A chemical compound is a chemical substance composed of many identical molecules composed of atoms from more than one element held together by chemical bonds. A chemical element bonded to an identical chemical element is not a chemical compound since only one element, not two different elements, is involved. There are four types of compounds, depending on how the constituent atoms are held together: molecules held together by covalent bonds ionic compounds held together by ionic bonds intermetallic compounds held together by metallic bonds certain complexes held together by coordinate covalent bonds. A chemical formula is a way of expressing information about the proportions of atoms that constitute a particular chemical compound, using the standard abbreviations for the chemical elements, subscripts to indicate the number of atoms involved. For example, water is composed of two hydrogen atoms bonded to one oxygen atom: the chemical formula is H2O. Many chemical compounds have a unique numerical identifier assigned by the Chemical Abstracts Service: its CAS number.
A compound can be converted to a different chemical composition by interaction with a second chemical compound via a chemical reaction. In this process, bonds between atoms are broken in both of the interacting compounds, bonds are reformed so that new associations are made between atoms. Any substance consisting of two or more different types of atoms in a fixed stoichiometric proportion can be termed a chemical compound, it follows from their being composed of fixed proportions of two or more types of atoms that chemical compounds can be converted, via chemical reaction, into compounds or substances each having fewer atoms. The ratio of each element in the compound is expressed in a ratio in its chemical formula. A chemical formula is a way of expressing information about the proportions of atoms that constitute a particular chemical compound, using the standard abbreviations for the chemical elements, subscripts to indicate the number of atoms involved. For example, water is composed of two hydrogen atoms bonded to one oxygen atom: the chemical formula is H2O.
In the case of non-stoichiometric compounds, the proportions may be reproducible with regard to their preparation, give fixed proportions of their component elements, but proportions that are not integral. Chemical compounds have a unique and defined chemical structure held together in a defined spatial arrangement by chemical bonds. Chemical compounds can be molecular compounds held together by covalent bonds, salts held together by ionic bonds, intermetallic compounds held together by metallic bonds, or the subset of chemical complexes that are held together by coordinate covalent bonds. Pure chemical elements are not considered chemical compounds, failing the two or more atom requirement, though they consist of molecules composed of multiple atoms. Many chemical compounds have a unique numerical identifier assigned by the Chemical Abstracts Service: its CAS number. There is varying and sometimes inconsistent nomenclature differentiating substances, which include non-stoichiometric examples, from chemical compounds, which require the fixed ratios.
Many solid chemical substances—for example many silicate minerals—are chemical substances, but do not have simple formulae reflecting chemically bonding of elements to one another in fixed ratios. It may be argued that they are related to, rather than being chemical compounds, insofar as the variability in their compositions is due to either the presence of foreign elements trapped within the crystal structure of an otherwise known true chemical compound, or due to perturbations in structure relative to the known compound that arise because of an excess of deficit of the constituent elements at places in its structure. Other compounds regarded as chemically identical may have varying amounts of heavy or light isotopes of the constituent elements, which changes the ratio of elements by mass slightly. Compounds are held together through a variety of different types of bonding and forces; the differences in the types of bonds in compounds differ based on the types of elements present in the compound.
London dispersion forces are the weakest force of all intermolecular forces. They are temporary attractive forces that form when the electrons in two adjacent atoms are positioned so that they create a temporary dipole. Additionally, London dispersion forces are responsible for condensing non polar substances to liquids, to further freeze to a solid state dependent on how low the temperature of the environment is. A covalent bond known as a molecular bond, involves the sharing of electrons between two atoms; this type of bond occurs between elements that fall close to each other on the periodic table of elements, yet it is observed between some metals and nonmetals. This is due to the mechanism of this type of bond. Elements that fall close to each other on the periodic table tend to have similar electronegativities, which means they have a similar affinity for electrons. Since neither element has a stronger affinity to donate or gain electrons, it causes the elements to share electrons so both elements have a more stable octet.
