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From Wikipedia

Bond length

In molecular geometry, bond length or bond distance is the average distance between nuclei of two bondedatoms in a molecule.


Bond length is related to bond order, when more electrons participate in bond formation the bond will get shorter. Bond length is also inversely related to bond strength and the bond dissociation energy, as a stronger bond is also a shorter bond, however, there are also few exceptions (Ex, H-H and H-O, the latter one has longer and stronger bond). In a bond between two identical atoms half the bond distance is equal to the covalent radius. Bond lengths are measured in molecules by means of X-ray diffraction. A set of two atoms sharing a bond is unique going from one molecule to the next. For example the carbon to hydrogen bond in methane is different from that in methyl chloride. It is however possible to make generalizations when the general structure is the same.

Bond lengths of carbon with other elements

A table with experimental single bonds for carbon to other elements is given below. Bond lengths are given in picometers. By approximation the bond distance between two different atoms is the sum of the individual covalent radii (these are given in the chemical element articles for each element). As a general trend, bond distances decrease across the row in the periodic table and increase down a group. This trend is identical to that of the atomic radius.

Bond lengths in organic compounds

The actual bond length between two atoms in a molecule depends on such factors as the orbital hybridization and the electronic and steric nature of the substituents. The carbon-carbon bond length in diamond is 154 pm which is also the largest bond length that exists for ordinary carbon covalent bonds.

Unusually long bond lengths do exist. In one, tricyclobutabenzene, a bond length of 160 pm is reported. The current record holder is another cyclobutabenzene with length 174 pm based on X-ray crystallography. In this type of compounds the cyclobutane ring would force 90° angles on the carbon atoms connected to the benzene ring where they ordinarily have angles of 120°.

The existence of a very long C-C bond length of up to 290 pm is claimed in a dimer of two tetracyanoethylenedianions although this concerns a 2-electron-4-center bond. This type of bonding has also been observed in dimers of neutral phenalene dimers. The bond lengths of these so-called pancake bonds are up to 305 pm.

Shorter than average carbon carbon bonds distances are also possible, alkenes and alkynes have bond lengths of respectively 133 and 120 pm due to increased s-character of the sigma bond. In benzene all bonds have the same length: 139 pm. In carbon carbon single bonds increased s-character is also notable in the central bond of diacetylene (137 pm) and that of a certain tetrahedrane dimer (144 pm).

In propionitrile the cyano group withdraws electrons also resulting in a reduced bond length (144 pm). Squeezing a CC bond is also possible by application of strain. An unusual organic compound exists called In-Methylcyclophane with a very short bond distance of 147 pm for the methyl group being squeezed between a trypticene and a phenyl group. In an in silico experiment a bond distance of 136 pm is estimated for neopentane locked up in fullerene. The smallest theoretical CC single bond obtained in this study is 131 pm for a hypothetical tetrahedrane derivative.

In the same study, it is estimated that for ethane it takes 2.8 kJ/mol to stretch the CC bond by 5 pm from its equilibrium value and only 3.5 kJ/mol to squeeze it by the same amount. On the other hand, stretching and squeezing by 15 pm requires 21.9 and 37.7 kJ/mol.

Ionic bond

An ionic bond is a type of chemical bond that involves a metal and a nonmetalion (or polyatomic ions such as ammonium) through electrostatic attraction. In short, it is a bond formed by the attraction between two oppositely charged ions.

The metal donates one or more electrons, forming a positively charged ion or cation with a stable electron configuration. These electrons then enter the non metal, causing it to form a negatively charged ion or anion which also has a stable electron configuration. The electrostatic attraction between the oppositely charged ions causes them to come together and form a bond.

For example, common table salt is sodium chloride. When sodium (Na) and chlorine (Cl) are combined, the sodium atoms each lose an electron, forming cations (Na+), and the chlorine atoms each gain an electron to form anions (Cl−). These ions are then attracted to each other in a 1:1 ratio to form sodium chloride (NaCl).

Na + Cl → Na+ + Cl−→ NaCl

The removal of electrons from the atoms is endothermic and causes the ions to have a higher energy. There may also be energy changes associated with breaking of existing bonds or the addition of more than one electron to form anions. However, the attraction of the ions to each other lowers their energy. Ionic bonding will occur only if the overall energy change for the reaction is favourable – when the bonded atoms have a lower energy than the free ones. The larger the resulting energy change the stronger the bond. The low electronegativity of metals and high electronegativity of non-metals means that the energy change of the reaction is most favorable when metals lose electrons and non-metals gain electrons.

