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
Langmuir adsorption model
The Langmuir adsorption model explains adsorption by assuming an adsorbate behaves as an ideal gas at isothermal conditions. At these conditions the adsorbate's partial pressure, p A, is related to the volume of it, V, adsorbed onto a solid adsorbent; the adsorbent, as indicated in the figure, is assumed to be an ideal solid surface composed of series of distinct sites capable of binding the adsorbate. The adsorbate binding is treated as a chemical reaction between the adsorbate molecule A g and an empty site, S; this reaction yields an adsorbed complex A ad with an associated equilibrium constant K eq: A g + S ↽ − − ⇀ A ad From these assumptions the Langmuir isotherm can be derived, which states that θ A = V V m = K eq A p A 1 + K eq A p A, where θ A is the fractional occupancy of the adsorption sites, V m is the volume of the monolayer. A continuous monolayer of adsorbate molecules surrounding a homogeneous solid surface is the conceptual basis for this adsorption model. In 1916, Irving Langmuir presented his model for the adsorption of species onto simple surfaces.
Langmuir was awarded the Nobel Prize in 1932 for his work concerning surface chemistry. He hypothesized that a given surface has a certain number of equivalent sites to which a species can “stick”, either by physisorption or chemisorption, his theory began when he postulated that gaseous molecules do not rebound elastically from a surface, but are held by it in a similar way to groups of molecules in solid bodies. Langmuir published two papers that proved the assumption that adsorbed films do not exceed one molecule in thickness; the first experiment involved observing electron emission from heated filaments in gases. The second, a more direct proof and measured the films of liquid on an adsorbent surface layer, he noted that the attractive strength between the surface and the first layer of adsorbed substance is much greater than the strength between the first and second layer. However, there are instances where the subsequent layers may condense given the right combination of temperature and pressure.
Inherent within this model, the following assumptions are valid for the simplest case: the adsorption of a single adsorbate onto a series of equivalent sites on the surface of the solid. The surface containing the adsorbing sites is a flat plane with no corrugations; the adsorbing gas adsorbs into an immobile state. All sites are equivalent; each site can hold at most one molecule of A. There are no interactions between adsorbate molecules on adjacent sites; this section provides a kinetic derivation for a single adsorbate case. The multiple adsorbate case is covered in the Competitive adsorption sub-section; the model assumes adsorption and desorption as being elementary processes, where the rate of adsorption rad and the rate of desorption rd are given by r ad = k ad p A, r d = k d, where PA is the partial pressure of A over the surface, is the concentration of bare sites in number/m2, is the surface concentration of A in molecules/m2, kad and kd are constants of forward adsorption reaction and backward desorption reaction in the above reactions.
At equilibrium, the rate of adsorption equals the rate of desorption. Setting rad = rd and rearranging, we obtain p A = k ad k d = K eq A; the concentration of sites is given by dividing the total number of sites by the area of the adsorbate: = S 0 / a. We can calculate the concentration of all sites by summing the concentration of free sites and occupied sites: = +. Combining this with the equilibrium equation, we get = K eq A
An exothermic reaction is a chemical reaction that releases energy through light or heat. It is the opposite of an endothermic reaction. Expressed in a chemical equation: reactants → products + energy. Exothermic Reaction means "thermic" means heat. So the reaction in which there is release of heat with or without light is called exothermic reaction. An exothermic reaction is a chemical reaction, it gives net energy to its surroundings. That is, the energy needed to initiate the reaction is less than the energy released; when the medium in which the reaction is taking place collects heat, the reaction is exothermic. When using a calorimeter, the total amount of heat that flows into the calorimeter is the negative of the net change in energy of the system; the absolute amount of energy in a chemical system is difficult to calculate. The enthalpy change, ΔH, of a chemical reaction is much easier to work with; the enthalpy change equals the change in internal energy of the system plus the work needed to change the volume of the system against constant ambient pressure.
