Manila City

Oxidation number In chemistry, the oxidation state is an indicator of the degree of oxidation of an atom in a chemical compound. The formal oxidation state is the hypothetical charge that an atom would have if all bonds to atoms of different elements were 100% ionic. Oxidation states are represented by Arabic numerals and can be positive, negative, or zero. Thus, H+ would have an oxidation state of 1+. The increase in oxidation state of an atom is known as an oxidation: a decrease in oxidation state is known as a reduction. Such reactions involve the transfer of electrons, a net gain in electrons being a reduction and a net loss of electrons being an oxidation. The oxidation number of an element in a molecule or complex is the charge that it would have if all the ligands (basically, atoms that donate electrons) were removed along with the electron pairs that were shared with the central atom[1]. It means that the oxidation number is the charge an atom had if it was in a compound composed of ions. It's used in the inorganic nomenclature of inorganic compounds. It is represented by a Roman numeral; the plus sign is omitted for positive oxidation numbers. The oxidation number is placed either as a right superscript to the element symbol, e.g. FeIII, or in parentheses after the name of the element, e.g. iron(III): in the latter case, there is no space between the element name and the oxidation number. The oxidation number can also be written with a number and either a + or - sign after it. If the element creates a positively charged ion, the oxidation number will have a + sign after it, (example-hydrogen 1+). If the element creates a negatively charged ion, the oxidation number will have a - sign after it, (example-oxygen 2-). The change in the oxidation number represents the number of electrons gained or lost in a chemical reaction. The oxidation number is usually numerically equal to the oxidation state of the central atom. However, for a variety of reasons, the oxidation state of transition metals can be difficult to determine[2]. The most-accepted answer is that the electron pairs forming the coordination bonds are mostly associated with the ligands: this is a good approximation for most Wernertype complexes, but much less true for organometallic compounds as well as for certain hydrido complexes, dithiolene complexes and nitrosyl complexes. Rules for the assigning of oxidation numbers 1. All species in their elemental form are given the oxidation number of zero. 2. All monoatomic ions have the same oxidation number as the charge on the ion. e.g. Mg2+ has the oxidation number of +2. 3. All combined hydrogen has an oxidation number of +1 (except metal hydrides where its oxidation number is -1). 4. All combined oxygen has an oxidation number of -2 (except peroxides where the oxidation number is -1, and compounds with fluorine where it can be positive). 5. Fluorine always has an oxidation number of -1. 6. In polyatomic species, the sum of the oxidation numbers of the element in the ion equals the charge on that species (we can use this to find the oxidation number of elements in polyatomic species). 7. Group 1 elements such as K and Na and Group 2 elements such as Mg always have a +1 and +2 oxidation state in compounds, respectively. 8. Cl has a range of oxidation states when bonded to O. However, its oxidation number is always -1 when bonded to ionic compounds.

Formal vs. spectroscopic oxidation states Although formal oxidation states can be helpful for classifying compounds, they are unmeasureable and their physical meaning can be ambiguous. Formal oxidation states require particular caution for molecules where the bonding is covalent, since the formal oxidation states require the heterolytic removal of ligands, which essentially denies covalency. Spectroscopic oxidation states, as defined by Jorgenson and reiterated by Wieghart, are measureables that are bench-marked using spectroscopic and crystallographic data.[1] Like many concepts in chemistry, spectroscopic oxidation states are powerful but require collateral

measurements. Formal oxidation states, on the other hand, result from arithmetic rules, not bonding. Skill in assigning formal oxidation states is considered essential, especially in inorganic chemistry. Calculation of formal oxidation states There are two common ways of computing the oxidation state of an atom in a compound. The first one is used for molecules when one has a Lewis structure, as is often the case for organic molecules, while the second one is used for simple compounds (molecular or not) and does not require a Lewis structure. It should be remembered that the oxidation state of an atom does not represent the "real" charge on that atom: this is particularly true of high oxidation states, where the ionization energy required to produce a multiply positive ion are far greater than the energies available in chemical reactions. The assignment of electrons between atoms in calculating an oxidation state is purely a formalism, albeit a useful one for the understanding of many chemical reactions. For more about issues with calculating atomic charges, see partial charge. From a Lewis structure When a Lewis structure of a molecule is available, the oxidation states may be assigned unambiguously by computing the difference between the number of valence electrons that a neutral atom of that element would have and the number of electrons that "belong" to it in the Lewis structure. For purposes of computing oxidation states, electrons in a bond between atoms of different elements belong to the most electronegative atom; electrons in a bond between atoms of the same element are split equally, and electrons in lone pair belong only to the atom with the lone pair. For example, consider acetic acid:

