form solid magnesium oxide (MgO). In this reaction, magnesium (the reducing agent) is oxidized, and oxygen (the oxidizing agent) is reduced.
The reaction between metallic sodium and chlorine gas is an oxidation-reduction reaction that does not involve oxygen:
This way of writing the oxidation-reduction reaction illustrates that both elements attain a noble-gas configuration (completely filled outer shells). Sodium loses an electron, achieving the noble gas configuration of neon, and chlorine gains an electron, achieving the noble gas configuration of argon.
C3 Electron-Sharing Reactions
Electron-sharing reactions involve breaking the covalent bonds between atoms in the reactants to form new covalent bonds with different atoms. The reaction between iodine and chlorine is an example of such a reaction: I2 + Cl2 → 2ICl.
Another type of electron-sharing reaction is an addition reaction, which increases the number of groups bonded to a molecule by breaking a double or triple bond (see Chemical Bond). An example of this type of reaction is the following: CH2 = CH2 + Br2 (combined in CCl4) → CH2Br-CH2Br.
Substitution reactions, which redistribute the way electrons are shared, occur when one chemical group replaces another group on a compound: CH3-Cl + NaOH (combined in H2O(l)) → CH3-OH + NaCl.
Hydrolysis reactions are a type of electron-sharing reaction that involves the cleavage of a molecule by water: HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq).
VIII BRANCHES OF CHEMISTRY
Chemists have divided chemistry into a number of different branches. These branches are somewhat arbitrary and do not have sharply defined boundaries. They often overlap with each other or with other sciences, such as physics, geology, or biology.
Inorganic chemistry is the study of the chemical nature of the elements and their compounds (except hydrocarbons—compounds composed of carbon and hydrogen).
Organic chemistry is the study of compounds consisting largely of hydrocarbons, which provide the parent material of all other organic compounds. Since carbon atoms can form rings and long branched chains, hundreds of thousands of carbon-based molecules exist. Organic compounds are of special importance, because they make up the majority of compounds in living organisms. Organic compounds form coal and petroleum. Organic chemists have learned how to convert raw materials from coal, petroleum, and grain into synthetic textiles, pesticides, dyes, drugs, plastics, and many other products.
Radiochemistry is the study of the chemical effects of high-energy radiation and the behavior of radioactive isotopes, atoms of the same element that vary in the number of neutrons they contain. For example, the heaviest known element, Element 112 (ununbium, or Uub) was first created by scientists at the Heavy-Ion Research Laboratory in Darmstadt, Germany in 1996. These scientists created an atom of ununbium containing 165 neutrons, labeled ununbium-277 (112 protons + 165 neutrons = ununbium-277). Because the ununbium nucleus contains so many particles, the atom becomes unstable and splits into smaller, so-called daughter components. As the atom breaks apart, energy is released in the form of electromagnetic waves and electrically charged bits of matter. This energy is known as radiation (Radioactivity; Nuclear Energy).
Physical chemistry is fundamental to all chemistry and deals with the application of physical laws to chemical systems and chemical change. Much of physical chemistry is concerned with the role of energy in chemical reactions; this branch of physical chemistry is known as thermodynamics. Other major areas of study in physical chemistry are the rates and mechanisms of reactions, called chemical kinetics. A third area of physical chemistry studies molecular structure. Physical chemists study molecular structure by examining the spectrum of electromagnetic energy emitted by molecules and explain structure using principles of quantum mechanics (see Quantum Theory).
Important subfields of physical chemistry include electrochemistry, which deals with the behavior of chemical substances subjected to electric current and the production of electrical energy by chemical systems. Other subfields of physical chemistry are colloid chemistry, which is concerned with the behavior of finely divided particles of matter; surface chemistry, which deals with the nature of surfaces and adsorption on them (see Photochemistry); and statistical mechanics, which applies the laws of probability to large numbers of particles.
Analytical chemistry is the science of separating complex materials into simpler ones and detecting and measuring the constituents. In a sense, analytical chemistry is the oldest branch of chemistry. A major feature of chemical analysis today is the wide use of physical instruments and computer control to automate the analysis of complex materials.
Biochemistry is the chemistry of living organisms and life processes. Even the simplest living thing is a complex chemical factory. Biochemists must have a detailed knowledge of organic chemistry. In some aspects of biochemistry, advanced physical chemistry is used, and biophysics and molecular biology are companion sciences.
Geochemistry is the application of chemistry (and, inevitably, physics) to processes taking place in the earth, such as mineral formation, the metamorphosis of rocks, and the formation and migration of petroleum.
Fields such as biochemistry, geochemistry, and materials science reveal the unity of the sciences. The divisions between chemistry, physics, biology, and geology are arbitrarily created for the convenience of humans—nature takes little account of these divisions.
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