A bomb calorimeter is suitable for measuring the energy change, ΔH, of a combustion reaction. Measured and calculated ΔH values are related to bond energies by: ΔH = − In an exothermic reaction, by definition, the enthalpy change has a negative value: ΔH < 0since a larger value is subtracted from a smaller value. For example, when hydrogen burns: 2H2 + O2 → 2H2O ΔH = −483.6 kJ/mol of O2 In an adiabatic system, the temperature raise due to enthalpy change can be expressed as −ΔH298.15 K = ∫T1T0Cp, pdT + ∫T0298 KdTwhere ΔH298.15 K is the standard enthalpy of reaction at 298 K, T0 and T1 are the initial and final temperature of the system and Cp,p and Cp,r are the heat capacities of the product and reactant, respectively. Assuming the heat capacity of the system remains as a constant value Cp,p=Cp,r=Cp, the change of temperature ΔT=T1−T0 can be expressed as −ΔH298.15 K = ∫T0+ΔTT0Cp, pdT = ΔTCp, pThe most available hand warmers make use of the oxidation of iron to achieve an exothermic reaction: 4Fe + 3O2 → 2Fe2O3 .
Combustion reactions of fuels or a substance e.g. Burning of natural gas: CH 4 + 2 O 2 ⟶ CO 2 + 2 H 2 O C 6 H 12 O 6 + 6 O 2 ⟶ 6 CO 2 + 6 H 2 O Neutralization The thermite reaction Reactions taking place in a self-heating can based on lime aluminium Many corrosion reactions such as oxidation of metals Most polymerization reactions The Haber process of ammonia production Respiration Decomposition of vegetable matter into compost Solution of sulfuric acid into water Dehydration of sugars upon contact with sulfuric acid Detonation of nitroglycerin Nuclear fission of uranium-235 The concept and its opposite number endothermic relate to the enthalpy change in any process, not just chemical reactions. In endergonic reactions and exergonic reactions it is the sign of the Gibbs free energy that determines the equilibrium point, not enthalpy; the related concepts endergonic and exergonic apply to all physical processes. The conceptually related endotherm and ectotherm are concepts in animal physiology.
In quantum numbers, when any excited energy level goes down to its original level for example: when n=4 fall to n=2, energy is released so, it is exothermic. Where an exothermic reaction causes heating of the reaction vessel, not controlled, the rate of reaction can increase, in turn causing heat to be evolved more quickly; this positive feedback situation is known as thermal runaway. An explosion can result from the problem. Heat production or absorption in either a physical process or chemical reaction is measured using calorimetry. One common laboratory instrument is the reaction calorimeter, where the heat flow into or from the reaction vessel is monitored; the technique can be used to follow chemical reactions as well as physical processes such as crystallization and dissolution. Energy released is measured in Joule per mole; the reaction has a negative ΔH value due to heat loss. E.g.: -123 J/mol Chemical thermodynamics Differential scanning calorimetry Endergonic Exergonic Endergonic reaction Exergonic reaction Exothermic process Endothermic reaction Endotherm
Gibbs free energy
In thermodynamics, the Gibbs free energy is a thermodynamic potential that can be used to calculate the maximum of reversible work that may be performed by a thermodynamic system at a constant temperature and pressure. The Gibbs free energy is the maximum amount of non-expansion work that can be extracted from a thermodynamically closed system; when a system transforms reversibly from an initial state to a final state, the decrease in Gibbs free energy equals the work done by the system to its surroundings, minus the work of the pressure forces. The Gibbs energy is the thermodynamic potential, minimized when a system reaches chemical equilibrium at constant pressure and temperature, its derivative with respect to the reaction coordinate of the system vanishes at the equilibrium point. As such, a reduction in G is a necessary condition for the spontaneity of processes at constant pressure and temperature; the Gibbs free energy called available energy, was developed in the 1870s by the American scientist Josiah Willard Gibbs.