The carbon atom on the left has 6 valence electrons from its bonds to the hydrogen atoms, because carbon is more electronegative than hydrogen, and 1 electron from its bond with the other carbon atom, because the electron pair in the C–C bond is split equally, for a total of 7 electrons. A neutral carbon atom would have 4 valence electrons, because carbon is in group 14 of the periodic table. The difference, 4 – 7 = –3, is the oxidation state of that carbon atom. That is, if it is assumed that all the bonds were 100% ionic (which in fact they are not), the carbon would described as C3-. Following the same rules, the carbon on the right has an oxidation state of +3 (it only gets one valence electron from the C–C bond; the oxygen atoms get all the other electrons because oxygen is more electronegative than carbon). The oxygen atoms both have an oxidation state of –2; they get 8 electrons each (4 from the lone pairs and 4 from the bonds), while a neutral oxygen atom would have 6. The hydrogen atoms all have oxidation state +1, because they surrender their electron to the more electronegative atoms to which they are bonded. Quantum number Quantum numbers describe values of conserved quantity in the dynamics of the quantum system. They often describe specifically the energies of electrons in atoms, but other possibilities include angular momentum, spin etc. Since any quantum system can have one or more quantum numbers, it is a futile job to list all possible quantum numbers. How many quantum numbers?

The question of how many quantum numbers are needed to describe any given system has no universal answer, although for each system one must find the answer for a full analysis of the system. The dynamics of any quantum system are described by a quantum Hamiltonian, H. There is one quantum number of the system corresponding to the energy, i.e., the eigenvalue of the Hamiltonian. There is also one quantum number for each operator O that commutes with the Hamiltonian (i.e. satisfies the relation OH = HO). These are all the quantum numbers that the system can have. Note that the operators O defining the quantum numbers should be independent of each other. Often there is more than one way to choose a set of independent operators. Consequently, in different situations different sets of quantum numbers may be used for the description of the same system. Single electron in an atom This section is not meant to be a full description of this problem. For that, see the article on the Hydrogen-like atom, Bohr atom, Schrödinger equation and the Dirac equation. The most widely studied set of quantum numbers is that for a single electron in an atom: not only because it is useful in chemistry, being the basic notion behind the periodic table, Valence and a host of other properties, but also because it is a solvable and realistic problem, and, as such, finds widespread use in textbooks. In non-relativistic quantum mechanics the Hamiltonian of this system consists of the kinetic energy of the electron and the potential energy due to the Coulomb force between the nucleus and the electron. The kinetic energy can be separated into a piece which is due to angular momentum, J, of the electron around the nucleus, and the remainder. Since the potential is spherically symmetric, the full Hamiltonian commutes with J2. J2 itself commutes with any one of the components of the angular momentum vector, conventionally taken to be Jz. These are the only mutually commuting operators in this problem; hence, there are three quantum numbers. These are conventionally known as •

•

The principal quantum number (n = 1, 2, 3, 4 ...) denotes the eigenvalue of H with the J2 part removed. This number therefore has a dependence only on the distance between the electron and the nucleus (ie, the radial coordinate, r). The average distance increases with n, and hence quantum states with different principal quantum numbers are said to belong to different shells. The azimuthal quantum number (l = 0, 1 ... n−1) (also known as the angular quantum number or orbital quantum number) gives the orbital angular momentum through the relation . In chemistry, this quantum number is very important, since it specifies the shape of an atomic orbital and strongly influences chemical bonds and bond angles. In some contexts, l=0 is called an s orbital, l=1, a p orbital, l=2, a d orbital and l=3, an f orbital.

•

The magnetic quantum number (ml = −l, −l+1 ... 0 ... l−1, l) is the eigenvalue, This is the projection of the orbital angular momentum along a specified axis.

.

Results from spectroscopy indicated that up to two electrons can occupy a single orbital. However two electrons can never have the same exact quantum state. A fourth quantum number with two possible values was added to as an ad hoc assumption to resolve the conflict. It could later be explained in detail by relativistic quantum mechanics. •

The spin quantum number (ms = −1/2 or +1/2), the intrinsic angular momentum of the electron. This is the projection of the spin s=1/2 along the specified axis.

To summarize, the quantum state of an electron is determined by the quantum numbers: name

orbital range symbol value meaning of values example

principal quantum number azimuthal quantum number (angular : momentum)

: of magnetic quantum number, (projection energy shift angular momentum) spin quantum number

always only:

Example: The quantum numbers used to refer to the outermost valence electron of the Fluorine (F) atom, which is located in the 2p atomic orbital, are; n = 2, l = 1, ml = 1, or 0, or −1, ms = −1/2 or 1/2. Note that molecular orbitals require totally different quantum numbers, because the Hamiltonian and its symmetries are quite different. Quantum numbers with spin-orbit interaction For more details on this topic, see Clebsch-Gordan coefficients. When one takes the spin-orbit interaction into consideration, l, m and s no longer commute with the Hamiltonian, and their value therefore changes over time. Thus another set of quantum numbers should be used. This set includes • • •