Ionic bonding occurs when valence electrons are transferred between elements. Opposite to covalent bonding, this chemical bond creates two oppositely charged ions; the metals in ionic bonding
Hydrolysis is a term used for both an electro-chemical process and a biological one. The hydrolysis of water is the separation of water molecules into hydrogen and oxygen atoms using electricity. Biological hydrolysis is the cleavage of biomolecules where a water molecule is consumed to effect the separation of a larger molecule into component parts; when a carbohydrate is broken into its component sugar molecules by hydrolysis, this is termed saccharification. Hydrolysis or saccharification is a step in the degradation of a substance. Hydrolysis can be the reverse of a condensation reaction in which two molecules join together into a larger one and eject a water molecule, thus hydrolysis adds water to break down, whereas condensation builds up by removing water and any other solvents. Some hydration reactions are hydrolysis. Hydrolysis is a chemical process in which a molecule of water is added to a substance. Sometimes this addition causes both water molecule to split into two parts. In such reactions, one fragment of the target molecule gains a hydrogen ion.
It breaks a chemical bond in the compound. A common kind of hydrolysis occurs when a salt of weak base is dissolved in water. Water spontaneously ionizes into hydroxide anions and hydronium cations; the salt dissociates into its constituent anions and cations. For example, sodium acetate dissociates in water into acetate ions. Sodium ions react little with the hydroxide ions whereas the acetate ions combine with hydronium ions to produce acetic acid. In this case the net result is a relative excess of hydroxide ions. Strong acids undergo hydrolysis. For example, dissolving sulfuric acid in water is accompanied by hydrolysis to give hydronium and bisulfate, the sulfuric acid's conjugate base. For a more technical discussion of what occurs during such a hydrolysis, see Brønsted–Lowry acid–base theory. Acid–base-catalysed hydrolyses are common, their hydrolysis occurs when the nucleophile attacks the carbon of the carbonyl group of the ester or amide. In an aqueous base, hydroxyl ions are better nucleophiles than polar molecules such as water.
In acids, the carbonyl group becomes protonated, this leads to a much easier nucleophilic attack. The products for both hydrolyses are compounds with carboxylic acid groups; the oldest commercially practiced example of ester hydrolysis is saponification. It is the hydrolysis of a triglyceride with an aqueous base such as sodium hydroxide. During the process, glycerol is formed, the fatty acids react with the base, converting them to salts; these salts are called soaps used in households. In addition, in living systems, most biochemical reactions take place during the catalysis of enzymes; the catalytic action of enzymes allows the hydrolysis of proteins, fats and carbohydrates. As an example, one may consider proteases, they catalyse the hydrolysis of interior peptide bonds in peptide chains, as opposed to exopeptidases. However, proteases do not catalyse the hydrolysis of all kinds of proteins, their action is stereo-selective: Only proteins with a certain tertiary structure are targeted as some kind of orienting force is needed to place the amide group in the proper position for catalysis.
The necessary contacts between an enzyme and its substrates are created because the enzyme folds in such a way as to form a crevice into which the substrate fits. Therefore, proteins that do not fit into the crevice will not undergo hydrolysis; this specificity preserves the integrity of other proteins such as hormones, therefore the biological system continues to function normally. Upon hydrolysis, an amide converts into an amine or ammonia. One of the two oxygen groups on the carboxylic acid are derived from a water molecule and the amine gains the hydrogen ion; the hydrolysis of peptides gives amino acids. Many polyamide polymers such as nylon 6,6 hydrolyse in the presence of strong acids; the process leads to depolymerization. For this reason nylon products fail by fracturing. Polyesters are susceptible to similar polymer degradation reactions; the problem is known as environmental stress cracking. Hydrolysis is related to energy storage. All living cells require a continual supply of energy for two main purposes: the biosynthesis of micro and macromolecules, the active transport of ions and molecules across cell membranes.