In 1873, Gibbs described this "available energy" as the greatest amount of mechanical work which can be obtained from a given quantity of a certain substance in a given initial state, without increasing its total volume or allowing heat to pass to or from external bodies, except such as at the close of the processes are left in their initial condition. The initial state of the body, according to Gibbs, is supposed to be such that "the body can be made to pass from it to states of dissipated energy by reversible processes". In his 1876 magnum opus On the Equilibrium of Heterogeneous Substances, a graphical analysis of multi-phase chemical systems, he engaged his thoughts on chemical free energy in full. According to the second law of thermodynamics, for systems reacting at STP, there is a general natural tendency to achieve a minimum of the Gibbs free energy. A quantitative measure of the favorability of a given reaction at constant temperature and pressure is the change ΔG in Gibbs free energy, caused by the reaction.
As a necessary condition for the reaction to occur at constant temperature and pressure, ΔG must be smaller than the non-PV work, equal to zero. ΔG equals the maximum amount of non-PV work that can be performed as a result of the chemical reaction for the case of reversible process. If the analysis indicated a positive ΔG for the reaction energy — in the form of electrical or other non-PV work — would have to be added to the reacting system for ΔG to be smaller than the non-PV work and make it possible for the reaction to occur. We can think of ∆G as the amount of "free" or "useful" energy available to do work; the equation can be seen from the perspective of the system taken together with its surroundings. First, assume that the given reaction at constant temperature and pressure is the only one, occurring; the entropy released or absorbed by the system equals the entropy that the environment must absorb or release, respectively. The reaction will only be allowed if the total entropy change of the universe is positive.
This is reflected in a negative ΔG, the reaction is called exergonic. If we couple reactions an otherwise endergonic chemical reaction can be made to happen; the input of heat into an inherently endergonic reaction, such as the elimination of cyclohexanol to cyclohexene, can be seen as coupling an unfavourable reaction to a favourable one such that the total entropy change of the universe is greater than or equal to zero, making the total Gibbs free energy difference of the coupled reactions negative. In traditional use, the term "free" was included in "Gibbs free energy" to mean "available in the form of useful work"; the characterization becomes more precise if we add the qualification that it is the energy available for non-volume work.. However, an increasing number of books and journal articles do not include the attachment "free", referring to G as "Gibbs energy"; this is the result of a 1988 IUPAC meeting to set unified terminologies for the international scientific community, in which the adjective "free" was banished.
This standard, has not yet been universally adopted. The quantity called "free energy" is a more advanced and accurate replacement for the outdated term affinity, used by chemists in the earlier years of physical chemistry to describe the force that caused chemical reactions. In 1873, Willard Gibbs published A Method of Geometrical Representation of the Thermodynamic Properties of Substances by Means of Surfaces, in which he sketched the principles of his new equation, able to predict or estimate the tendencies of various natural processes to ensue when bodies or systems are brought into contact. By studying the interactions of homogeneous substances in contact, i.e. bodies composed of part solid, part liquid, part vapor, by using a three-dimensional volume-entropy-internal energy graph, Gibbs was able to determine three states of equilibrium, i.e. "necessarily stable", "neutral", "unstable", whether or not changes woul
In Newtonian mechanics, linear momentum, translational momentum, or momentum is the product of the mass and velocity of an object. It is a vector quantity, possessing a direction in three-dimensional space. If m is an object's mass and v is the velocity the momentum is p = m v, In SI units, it is measured in kilogram meters per second. Newton's second law of motion states that a body's rate of change in momentum is equal to the net force acting on it. Momentum depends on the frame of reference, but in any inertial frame it is a conserved quantity, meaning that if a closed system is not affected by external forces, its total linear momentum does not change. Momentum is conserved in special relativity and, in a modified form, in electrodynamics, quantum mechanics, quantum field theory, general relativity, it is an expression of one of the fundamental symmetries of time: translational symmetry. Advanced formulations of classical mechanics and Hamiltonian mechanics, allow one to choose coordinate systems that incorporate symmetries and constraints.