The total angular momentum quantum number (j = 1/2,3/2 ... n−1/2) gives the total angular momentum through the relation . The projection of the total angular momentum along a specified axis (mj = -j,-j+1... j), which is analogous to m, and satisfies mj = ml + ms. Parity. This is the eigenvalue under reflection, and is positive (i.e. +1) for states which came from even l and negative (i.e. -1) for states which came from odd l. The former is also known as even parity and the latter as odd parity

For example, consider the following eight states, defined by their quantum numbers: • • • • • • • •

(1) (2) (3) (4) (5) (6) (7) (8)

l l l l l l l l

= = = = = = = =

1, 1, 1, 1, 1, 1, 0, 0,

ml ml ml ml ml ml ml ml

= = = = = = = =

1, ms = +1/2 1, ms = -1/2 0, ms = +1/2 0, ms = -1/2 -1, ms = +1/2 -1, ms = -1/2 0, ms = +1/2 0, ms = -1/2

The quantum states in the system can be described as linear combination of these eight states. However, in the presence of spin-orbit interaction, if one wants to describe the same system by eight states which are eigenvectors of the Hamiltonian (i.e. each represents a state which does not mix with others over time), we should consider the following eight states: Periodicity Periodicity is the quality of occurring at regular intervals or periods (in time or space) and can occur in different contexts: • • • • • •

A clock marks time at periodic intervals. A metronome ticks at periodic intervals of time. A publication published at periodic intervals can be called a "periodical", for example a magazine. In mathematics, a function whose output contains values that repeat periodically is called a periodic function. In chemistry, the periodic table is a table which classifies the chemical elements by means of the periodicity of their chemical properties. In physics, period is the number of cycles as a result of time (time/cycle). The amount of time it takes to complete one full revolution. Period is also the inverse of frequency.

• • •

In music theory, periodicity is described as "predictability gives rise to expectations". Standing waves crest at periodic intervals of distance. In finance, the periodicity of a loan describes the interval between payments.

The measure of periodicity in time is frequency which has the Metric units of Hertz. Chemical bonding Almost everything a person sees or touches in daily life—the air we breathe, the food we eat, the clothes we wear, and so on—is the result of a chemical bond, or, more accurately, many chemical bonds. Though a knowledge of atoms and elements is essential to comprehend the subjects chemistry addresses, the world is generally not composed of isolated atoms; rather, atoms bond to one another to form molecules and hence chemical compounds. Not all chemical bonds are created equal: some are weak, and some very strong, a difference that depends primarily on the interactions of electrons between atoms. How It Works Early Ideas of Bonding The theory that all of matter is composed of atoms did not originate in modern times: the atomic model actually dates back to the fifth century B.C. in Greece. The leading exponent of atomic theory in ancient times was Democritus (c. 460-370 B.C.), who proposed that matter could not be infinitely subdivided. At its deepest substructure, Democritus maintained, the material world was made up of tiny fragments he called atomos, a Greek term meaning "no cut" or "indivisible" Forward-thinking though it was, Democritus's idea was not what modern scientists today would describe as a proper scientific hypothesis. His "atoms" were not purely physical units, but rather idealized philosophical constructs, and thus, he was not really approaching the subject from the perspective of a scientist. In any case, there was no way for Democritus to test his hypothesis even if he had wanted to: by their very nature, the atoms he described were far too small to observe. Even today, what scientists know about atomic behavior comes not from direct observation, but indirect means. Hence, Democritus and the few other ancients who subscribed to atomic theory went more on instinct than by scientific methods. Yet, some of them were remarkably prescient in their description of the bonding of atoms, in view of the primitive scientific methods they had at their disposal. No other scientist came close to the accuracy of their theory for about 2,000 years. Asclepiades and Lucretius Discuss Bonds The physician Asclepiades of Prusa (c.130-40 B.C.) drew on the ideas of the Greek philosopher Epicurus (341-270 B.C., another proponent of atomism. Asclepiades speculated on the ways in which atoms interact, and discussed "clusters of atoms," though, of course, he had no idea what force attracted the atoms to one another. A few years after Asclepiades, the Roman philosopher and poet Lucretius (c.95-c.55 B.C. espoused views that combined atomism with the idea of the "four elements"—earth, air, fire, and water. In his great work De rerum natura ("On the Nature of Things"), Lucretius described atoms as tiny spheres attached to each other by fishhook-like appendages that became entangled with one another. Lack of Progress Until 1800 Unfortunately, the competing idea of the four elements, handed down by the great philosopher Aristotle (384-322 B.C.), prevailed over the atomic model. As the Roman Empire began to decline after A.D. 200, the pace of scientific inquiry slowed and—in Western Europe at least—eventually came to a virtual halt. Hence, the four elements theory, which had its own