The energy derived from the oxidation of nutrients is not used directly but, by means of a complex and long sequence of reactions, it is channelled into a special energy-storage molecule, adenosine triphosphate. The ATP molecule contains pyrophosphate linkages. ATP can undergo hydrolysis in two ways: the removal of terminal phosphate to form adenosine diphosphate and inorganic phosphate, or the removal of a terminal diphosphate to yield adenosine monophosphate and pyrophosphate; the latter undergoes further cleavage in
Nitroso refers to a functional group in organic chemistry which has the NO group attached to an organic moiety. As such, various nitroso groups can be categorized as C-nitroso compounds, S-nitroso compounds, N-nitroso compounds, O-nitroso compounds. Nitrosobenzene ( is prepared by reduction of nitrobenzene to phenylhydroxylamine, oxidized. Nitrosoarenes participate in a monomer-dimer equilibrium; the dimers, which are pale-yellow, are favored in the solid state, whereas the deep-green monomers are favored in dilute solution or at higher temperatures. They exist as cis- and trans-isomers. Nitroso compounds can be prepared by the reduction of nitro compounds or by the oxidation of hydroxylamines. A good example is 3CNO, known formally as 2-methyl-2-nitrosopropane, or t-BuNO. 3CNO is blue and exists in solution in equilibrium with its dimer, colorless, m.p. 80–81 °C. In the Fischer–Hepp rearrangement aromatic 4-nitrosoanilines are prepared from the corresponding nitrosamines. Another named reaction involving a nitroso compound is the Barton reaction.
Organonitroso compounds serve as a ligands for transition metals. Due to the stability of the nitric oxide free radical, nitroso organyls tend to have low C–N bond dissociation energies: nitrosoalkanes have BDEs on the order of 30 to 40 kcal/mol, while nitrosoarenes have BDEs on the order of 50 to 60 kcal/mol; as a consequence, they are heat- and light-sensitive. Compounds containing O– or N– bonds have lower bond dissociation energies. For instance, N-nitrosodiphenylamine, Ph2N–N=O, has a N–N bond dissociation energy of only 23 kcal/mol. Nitrite can enter two kinds of reaction, depending on the physico-chemical environment. Nitrosylation is adding a nitrosyl ion NO− to a metal or a thiol, leading to nitrosyl iron Fe–NO or S-nitrosothiols. Nitrosation is adding a nitrosonium ion NO+ to an amine –NH2 leading to a nitrosamine; this conversion occurs at acidic pH in the stomach, as shown in the equation for the formation of N-phenylnitrosamine: NO−2 + H+ ⇌ HONO HONO + H+ ⇌ H2O + NO+ C6H5NH2 + NO+ → C6H5NNO + H+Many primary alkyl N-nitroso compounds, such as CH3NNO, tend to be unstable with respect to hydrolysis to the alcohol.
Those derived from secondary amines are more robust. It is these N-nitrosamines. Nitrosyls are non-organic compounds containing the NO group, for example directly bound to the metal via the N atom, giving a metal–NO moiety. Alternatively, a nonmetal example is the common reagent nitrosyl chloride. Nitric oxide is a stable radical, having an unpaired electron. Reduction of nitric oxide gives the hyponitrite anion, NO−: NO + e− → NO−Oxidation of NO yields the nitrosonium cation, NO+: NO → NO+ + e−Nitric oxide can serve as a ligand forming metal nitrosyl complexes or just metal nitrosyls; these complexes can be viewed as adducts of NO −, or some intermediate case. In foodstuffs and in the gastro-intestinal tract and nitrosylation do not have the same consequences on consumer health. In cured meat: Meat processed by curing contains nitrite and has a pH of 5 where all nitrite is present as NO−2. Cured meat is added with sodium ascorbate; as demonstrated by S. Mirvish, ascorbate inhibits nitrosation of amines to nitrosamine, because ascorbate reacts with NO−2 to form NO.