Pure ionic bonding is not known to exist. All ionic compounds have a degree of covalent bonding. The larger the difference in electronegativity between two atoms, the more ionic the bond. Ionic compounds conduct electricity when molten or in solution. They generally have a high melting point and tend to be soluble in water.

Ionic structure

Ionic compounds in the solid state form lattice structures. The two principal factors in determining the form of the lattice are the relative charges of the ions and their relative sizes. Some structures are adopted by a number of compounds; for example, the structure of the rock salt sodium chloride is also adopted by many alkali halides, and binary oxides such as MgO.

Strength of an ionic bond

For a solid crystalline ionic compound the enthalpy change in forming the solid from gaseous ions is termed the lattice energy. The experimental value for the lattice energy can be determined using the Born-Haber cycle. It can also be calculated using the Born-Landé equation as the sum of the electrostatic potential energy, calculated by summing interactions between cations and anions, and a short range repulsive potential energy term. The electrostatic potential can be expressed in terms of the inter-ionic separation and a constant (Madelung constant) that takes account of the geometry of the crystal. The Born-Landé equation gives a reasonable fit to the lattice energy of e.g. sodium chloride where the calculated value is −756 kJ/mol which compares to −787 kJ/mol using the Born-Haber cycle.

Polarization effects

Ions in crystal lattices of purely ionic compounds are spherical; however, if the positive ion is small and/or highly charged, it will distort the electron cloud of the negative ion, an effect summarised in Fajans' rules. This polarization of the negative ion leads to a build-up of extra charge density between the two nuclei, i.e., to partial covalency. Larger negative ions are more easily polarized, but the effect is usually only important when positive ions with charges of 3+ (e.g., Al3+) are involved. However, 2+ ions (Be2+) or even 1+ (Li+) show some polarizing power because their sizes are so small (e.g., LiI is ionic but has some covalent bonding present). Note that this is not the ionic polarization effect which refers to displacement of ions in the lattice due to the application of an electric field.

Ionic versus covalent bonds

In an ionic bond, the atoms are bound by attraction of opposite ions, whereas, in a covalent bond, atoms are bound by sharing electrons. In covalent bonding, the molecular geometry around each atom is determined by VSEPR rules, whereas, in ionic materials, the geometry follows maximum packing rules.

In reality, purely ionic bonds do no

From Encyclopedia


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. 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. 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. 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. 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 all-important 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. 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 (1776-1856). 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. 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. 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 (1856-1940). 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 critic


energy in physics, the ability or capacity to do work or to produce change. Forms of energy include heat , light , sound , electricity , and chemical energy. Energy and work are measured in the same units—foot-pounds, joules, ergs, or some other, depending on the system of measurement being used. When a force acts on a body, the work performed (and the energy expended) is the product of the force and the distance over which it is exerted. Potential and Kinetic Energy Potential energy is the capacity for doing work that a body possesses because of its position or condition. For example, a stone resting on the edge of a cliff has potential energy due to its position in the earth's gravitational field. If it falls, the force of gravity (which is equal to the stone's weight; see gravitation ) will act on it until it strikes the ground; the stone's potential energy is equal to its weight times the distance it can fall. A charge in an electric field also has potential energy because of its position; a stretched spring has potential energy because of its condition. Chemical energy is a special kind of potential energy; it is the form of energy involved in chemical reactions. The chemical energy of a substance is due to the condition of the atoms of which it is made; it resides in the chemical bonds that join the atoms in compound substances (see chemical bond ). Kinetic energy is energy a body possesses because it is in motion. The kinetic energy of a body with mass m moving at a velocity v is one half the product of the mass of the body and the square of its velocity, i.e., KE = 1/2 mv2 . Even when a body appears to be at rest, its atoms and molecules are in constant motion and thus have kinetic energy. The average kinetic energy of the atoms or molecules is measured by the temperature of the body. The difference between kinetic energy and potential energy, and the conversion of one to the other, is demonstrated by the falling of a rock from a cliff, when its energy of position is changed to energy of motion. Another example is provided in the movements of a simple pendulum (see harmonic motion ). As the suspended body moves upward in its swing, its kinetic energy is continuously being changed into potential energy; the higher it goes the greater becomes the energy that it owes to its position. At the top of the swing the change from kinetic to potential energy is complete, and in the course of the downward motion that follows the potential energy is in turn converted to kinetic energy. Conversion and Conservation of Energy It is common for energy to be converted from one form to another; however, the law of conservation of energy, a fundamental law of physics, states that although energy can be changed in form it can be neither created nor destroyed (see conservation laws ). The theory of relativity shows, however, that mass and energy are equivalent and thus that one can be converted into the other. As a result, the law of conservation of energy includes both mass and energy. Many transformations of energy are of practical importance. Combustion of fuels results in the conversion of chemical energy into heat and light. In the electric storage battery chemical energy is converted to electrical energy and conversely. In the photosynthesis of starch, green plants convert light energy from the sun into chemical energy. Hydroelectric facilities convert the kinetic energy of falling water into electrical energy, which can be conveniently carried by wires to its place of use (see power, electric ). The force of a nuclear explosion results from the partial conversion of matter to energy (see nuclear energy ).

