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Acetylene (systematic name: ethyne) is the chemical compound with the formula C2H2. It is a hydrocarbon and the simplest alkyne. This colorless gas is widely used as a fuel and a chemical building block. It is unstable in pure form and thus is usually handled as a solution.
As an alkyne, acetylene is unsaturated because its two carbon atoms are bonded together in a triple bond. The carbon-carbon triple bond places all four atoms in the same straight line, with CCH bond angles of 180Â°. Since acetylene is a linear symmetrical molecule, it possesses the Dâˆžhpoint group.
Acetylene was discovered in 1836 by Edmund Davy who identified it as a "new carburet of hydrogen". It was rediscovered in 1860 by French chemist Marcellin Berthelot, who coined the name "acetylene". Berthelot was able to prepare this gas by passing vapours of organic compounds (methanol, ethanol, etc.) through a red-hot tube and collecting the effluent. He also found acetylene was formed by sparking electricity through mixed cyanogen and hydrogen gases. Berthelot later obtained acetylene directly by passing hydrogen between the poles of a carbon arc.
Today acetylene is mainly manufactured by the partial combustion of methane or appears as a side product in the ethylene stream from cracking of hydrocarbons. Approximately 400,000 tonnes are produced this way annually. Its presence in ethylene is usually undesirable because of its explosive character and its ability to poison Ziegler-Natta catalysts. It is selectively hydrogenated into ethylene, usually using Pd-Ag catalysts.
Until the 1950s, when oil supplanted coal as the chief source of carbon, acetylene (and the aromatic fraction from coal tar) was the main source of organic chemicals in the chemical industry. It was prepared by the hydrolysis of calcium carbide, a reaction discovered by Friedrich WÃ¶hler in 1862 and still familiar to students:
- CaC2 + 2H2O â†’ Ca(OH)2 + C2H2
Calcium carbide production requires extremely high temperatures, ~2000 Â°C, necessitating the use of an electric arc furnace. In the US, this process was an important part of the late-19th century revolution in chemistry enabled by the massive hydroelectric power project at Niagara Falls.
In terms of valence bond theory, in each carbon atom the 2s orbitalhybridizes with one 2p orbital thus forming an sp hybrid. The other two 2p orbitals remain unhybridized. The two ends of the two sp hybrid orbitals overlap to form a strong Ïƒ valence bond between the carbons, while on each of the other two ends hydrogen atoms attach also by Ïƒ bonds. The two unchanged 2p orbitals form a pair of weaker Ï€ valence bonds.
One new application is the conversion of acetylene to ethylene for use in making a variety of polyethylene plastics. An important reaction of acetylene is its combustion, the basis of the acetylene welding technologies. Otherwise, its major applications involve its conversion to acrylic acid derivatives.
Compared to most hydrocarbons, acetylene is relatively acidic, though it is still much less acidic than water or ethanol. Thus it reacts with strong bases to form acetylide salts. For example, acetylene reacts with sodium amide in liquid ammonia to form sodium acetylide, and with butyllithium in cold THF to give lithium acetylide.
Acetylides of heavy metals are easily formed by reaction of acetylene with the metal ions. Several, e.g., silver acetylide and copper acetylide, are powerful and very dangerous explosives. Copper acetylide is also formed by reacting acetylene with metallic copper or its alloys; these materials are therefore unsuitable for installations for handling acetylene.
- With chemistry is a chemical bond between two chemical elements involving six bonding electrons instead of the usual two in a covalent single bond. The most common triple bond, that between two carbon atoms, can be found in alkynes. Other functional groups containing a triple bond are cyanides and isocyanides. Some diatomic molecules, such as dinitrogen and carbon monoxide are also triple bonded. In skeletal formula the triple bond is drawn as three parallel lines between the two connected atoms.
The type of bonding can be explained in terms of orbital hybridization. In the case of acetylene each carbon atom has two sp orbitals and two p-orbitals. The two sp orbitals are linear with 180Â° angles and occupy the x-axis (cartesian coordinate system). The p-orbitals are perpendicular on the y-axis and the z-axis. When the carbon atoms approach each other the sp orbitals overlap to form a sp-sp sigma bond. At the same time the pz-orbitals approach and together they form a pz-pzpi-bond. Likewise, the other pair of py-orbitals form a py-py pi-bond. The result is formation of one sigma bond and two pi bonds.
In bent bond theory the triple bond can also formed by the overlapping of three sp3 lobes without the need to invoke a pi-bond.
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 90Â° angles on the carbon atoms connected to the benzene ring where they ordinarily have angles of 120Â°.
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.
Almost everything a person sees or touches in daily lifeâ€”the air we breathe, the food we eat, the clothes we wear, and so onâ€”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"â€”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 andâ€”in Western Europe at leastâ€”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 particlesâ€”protons and neutronsâ€”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 structureâ€”still very small, but larger than the atomâ€”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 KekulÃ© (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, KekulÃ©described 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. KekulÃ© 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 bondâ€”he is usually given equal credit with KekulÃ© for these ideasâ€”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
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Answers:2 pi bonds can result from if the molecule has either 2 double bonds or 1 triple bond. carbon can at most form 4 bonds because it has four valence electrons to satisfy octet rule. c2h8 has 2 carbons single bonded and 4 hydrogens attached to each carbon? i don't even think that's possible. c3h8 has the carbons still single bonded c-c-c Answer----->c2h2 has only 1 hydrogen attached to each carbon so that means that the carbons can form 3 more bonds. so the carbons form a triple bond. So thats 2 pi bonds. c2h6 only forms a single bond btwn the carbons
Answers:Ethene H2C=CH2 Ethene has a double bond between the two carbon atoms. A double bond consists of a sigma bond, which is the end-to-end overlap of sp2 hybrid orbitals, and a pi bond, which is the side-to-side overlap of unhybridized p-orbitals. So there is 1 pi bond in ethene.
Answers:6 single, 1 double. You had to right a book for that eh guy down there.
Answers:CH3- CO- OH the oxygen on the seccond carbon is bonded with a double bond. I hope this helps