In these systems the conserved quantity is generalized momentum, in general this is different from the kinetic momentum defined above. The concept of generalized momentum is carried over into quantum mechanics, where it becomes an operator on a wave function; the momentum and position operators are related by the Heisenberg uncertainty principle. In continuous systems such as electromagnetic fields and deformable bodies, a momentum density can be defined, a continuum version of the conservation of momentum leads to equations such as the Navier–Stokes equations for fluids or the Cauchy momentum equation for deformable solids or fluids. Momentum is a vector quantity: it has both magnitude and direction. Since momentum has a direction, it can be used to predict the resulting direction and speed of motion of objects after they collide. Below, the basic properties of momentum are described in one dimension; the vector equations are identical to the scalar equations. The momentum of a particle is conventionally represented by the letter p.
It is the product of two quantities, the particle's mass and its velocity: p = m v. The unit of momentum is the product of the units of velocity. In SI units, if the mass is in kilograms and the velocity is in meters per second the momentum is in kilogram meters per second. In cgs units, if the mass is in grams and the velocity in centimeters per second the momentum is in gram centimeters per second. Being a vector, momentum has direction. For example, a 1 kg model airplane, traveling due north at 1 m/s in straight and level flight, has a momentum of 1 kg⋅m/s due north measured with reference to the ground; the momentum of a system of particles is the vector sum of their momenta. If two particles have respective masses m1 and m2, velocities v1 and v2, the total momentum is p = p 1 + p 2 = m 1 v 1 + m 2 v 2; the momenta of more than two particles can be added more with the following: p = ∑ i m i v i. A system of particles has a center of mass, a point determined by the weighted sum of their positions: r cm = m 1 r 1 + m 2 r 2 + ⋯ m 1 + m 2 + ⋯ = ∑ i m i r i ∑ i m i.
If all the particles are moving, the center of mass will be moving as well. If the center of mass is moving at velocity vcm, the momentum is: p = m v cm; this is known as Euler's first law. If the net force applied to a particle is a constant F, is applied for a time interval Δt, the momentum of the particle changes by an amount Δ p = F Δ t. In differential form, this is Newton's second law. If the net force experienced by a particle changes as a function of time, F, the change in momentum between times t1 and t2 is Δ p = J = ∫ t 1
Physisorption called physical adsorption, is a process in which the electronic structure of the atom or molecule is perturbed upon adsorption. The fundamental interacting force of physisorption is caused by van der Waals force. Though the interaction energy is weak, physisorption plays an important role in nature. For instance, the van der Waals attraction between surfaces and foot-hairs of geckos provides the remarkable ability to climb up vertical walls. Van der Waals forces originate from the interactions between induced, permanent or transient electric dipoles. In comparison with chemisorption, in which the electronic structure of bonding atoms or molecules is changed and covalent or ionic bonds form, physisorption speaking, can only be observed in the environment of low temperature and the absence of the strong chemisorptions. In practice, the categorisation of a particular adsorption as physisorption or chemisorption depends principally on the binding energy of the adsorbate to the substrate.
To give a simple illustration of physisorption, we can first consider an adsorbed hydrogen atom in front of a perfect conductor, as shown in Fig. 1. A nucleus with positive charge is located at R =, the position coordinate of its electron, r = is given with respect to the nucleus; the adsorption process can be viewed as the interaction between this hydrogen atom and its image charges of both the nucleus and electron in the conductor. As a result, the total electrostatic energy is the sum of attraction and repulsion terms: V = e 2 4 π ε 0; the first term is the attractive interaction of nucleus and its image charge, the second term is due to the interaction of the electron and its image charge. The repulsive interaction is shown in the third and fourth terms arising from the interaction of nucleus-image electron and electron-image nucleus, respectively. By Taylor expansion in powers of |r| / |R|, this interaction energy can be further expressed as: V = − e 2 16 π ε 0 Z 3 + 3 e 2 32 π ε 0 Z 4 + O.