From Yahoo Answers

Question:for the industrial synthesis of isopropyl alcohol (rubbing alcohol) by reaction of water with propene.

Answers:76 kJ

Question:How much energy would be required to break a single C-H bond? A. 3.98 x10^-32 kJ B. 3.54 x10^-22 kJ C. 213 kJ D. None of these

Answers:213kJ/mol/6.02 E23/mole = 35.4 E-23 = 3.54 E-22 KJ answerB

Question:Since hydrochloric acid(HCL) has stong covalent bonds, why does it dissociate so easily into H+ ions and Cl- ions? (I know it's an acid, but it has strong covalent bonding among atoms!)

Answers:the term HCl ---> H+ + Cl- is the most misleading term that someone could learn in highschool level. But first of all VERY GOOD QUESTION! The term that HCl 'dissociates' in water actually doesn't mean that suddenly HCl molecule breaks into H+ and Cl-, because if that is the case one will need to supply enormous amount of energy. What actually happens is that HCl REACTS with water. since water is considered to be strong basic when compared to HCl, so what really happens is HCl + H2O --> [H3O+][Cl-] and this reaction is an exothermal reaction, that's why when you dilute HCl with water the you'll notice that the reaction glass will be slightly warmer. Since in the past time the most common solvent used in chemistry was water, the definition of acid was also used relatively to water (bronsted acidity concept) and in many case people start to leave out the complete reaction equation and shortened it with HCl ---> H+ + Cl- , it's not wrong though as long as you know that it only applies in aqueous solution and that what really happens is reaction between HCl and water and NOT self-breaking covalent bond of HCl in the more modern chemistry time, more organic solvent are involved and that's why the definition of more universal Lewis acidity become more useful, but this is not in the scope of your question. I hope that would help

Question:In an organic molecule such as Toluene, one of the C-H bonds in the Methyl group attached to the Benzene ring breaks in order to leave a lone pair of electrons on carbon which causes the Resonance to begin. However, C-H bond dissociation energy if around 104kCal. Where does this energy come from?

Answers:Depends. What reaction are you doing? How are you breaking the bond? Toluene is stable. It doesn't spontaneously generate radicals at room temperature by homolyzing that C-H bond. The energy doesn't "come from" anywhere, which is why the bond ain't breaking. Your question implies that you've done something to break the bond, but unless you specify the exact chemical change and the conditions under which it occurs, it's impossible to say what the driving force is. Maybe you're heating it. Maybe you're bombarding it with high energy photons. Maybe you're reacting it with something that wants that H atom more that toluene does. Dunno. Your question is also confusing two concepts. The C-H BDE refers to a very specific chemical process: gas phase toluene turns into a gas phase hydrogen atom (H-dot) and a gas phase toluene radical (C6H5CH2-dot). Uphill by 104 kcal/mol. Your mention of a "lone pair of electrons on C" implies not C-H homolysis (bond strength), but H+ dissociation (acidity), which is a different reaction: products of H+ and [C6H5CH2] . You can't use bond strength data (alone) to predict acidities or energies of heterolytic bond cleavage: the gas phase loss of H+ to leave a lone pair on C is not an enegy change of 104 kcal.

From Youtube

Radical Stability Trends, BDEs, & Initiation :This webcast discusses radical stability trends, the importance of considering bond dissociation energies, and introduces radical initiators, which kick off radical reactions by undergoing homolytic cleavage.

37. Potential Energy Surfaces, Transition State Theory and Reaction Mechanism :Freshman Organic Chemistry (CHEM 125) After discussing the statistical basis of the law of mass action, the lecture turns to developing a framework for understanding reaction rates. A potential energy surface that associates energy with polyatomic geometry can be realized physically for a linear, triatomic system, but it is more practical to use collective energies for starting material, transition state, and product, together with Eyring theory, to predict rates. Free-radical chain halogenation provides examples of predicting reaction equilibria and rates from bond dissociation energies. The lecture concludes with a summary of the semester's topics from the perspective of physical-organic chemistry. Complete course materials are available at the Open Yale Courses website: open.yale.edu This course was recorded in Fall 2008.