fanciful explanations as to why certain "elements" bonded with one another, held sway in Europe until the beginning of the modern era. During the seventeenth century, a mounting array of facts from the realms of astronomy and physics collectively disproved the Aristotelian model. In the area of chemistry, English physicist and chemist Robert Boyle (1627-1691) showed that the four elements were not elements at all, because they could be broken down into simpler substances. Yet, no one really understood what constituted an element until the very beginning of the nineteenth century, and until that question was addressed, it was difficult to move on to the mystery of why certain atoms bonded with one another. Early Modern Advances in Bonding Theory Dalton's Atomic Theory The birth of atomic theory in modern times occurred in 1803, when English chemist John Dalton (1766-1844) formulated the idea that all elements are composed of tiny, indestructible particles. These he called by the name Democritus had given them nearly 23 centuries earlier: atoms. All known substances, he said, are composed of some combination of atoms, which differ from one another only in mass. Though Dalton's theory paved the way for enormous advances in the years that followed, there were a number of flaws in it. Mass alone, for instance, is not really what differentiates one atom from another: differences in mass reflect the presence of subatomic particles— protons and neutrons—of whose existence scientists were unaware at the time. Furthermore, the properties of atoms that cause them to bond relate to a third subatomic particle, the electron, which, though it contributes little to the mass of the atom, is allimportant to the energy it possesses. As for how atoms bond to one another, Dalton had little to say: in his conception of the atomic model, atoms simply sit adjacent to one another without forming true bonds, as such. Avogadro and the Molecule Though Dalton recognized that the structure of atoms in a particular element or compound is uniform, he maintained that compounds are made up of compound atoms: thus, water is a compound of "water atoms." However, water is not an element, and therefore, there had to be some structure—still very small, but larger than the atom—in which atoms coalesced to form the basic materials of a compound. That structure was the molecule, first described by Italian physicist Amedeo Avogadro (17761856). For several decades, Avogadro, who originated the idea of the mole as a means of comparing large groups of atoms or molecules, remained a more or less unsung hero. Only in 1860, four years after his death, was his idea of the molecule resurrected by Italian chemist Stanislao Cannizzaro (1826-1910). Cannizzaro's work was occasioned by disagreement among scientists regarding the determination of atomic mass; however, the establishment of the molecular model had far-reaching implications for theories of bonding. Symbolizing Atomic Bonds In 1858, German chemist Friedrich August Kekulé (1829-1896) made the first attempt to define the concept of valency, or the property an atom of one element possesses that determines its ability to bond with atoms of other elements. A pioneer in organic chemistry, which deals with chemical structures containing carbon, Kekulédescribed the carbon atom as tetravalent, meaning that it can bond to four other atoms. (The Latin prefix tetra-means "four.") He also speculated that carbon atoms are capable of bondingwith one another in long chains. This was one of the first attempts to examine the subject of bonding using modern scientific terminology, complete with hypotheses that could be tested by experimentation. Kekulé also recognized that in order to discuss bonds understandably, there needed to be some means of representing those bonds with symbols. He even went so far as to develop a system for showing the arrangement of bonds in space; however, his system was so elaborate that it was

replaced in favor of a simpler one developed by Scottish chemist Archibald Scott Couper (1831-1892). Couper, who also studied valency and the tetravalent carbon bond—he is usually given equal credit with Kekulé for these ideas—created an extremely straightforward schematic representation still in use by chemists today. In Couper's system, short dashed lines serve to designate chemical bonds. Hence, the bond between two hydrogen atoms and an oxygen atom in a water molecule would be represented thus: H-O-H. As the understanding of bonds progressed in modern times, this system was modified to take into account multiple bonds, discussed below. Real-Life Applications Atoms, Electrons, and Ions Today, chemical bonding is understood as the joining of atoms through electromagnetic force. Before that understanding could be achieved, however, scientists had to unlock the secret of the electromagnetic interactions that take place within an atom. The key to bonding is the electron, discovered in 1897 by English physicist J. J. Thomson (18561940). Atomic structure in general, and the properties of the electron in particular, are discussed at length elsewhere in this volume. However, because these specifics are critical to bonding, they will be presented here in the shortest possible form. At the center of an atom is a nucleus, consisting of protons, with a positive electrical charge; and neutrons, which have no charge. These form the bulk of the atom's mass, but they have little to do with bonding. In fact, the neutron has nothing to do with it, while the proton plays only a passive role, rather like a flower being pollinated by a bee. The "bee" is the electron, and, like a bee, it buzzes to and fro, carrying a powerful "sting"—its negative electric charge, which attracts it to the positively charged proton. Electrons and Ions Though the electron weighs much, much less than a proton, it possesses enough electric charge to counterbalance the positive charge of the proton. All atoms have the same number of protons as electrons, and hence the net electric charge is zero. However, as befits their highly active role, electrons are capable of moving from one atom to another under the proper circumstances. An atom that loses or acquires electrons has an electric charge, and is called an ion. The atom that has lost an electron or electrons becomes a positively charged ion, or cation. On the other hand, an atom that gains an electron or electrons becomes a negatively charged ion, or anion. As we shall see, ionic bonds, such as those that join sodium and chlorine atoms to form NaCl, or salt, are extremely powerful. Electron Configuration Even in covalent bonding, which does not involve ions, the configurations of electrons in two atoms are highly important. The basics of electron configuration are explained in the Electrons essay, though even there, this information is presented with the statement that the student should consult a chemistry textbook for a more exhaustive explanation. In the simplest possible terms, electron configuration refers to the distribution of electrons at various positions in an atom. However, because the behavior of electrons cannot be fully predicted, this distribution can only be expressed in terms of probability. An electron moving around the nucleus of an atom can be compared to a fly buzzing around some form of attractant (e.g., food or a female fly, if the moving fly is male) at the center of a sealed room. We can state positively that the fly is in the room, and we can predict that he will be most attracted to the center, but we can never predict his location at any given moment. As one moves along the periodic table of elements, electron configurations become ever more complex. The reason is that with an increase in atomic number, there is an increase in the energy levels of atoms. This indicates a greater range of energies that electrons can occupy,