Ascorbate and pH 5 thus favor nitrosylation of heme iron, forming nitrosylheme, a red pigment when included inside myoglobin, a pink pigment when it has been released by cooking. It participates to the "bacon flavor" of cured meat: nitrosylheme is thus considered a benefit for the meat industry and for consumers. In the stomach: secreted hydrogen chloride makes an acidic environment and ingested nitrite leads to nitrosation of amines, that yields nitrosamines. Nitrosation is low if vitamin C concentration is high. S-nitrosothiols are formed, that are stable at pH 2. In the colon: neutral pH does not favor nitrosation. No nitrosamine is formed in stools after addition of a secondary amine or nitrite. Neutral pH favors NO− release from S-nitrosothiols, nitrosylation of iron; the called NOC measured by Bingham's team in stools from red meat-fed volunteers were, according to Bingham and Kuhnle non-N-nitroso ATNC, e.g. S-nitrosothiols and nitrosyl iron. Nitrosamine, the functional group with the NO attached to an amine, such as R2N–NO Nitrosobenzene Nitric oxide Nitroxyl
In organic chemistry, a hydrocarbon is an organic compound consisting of hydrogen and carbon. Hydrocarbons are examples of group 14 hydrides. Hydrocarbons from which one hydrogen atom has been removed are functional groups called hydrocarbyls; because carbon has 4 electrons in its outermost shell carbon has four bonds to make, is only stable if all 4 of these bonds are used. Aromatic hydrocarbons, alkanes and alkyne-based compounds are different types of hydrocarbons. Most hydrocarbons found on Earth occur in crude oil, where decomposed organic matter provides an abundance of carbon and hydrogen which, when bonded, can catenate to form limitless chains; as defined by IUPAC nomenclature of organic chemistry, the classifications for hydrocarbons are: Saturated hydrocarbons are the simplest of the hydrocarbon species. They are composed of single bonds and are saturated with hydrogen; the formula for acyclic saturated hydrocarbons is CnH2n+2. The most general form of saturated hydrocarbons is CnH2n +2.
Those with one ring are the cycloalkanes. Saturated hydrocarbons are the basis of petroleum fuels and are found as either linear or branched species. Substitution reaction is their characteristics property. Hydrocarbons with the same molecular formula but different structural formulae are called structural isomers; as given in the example of 3-methylhexane and its higher homologues, branched hydrocarbons can be chiral. Chiral saturated hydrocarbons constitute the side chains of biomolecules such as chlorophyll and tocopherol. Unsaturated hydrocarbons have one or more triple bonds between carbon atoms; those with double bond are called alkenes. Those with one double bond have the formula CnH2n; those containing triple bonds are called alkyne. Those with one triple bond have the formula CnH2n−2. Aromatic hydrocarbons known as arenes, are hydrocarbons that have at least one aromatic ring. Hydrocarbons can be gases, waxes or low melting solids or polymers; because of differences in molecular structure, the empirical formula remains different between hydrocarbons.
This inherent ability of hydrocarbons to bond to themselves is known as catenation, allows hydrocarbons to form more complex molecules, such as cyclohexane, in rarer cases, arenes such as benzene. This ability comes from the fact that the bond character between carbon atoms is non-polar, in that the distribution of electrons between the two elements is somewhat due to the same electronegativity values of the elements, does not result in the formation of an electrophile. With catenation comes the loss of the total amount of bonded hydrocarbons and an increase in the amount of energy required for bond cleavage due to strain exerted upon the molecule. In simple chemistry, as per valence bond theory, the carbon atom must follow the 4-hydrogen rule, which states that the maximum number of atoms available to bond with carbon is equal to the number of electrons that are attracted into the outer shell of carbon. In terms of shells, carbon consists of an incomplete outer shell, which comprises 4 electrons, thus has 4 electrons available for covalent or dative bonding.