as well as a greater range of motion. Electrons occupying the highest energy level in an atom are called valence electrons, and these are the only ones involved in chemical bonding. By contrast, the core electrons, or the ones closest to the nucleus, play no role in the bonding of atoms. Ionic and Covalent Bonds The Goal of Eight Valence Electrons The above discussion of the atom, and the electron's place in it, refers to much that was unknown at the time Thomson discovered the electron. Protons were not discovered for several more years, and neutrons several decades after that. Nonetheless, the electron proved the key to solving the riddle of how substances bond, and not long after Thomson's discovery, German chemist Richard Abegg (1869-1910) suggested as much. While studying noble gases, noted for their tendency not to bond, Abegg discovered that these gases always have eight valence electrons. His observation led to one of the most important principles of chemical bonding: atoms bond in such a way that they achieve the electron configuration of a noble gas. This has been shown to be the case in most stable chemical compounds. Two Different Types of Bonds Perhaps, Abegg hypothesized, atoms combine with one another because they exchange electrons in such a way that both end up with eight valence electrons. This was an early model of ionic bonding, which results from attractions between ions with opposite electric charges: when they bond, these ions "complete" one another. Ionic bonds, which occur when a metal bonds with a nonmetal are extremely strong. As noted earlier, salt is an example of an ionic bond: the metal sodium loses an electron, forming a cation; meanwhile, the nonmetal chlorine gains the electron to become an anion. Their ionic bond results from the attraction of opposite charges. Ionic bonding, however, could not explain all types of chemical bonds for the simple reason that not all compounds are ionic. A few years after Abegg's death, American chemist Gilbert Newton Lewis (1875-1946) discovered a very different type of bond, in which nonionic compounds share electrons. The result, once again, is eight valence electrons for each atom, but in this case, the nuclei of the two atoms share electrons. In ionic bonding, two ions start out with different charges and end up forming a bond in which both have eight valence electrons. In the type of bond Lewis described, a covalent bond, two atoms start out as atoms do, with a net charge of zero. Each ends up possessing eight valence electrons, but neither atom "owns" them; rather, they share electrons. Lewis Structures In addition to discovering the concept of covalent bonding, Lewis developed the Lewis structure, a means of showing schematically how valence electrons are arranged among the atoms in a molecule. Also known as the electron-dot system, Lewis structures represent the valence electrons as dots surrounding the chemical symbols of the atoms involved. These dots, which look rather like a colon, may be above or below, or on either side of, the chemical symbol. (The dots above or below the chemical symbol are side-by-side, like a colon turned at a 90°-angle.) To obtain the Lewis structure representing a chemical bond, it is first necessary to know the number of valence electrons involved. One pair of electrons is always placed between elements, indicating the bond between them. Sometimes this pair of valence electrons is symbolized by a dashed line, as in the system developed by Couper. The remaining electrons are distributed according to the rules by which specific elements bond. Multiple Bonds

Hydrogen bonds according to what is known as the duet rule, meaning that a hydrogen atom has only two valence electrons. In most other elements—there are exceptions, but these will not be discussed here—atoms end up with eight valence electrons, and thus are said to follow the octet rule. If the bond is covalent, the total number of valence electrons will not be a multiple of eight, however, because the atoms share some electrons. When carbon bonds to two oxygen atoms to form carbon dioxide (CO2), it is represented in the Couper system as O-C-O. The Lewis structure also uses dashed lines, which stand for two valence electrons shared between atoms. In this case, then, the dashed line to the left of the carbon atom indicates a bond of two electrons with the oxygen atom to the left, and the dashed line to the right of it indicates a bond of two electrons with the oxygen atom on that side. The non-bonding valence electrons in the oxygen atoms can be represented by sets of two dots above, below, and on the outside of each atom, for a total of six each. Combined with the two dots for the electrons that bond them to carbon, this gives each oxygen atom a total of eight valence electrons. So much for the oxygen atoms, but something is wrong with the representation of the carbon atom, which, up to this point, is shown only with four electrons surrounding it, not eight. In fact carbon in this particular configuration forms not a single bond, but a double bond, which is represented by two dashed lines—a symbol that looks like an equals sign. By showing the double bonds joining the carbon atom to the two oxygen atoms on either side, the carbon atom has the required number of eight valence electrons. The carbon atom may also form a triple bond (represented by three dashed lines, one above the other) with an oxygen atom, in which case the oxygen atom would have only two other valence electrons. Electronegativity and Polar Covalent Bonds Today, chemists understand that most bonds are neither purely ionic nor purely covalent; rather, there is a wide range of hybrids between the two extremes. Credit for this discovery belongs to American chemist Linus Pauling (1901-1994), who, in the 1930s, developed the concept of electronegativity—the relative ability of an atom to attract valence electrons. Elements capable of bonding are assigned an electronegativity value ranging from a minimum of 0.7 for cesium to a maximum of 4.0 for fluorine. Fluorine is capable of bonding with some noble gases, which do not bond with any other elements or each other. The greater the electronegativity value, the greater the tendency of an element to draw valence electrons to itself. If fluorine and cesium bond, then, the bond would be purely ionic, because the fluorine exerts so much more attraction for the valence electrons. But if two elements have equal electronegativity values—for instance, cobalt and silicon, both of which are rated at 1.9—the bond is purely covalent. Most bonds, as stated earlier, fall somewhere in between these two extremes. Polar Covalent Bonding When substances of differing electronegativity values form a covalent bond, this is described as polar covalent bonding. Sometimes these are simply called "polar bonds," but that is not as accurate: all ionic bonds, after all, are polar, due to the extreme differences in electronegativity. The term "polar covalent bond" is much more specific, describing a bond, for instance, between hydrogen (2.1) and sulfur (2.6). Because sulfur has a slightly greater electronegativity value, the valence electrons will be slightly more attracted to the sulfur atom than to the hydrogen atom. Another example of a polar covalent bond is the one that forms between hydrogen and oxygen (3.5) to form H2O or water, which has a number of interesting properties. For instance, the polar quality of a water molecule gives it a great attraction for ions, and thus ionic substances such as salt dissolve easily in water. "Pure" water from a mountain stream is actually filled with traces of the rocks over which it has flowed. In fact, water—sometimes called the "universal solvent"—is almost impossible to find in pure form, except when it is purified in a laboratory.

By contrast, molecules of petroleum (CH2) tend to be nonpolar, because carbon and hydrogen have almost identical electronegativity values—2.5 and 2.1 respectively. Thus, an oil molecule offers no electric charge to bond it with a water molecule, and for this reason, oil and water do not mix. It is a good thing that water molecules attract each other so strongly, because this means that a great amount of energy is required to change water from a liquid to a gas. If this were not so, the oceans and rivers would vaporize, and life on Earth could not exist as it does. Bond Energy The last two paragraphs allude to attractions between molecules, which is not the same as (nor is it as strong as) the attraction between atoms within a molecule. In fact, the bond energy—the energy required to pull apart the atoms in a chemical bond—is low for water. This is due to the presence of hydrogen atoms, with their two (rather than eight) valence electrons. It is thus relatively easy to separate water into its constituent parts of hydrogen and oxygen, through a process known as electrolysis. Covalent bonds that involve hydrogen are among the weakest bonds between atoms. (Again, this is different from bonds between molecules.) Stronger than hydrogen bonds are regular, octet-rule covalent bonds: as one might expect, double covalent bonds are stronger than single ones, and triple covalent bonds are stronger still. Strongest of all are ionic bonds, involved in the bonding of a metal to a metal, or a metal to a nonmetal, as in salt. The strength of the bond energy in salt is reflected by its boiling point of 1,472°F (800°C), much higher than that of water, at 212°F (100°C).

The force that holds atoms together in molecules and solids. Chemical bonds are very strong. To break one bond in each molecule in a mole of material typically requires an energy of many tens of kilocalories. It is convenient to classify chemical bonding into several types, although all real cases are mixtures of these idealized cases. The theory of the various bond types has been well developed and tested by theoretical chemists. See also Computational chemistry; Molecular orbital theory; Quantum chemistry. The simplest chemical bonds to describe are those resulting from direct coulombic attractions between ions of opposite charge, as in most crystalline salts. These are termed ionic bonds. See also Structural chemistry. Other chemical bonds include a wide variety of types, ranging from the very weak van der Waals attractions, which bind neon atoms together in solid neon, to metallic bonds or metallike bonds, in which very many electrons are spread over a lattice of positively charged atom cores and give rise to a stable configuration for those cores. See also Intermolecular forces. The covalent bond, in which two electrons bind two atoms together, as in