Hydrocarbons are hydrophobic like lipids. Some hydrocarbons are abundant in the solar system. Lakes of liquid methane and ethane have been found on Titan, Saturn's largest moon, confirmed by the Cassini-Huygens Mission. Hydrocarbons are abundant in nebulae forming polycyclic aromatic hydrocarbon compounds. Hydrocarbons are a primary energy source for current civilizations; the predominant use of hydrocarbons is as a combustible fuel source. In their solid form, hydrocarbons take the form of asphalt. Mixtures of volatile hydrocarbons are now used in preference to the chlorofluorocarbons as a propellant for aerosol sprays, due to chlorofluorocarbons' impact on the ozone layer. Methane and ethane are gaseous at ambient temperatures and cannot be liquefied by pressure alone. Propane is however liquefied, exists in'propane bottles' as a liquid. Butane is so liquefied that it provides a safe, volatile fuel for small pocket lighters. Pentane is a colorless liquid at room temperature used in chemistry and industry as a powerful nearly odorless solvent of waxes and high molecular weight organic compounds, including greases.
Hexane is a used non-polar, non-aromatic solvent, as well as a significant fraction of common gasoline. The C6 through C10 alkanes and isomeric cycloalkanes are the top components of gasoline, jet fuel and specialized industrial solvent mixtures. With the progressive addition of carbon units, the simple non-ring structured hydrocarbons have higher viscosities, lubricating indices, boiling points, solidification temperatures, deeper color. At the opposite extreme from methane lie the heavy tars that remain as the lowest fraction in a crude oil refining retort, they are collected and utilized as roofing comp
Cucurbiturils are macrocyclic molecules made of glycoluril monomers linked by methylene bridges. The oxygen atoms are located along the edges of the band and are tilted inwards, forming a enclosed cavity; the name is derived from the resemblance of this molecule with a pumpkin of the family of Cucurbitaceae. Cucurbiturils are written as cucurbituril, where n is the number of glycoluril units. Two common abbreviations are CB, or CBn; these compounds are interesting to chemists because they are suitable hosts for an array of neutral and cationic species. The binding mode is thought to occur through hydrophobic interactions, and, in the case of cationic guests, through cation-dipole interactions as well; the dimensions of cucurbiturils are on the ~ 10 Å size scale. For instance, the cavity of cucurbituril has a height ~ 9.1 Å, an outer diameter ~ 5.8 Å, an inner diameter ~ 3.9 Å. Cucurbiturils were first synthesized in 1905 by Robert Behrend, by condensing glycoluril with formaldehyde, but their structure was not elucidated until 1981.
The field expanded as CB5, CB7, CB8 were discovered and isolated by Kim Kimoon in the year 2000. To date cucurbiturils composed of 5, 6, 7, 8, 10, 14 repeat units have all been isolated, which have internal cavity volumes of 82, 164, 279, 479, 870 Å3 respectively. A cucurbituril composed of 9 repeat units has yet to be isolated. Other common molecular capsules that share a similar molecular shape with cucurbiturils include cyclodextrins and pillararenes. Cucurbiturils are amidals and synthesized from urea 1 and a dialdehyde via a nucleophilic addition to give the intermediate glycoluril 3; this intermediate is condensed with formaldehyde to give hexamer cucurbituril above 110 °C. Ordinarily, multifunctional monomers such as 3 would undergo a step-growth polymerization that would give a distribution of products, but due to favorable strain and an abundance of hydrogen bonding, the hexamer is the only reaction product isolated after precipitation. Decreasing the temperature of the reaction to between 75 and 90 °C can be used to access other sizes of cucurbiturils including CB, CB, CB, CB.