is the most characteristic link in chemistry. The theory that accounts for it is a cornerstone of chemical science. The physical and chemical properties of any molecule are direct consequences of its particular detailed electronic structure. Yet the theory of any one covalent chemical bond, for example, the HH bond in the hydrogen molecule, has much in common with the theory of any other covalent bond, for example, the OH bond in the water molecule. The current theory of covalent bonds both treats their qualitative features and quantitatively accounts for the molecular properties which are a consequence of those features. The theory is a branch of quantum theory. See also Quantum chemistry. The problem of the proper description of chemical bonds in molecules that are more complicated than H2 has many inherent difficulties. The qualitative theory of chemical bonding in complex molecules preserves the use of many chemical concepts that predate quantum chemistry itself; among these are electrostatic and steric factors, tautomerism, and

electronegativity. The quantitative theory is highly computational in nature and involves extensive use of computers and supercomputers. The number of covalent bonds which an atom can form is called the covalence and is determined by the detailed electron configuration of the atom. An extremely important case is that of carbon. In most of its compounds, carbon forms four bonds. When these connect it to four other atoms, the directions of the bonds to these other atoms normally make angles of about 109° to one another, unless the attached atoms are crowded or constrained by other bonds. That is, covalent bonds have preferred directions. However, in accord with the idea that carbon forms four bonds, it is necessary to introduce the notion of double and triple bonds. Thus in the structural formula of ethylene, C2H4 (1), all lines denote covalent bonds, the double line connecting

the carbon atoms being a double bond. Such double bonds are distinctly shorter, almost twice as stiff, and require considerably more energy to break completely than do single bonds. However, they do not require twice as much energy to break as a single bond. Similarly, acetylene (2) is written with a triple bond, which is still shorter than a double bond. A carboncarbon single bond has a length close to 1.54 × 10−8 cm, whereas the triple bond is about 1.21 × 10−8 cm long. See also Bond angle and distance; Valence. Many substances have some bonds which are covalent and others which are ionic. Thus in crystalline ammonium chloride, NH4Cl, the hydrogens are bound to nitrogen by electron pairs, but the NH4 group is a positive ion and the chlorine is a negative ion. Both electrons of a covalent bond may come from one of the atoms. Such a bond is called a coordinate or dative covalent bond or semipolar double bond, and is one example of the combination of ionic and covalent bonding. The hydrogen bond is a special bond in which a hydrogen atom links a pair of other atoms. The linked atoms are normally oxygen, fluorine, chlorine, or nitrogen. These four elements are all quite electronegative, a fact which favors a partially ionic interpretation of this kind of bonding. See also Electronegativity; Hydrogen bond. Cat ion and anions

+1 A cation (pronounced cat-eye-on) is an atom or molecule that has gained a positive charge. This often happens in water, because of the unique nature of the water molecule which, like a magnet, has positive and negative ends (it is "polar").

WATER

This is what happens to table salt, NaCl, in water:

"The importance of this polar molecule is that the water molecules can reduce the attraction between the positive cation and negative anion by orientating with their negative poles towards the cation and their positive poles towards the anion, as shown above. It is this special property that allows the two (or more) charged particles of a molecule formed by electrostatic bonding to dissociate in aqueous solution." The ocean is full of cations (and anions). There are cations in fresh water, too, especially in "hard" water. Alkali Metals have just one electron in their outer shell, so they tend to "ionize" very quickly by shedding that once electron, thereby gaining a charge of +1. Alkaline Earth Metals tend to shed two electrons and have a charge of =+2. - 1 An anion (pronounced ann-ion) is an atom or molecule that has a negative charge. Cations are abundant in seawater and hard water. The Halogens are missing one electron in their last shell, so they very quickly absorb one electron and have a charge of -1. Most of the common anions are not single atoms, however, but groups of atoms, called molecules. Example include sulfates and carbonates. Note that there is a naming convention. The "-ide" ending refers to a simple anion. The anion of chlorine is chloride. The anion for fluorine is fluoride. Sulfur becomes sulfide. Anions that contain oxygen often end in "-ate". Carbonate, sulfate and phosphate are examples. This will help you remember which is which. It is worth memorizing the more common anions because they are so common and have such a great impact upon our lives. The cations that are most common react with the anions that are most common, especially when the water they are in dries up. The products of these reactions are called "salts". The name for the compound is the name of the cation followed by the name of the anion, for example, sodium chloride. 1 The total charge of the salt should equal zero. For example: Three sodium cations (Na+) are needed to make a neutral salt with one phosphate anion (PO43 ) and the formula for the salt is then Na3PO4 Ions from water molecules - a special case. Water, H2O, can be thought of like this: H-O-H

or

H+ and

OH-

Sometimes water has more H+ ions than OH- ions. This water is said to be "acidic". In fact, acidity on the pH scale is a measurement of the number of H+ in water.

On the other hand, water can have more OH- ions, or "hydroxides ions". The water is then said to be basic, alkaline, or caustic. The pH scale, above 7, measures the number of hydroxide ions. Both hydrogen and hydroxide ions are very reactive. That is why acids and bases burn you!

Common Cations (positive charge) hydrogen H+ sodium Na+ potassium K+ calcium Ca+2 magnesium Mg+2 ferrous (iron) Fe+2 ferric (iron) Fe+

Common Anions (negative charge) hydroxide OHchloride Clsulfide S-2 bicarbonate HCO3carbonate CO3-2 sulfate SO4-2 phosphate PO4-3

Nomenclature The necessity of giving each compound a unique name requires a richer variety of terms than is available with descriptive prefixes such as n- and iso-. The naming of organic compounds is facilitated through the use of formal systems of nomenclature. Nomenclature in organic chemistry is of two types: common and systematic. In biological classification, system of naming organisms. The species to which the organism belongs is indicated by two words, the genus and species names, which are Latinized words derived from various sources. This system, which is called the Linnaean system of binomial nomenclature, was established in the 1750s by Carolus Linnaeus. Like any other field in science, genetics has its own language. However, genetics is also a multidisciplinary field that encompasses expertise, and hence terminology, from diverse areas of science, including molecular biology, statistics, clinical medicine, and, most recently, bioinformatics. Despite all of the new and changing language in the field, two of the most frequently used terms in genetics are still "chromosomes" and "genes." Humans have twenty-three pairs of chromosomes. One member of each pair is inherited from the person's mother, and the other from the father. Of the pairs, twenty-two are known as autosomes. The remaining pair consists of the sex chromosomes, which determine a person's gender. Females have two X chromosomes, and males have one X chromosome and one Y chromosome. Chromosomes are located in the nucleus of a cell. During cell division, which is known as mitosis, the chromosomes' long strands coil up tightly, to the point where they can be seen as individual units under the microscope. At this stage, each chromosome is composed of two identical strands, called chromatids (each of which further consists of two strands of nucleotides). Chromatids are attached at a constriction point called the centromere. Chromosomes can be distinguished by their size and by their "banding pattern." Researchers use a chemical staining process in the laboratory to create the banding pattern, allowing them to see the chromosomes more easily. Each chromosome is divided into two sections, or "arms," with one arm on each side of the centromere. The short arm is called the p arm, and the long arm is the q arm. The bands on each arm are numbered. As new and better staining techniques are developed, the numbering system is also refined, so that band 32, for example, would be subdivided into 32.1, then 32.15, and so on. One particular position on the long arm of chromosome 5 would be referred to as 5q32.15.

Almost every human chromosome contains more than a thousand genes. Therefore, even a small extra piece or missing piece of a chromosome results in hundreds of genes being added or deleted from an individual. When researchers study a person's chromosomes, they try to determine if there are any missing or extra chromosomes or chromosome pieces. The addition or deletion of genes sometimes causes a recognizable genetic disorder. Down syndrome, for example, results when there are three copies of chromosome 21 rather than the normal two. Genes are very small structures that lie on chromosomes. They are the instructions, or blueprints, for producing proteins, which are the building blocks that our bodies use to grow, develop, and function. Humans have an estimated thirty thousand to forty thousand genes. As happens with other types of scientific discovery, the person who discovers a gene names it, and a scientist can name a gene anything he or she wants. This has led to some confusion, as different naming schemes are used by different groups. To bring order to the situation, several international working groups are trying to standardize the naming of genes. There are separate working groups that focus on naming genes from humans, mice, fruit flies, plants, and other organisms. Some scientists choose names based on the clinical disorder that is thought to be associated with changes in the gene. For example, one gene was named CFTR because changes in its sequence are associated with the disease cystic fibrosis. Geneticists studying fruit flies traditionally use single-word names, such as wingless, hunchback, and seven less, that refer to the effect of a mutation in the gene. "Seven less" refers to the absence of the R7 protein. For human genes, abbreviations are commonly used. Abbreviated names are especially useful for genes with long names. WNT2, for example, stands for "wingless-type MMTV integration site family member 2." Although the word "wingless" seems unnecessary (humans, of course, don't have wings!), WNT2 is named after similar genes in fruit flies. Genes are often named after genes they resemble in other organisms. Sometimes the gene name is actually a variation of the name of the protein it makes. For example, the RELN gene in the human encodes the "reelin" protein. The "reelin" protein was named for the reeling motion exhibited by mice that lack a functional copy of the protein. Other genes are classified based on what their proteins do. For example, HOX genes (short for homeobox) are genes involved in development. Individual HOX genes are named with additional letters and numbers, such as HOXA1 or HOXD9. The consistent naming system lets scientists know that any gene with the name HOX is likely to play a specific role in development. There are even playful gene names that have nothing to do with a disorder or protein. An example is the SHH gene, which is involved in the development of the brain, spinal cord, and limbs. The SHH gene is named after the cartoon character Sonic the Hedgehog!