CB is still the major product. The isolation of sizes other than CB requires fractional dissolution. CB, CB, CB, CB are all commercially available; the larger sizes are a active area of research since they can bind larger and more interesting guest molecules, thus expanding their potential applications. Cucurbituril is difficult to isolate, it was first discovered by Day and coworkers in 2002 as an inclusion complex containing CB by fractional crystallization of the cucurbituril reaction mixture. The CB•CB was unambiguously identified by single crystal X-ray structural analysis that revealed the complex resembled a molecular gyroscope. In this case, the free rotation of the CB within the CB cavity mimics the independent rotation of a flywheel within the frame of a gyroscope. Isolation of pure CB could not be accomplished by direct separation methods since the compound has such a high affinity for CB; the strong binding affinity for the CB can be understood since it has a complementary size and shape to the cavity of the CB.
Pure CB was isolated by Isaacs and coworkers in 2005 by introducing a more binding melamine diamine guest, capable of displacing the CB. The melamine diamine guest was separated from the CB by reaction with acetic anhydride that converted the positively charged amine groups to neutrally charged amides. Cucurbiturils bind cationic guests, but by removing the positive charge from the melamine diamine guest reduces the association constant to the point it can be removed by washing with methanol, DMSO, water; the CB has an unusually large cavity that's free and capable of binding extraordinarily large guests including a cationic calixarene. Cucurbiturils have been used by chemists for various applications, including drug delivery, asymmetric synthesis, molecular switching, dye tuning. Cucurbiturils are efficient host molecules in molecular recognition and have a high affinity for positively charged or cationic compounds. High association constants with positively charged molecules are attributed to the carbonyl groups that line each end of the cavity and can interact with cations in a similar fashion to crown ethers.
The affinity of cucurbiturils can be high. For example, the affinity equilibrium constant of cururbituril with the positively charged 1-aminoadamantane hydrochloride is experimentally determined at 4.23*1012. Host guest interactions significantly influence solubility behavior of cucurbiturils. Cucurbituril dissolves poorly in just about any solvent but solubility is improved in a solution of potassium hydroxide or in an acidic solution; the cavitand forms a positively charged inclusion compound with a potassium ion or a hydronium ion which have much greater solubility than the uncomplexed neutral molecule. CB is large enough to hold other molecular hosts such as a calixarene molecule. With a calixarene guest different chemical conformations are in rapid equilibrium. Allosteric control is provided when an adamantane molecule forces a cone conformation with a calixarene - adamantane inclusion complex within a CB molecule. Given their high affinities to form inclusion complexes cucurbiturils have been employed as the macrocycles component of a rotaxane.
After formation of the supramolecular assembly or threaded complex with a guest molecule
An aminal or aminoacetal is a functional group or type of organic compound that has two amine groups attached to the same carbon atom: -C-.. The aminal and the hemiaminal groups are analogous to hemiacetals and acetals with nitrogen replaced by oxygen. Aminals are encountered for instance, the Fischer indole synthesis. Cyclic aminals are well known, being derived by the condensation of a diamine and an aldehyde. An example is hexamethylenetetramine derived from formaldehyde. Hemiaminal ethers are sometimes called aminals although it is discouraged by the IUPAC; the ethers have the following structure: R‴-C-R⁗. The glycosylamines are examples of cyclic hemiaminal ethers. Acetal Hemiaminal
In organic chemistry, an alkene is an unsaturated hydrocarbon that contains at least one carbon–carbon double bond. The words alkene and olefin are used interchangeably. Acyclic alkenes, with only one double bond and no other functional groups, known as mono-enes, form a homologous series of hydrocarbons with the general formula CnH2n. Alkenes have two hydrogen atoms fewer than the corresponding alkane; the simplest alkene, with the International Union of Pure and Applied Chemistry name ethene, is the organic compound produced on the largest scale industrially. Aromatic compounds are drawn as cyclic alkenes, but their structure and properties are different and they are not considered to be alkenes. Like a single covalent bond, double bonds can be described in terms of overlapping atomic orbitals, except that, unlike a single bond, a carbon–carbon double bond consists of one sigma bond and one pi bond; this double bond is stronger than a single covalent bond and shorter, with an average bond length of 1.33 ångströms.
Each carbon of the double bond uses its three sp2 hybrid orbitals to form sigma bonds to three atoms. The unhybridized 2p atomic orbitals, which lie perpendicular to the plane created by the axes of the three sp² hybrid orbitals, combine to form the pi bond; this bond lies outside the main C–C axis, with half of the bond on one side of the molecule and half on the other. With a strength of 65 kcal/mol, the pi bond is weaker than the sigma bond. Rotation about the carbon–carbon double bond is restricted because it incurs an energetic cost to break the alignment of the p orbitals on the two carbon atoms; as a consequence, substituted alkenes may exist as one of called cis or trans isomers. More complex alkenes may be named with the E–Z notation for molecules with three or four different substituents. For example, of the isomers of butene, the two methyl groups of -but-2-ene appear on the same side of the double bond, in -but-2-ene the methyl groups appear on opposite sides; these two isomers of butene are different in their chemical and physical properties.
Twisting to a 90° dihedral angle between two of the groups on the carbons requires less energy than the strength of a pi bond, the bond still holds. The carbons of the double bond become pyramidal, which allows preserving some p orbital alignment—and hence pi bonding; the other two attached. This contradicts a common textbook assertion that the two carbons retain their planar nature when twisting, in which case the p orbitals would rotate enough away from each other to be unable to sustain a pi bond. In a 90°-twisted alkene, the p orbitals are only misaligned by 42° and the strain energy is only around 40 kcal/mol. In contrast, a broken pi bond has an energetic cost of around 65 kcal/mol; some pyramidal alkenes are stable. For example, trans-cyclooctene is a stable strained alkene and the orbital misalignment is only 19°, despite having a significant dihedral angle of 137° and a degree of pyramidalization of 18°. Trans-cycloheptene is stable at low temperatures; as predicted by the VSEPR model of electron pair repulsion, the molecular geometry of alkenes includes bond angles about each carbon in a double bond of about 120°.
The angle may vary because of steric strain introduced by nonbonded interactions between functional groups attached to the carbons of the double bond. For example, the C–C–C bond angle in propylene is 123.9°. For bridged alkenes, Bredt's rule states that a double bond cannot occur at the bridgehead of a bridged ring system unless the rings are large enough. Following Fawcett and defining S as the total number of non-bridgehead atoms in the rings, bicyclic systems require S ≥ 7 for stability and tricyclic systems require S ≥ 11. Many of the physical properties of alkenes and alkanes are similar: they are colourless and combustable; the physical state depends on molecular mass: like the corresponding saturated hydrocarbons, the simplest alkenes, ethene and butene are gases at room temperature. Linear alkenes of five to sixteen carbons are liquids, higher alkenes are waxy solids; the melting point of the solids increases with increase in molecular mass. Alkenes have stronger smells than the corresponding alkane.
Ethylene is described to have a "sweet" odor, for example. The binding of cupric ion to the olefin in the mammalian olfactory receptor MOR244-3 is implicated in the smell of alkenes. Strained alkenes, in particular, like norbornene and trans-cyclooctene are known to have strong, unpleasant odors, a fact consistent with the stronger π complexes they form with metal ions including copper. Alkenes are stable compounds, but are more reactive than alkanes, either because of the reactivity of the carbon–carbon pi-bond or the presence of allylic CH centers. Most reactions of alkenes involve additions to this pi bond. Alkenes serve as a feedstock for the petrochemical industry because they can participate in a wide variety of reactions, prominently polymerization and alkylation. Alkenes react in many addition reactions. Most of these addition reactions follow the mechanism of electrophilic addition. Examples are hydrohalogenation, halohydrin formation, hydroboration, dichlorocarbene addition, Simmons–Smith reaction, catalytic hydrogenation, epox