One can find from the first non-vanishing term that the physisorption potential depends on the distance Z between adsorbed atom and surface as Z−3, in contrast with the r−6 dependence of the molecular van der Waals potential, where r is the distance between two dipoles. The van der Waals binding energy can be analyzed by another simple physical picture: modeling the motion of an electron around its nucleus by a three-dimensional simple harmonic oscillator with a potential energy Va: V a = m e 2 ω 2, where me and ω are the mass and vibrational frequency of the electron, respectively; as this atom approaches the surface of a metal and forms adsorption, this potential energy Va will be modified due to the image charges by additional potential terms which are quadratic in the displacements: V a = m e 2 ω 2 − e 2 16 π ε 0 Z 3 ( x 2 + y 2 2 + z
Corrosion is a natural process, which converts a refined metal to a more chemically-stable form, such as its oxide, hydroxide, or sulfide. It is the gradual destruction of materials by chemical and/or electrochemical reaction with their environment. Corrosion engineering is the field dedicated to stopping corrosion. In the most common use of the word, this means electrochemical oxidation of metal in reaction with an oxidant such as oxygen or sulfates. Rusting, the formation of iron oxides, is a well-known example of electrochemical corrosion; this type of damage produces oxide or salt of the original metal, results in a distinctive orange colouration. Corrosion can occur in materials other than metals, such as ceramics or polymers, although in this context, the term "degradation" is more common. Corrosion degrades the useful properties of materials and structures including strength and permeability to liquids and gases. Many structural alloys corrode from exposure to moisture in air, but the process can be affected by exposure to certain substances.
Corrosion can be concentrated locally to form a pit or crack, or it can extend across a wide area more or less uniformly corroding the surface. Because corrosion is a diffusion-controlled process, it occurs on exposed surfaces; as a result, methods to reduce the activity of the exposed surface, such as passivation and chromate conversion, can increase a material's corrosion resistance. However, some corrosion mechanisms are less predictable. Galvanic corrosion occurs when two different metals have physical or electrical contact with each other and are immersed in a common electrolyte, or when the same metal is exposed to electrolyte with different concentrations. In a galvanic couple, the more active metal corrodes at an accelerated rate and the more noble metal corrodes at a slower rate; when immersed separately, each metal corrodes at its own rate. What type of metal to use is determined by following the galvanic series. For example, zinc is used as a sacrificial anode for steel structures. Galvanic corrosion is of major interest to the marine industry and anywhere water contacts pipes or metal structures.
Factors such as relative size of anode, types of metal, operating conditions affect galvanic corrosion. The surface area ratio of the anode and cathode directly affects the corrosion rates of the materials. Galvanic corrosion is prevented by the use of sacrificial anodes. In any given environment, one metal will be either more noble or more active than others, based on how its ions are bound to the surface. Two metals in electrical contact share the same electrons, so that the "tug-of-war" at each surface is analogous to competition for free electrons between the two materials. Using the electrolyte as a host for the flow of ions in the same direction, the noble metal will take electrons from the active one; the resulting mass flow or electric current can be measured to establish a hierarchy of materials in the medium of interest. This hierarchy is useful in predicting and understanding corrosion, it is possible to chemically remove the products of corrosion. For example, phosphoric acid in the form of naval jelly is applied to ferrous tools or surfaces to remove rust.
Corrosion removal should not be confused with electropolishing, which removes some layers of the underlying metal to make a smooth surface. For example, phosphoric acid may be used to electropolish copper but it does this by removing copper, not the products of copper corrosion; some metals are more intrinsically resistant to corrosion than others. There are various ways of protecting metals from corrosion including painting, hot dip galvanizing, combinations of these; the materials most resistant to corrosion are those for which corrosion is thermodynamically unfavorable. Any corrosion products of gold or platinum tend to decompose spontaneously into pure metal, why these elements can be found in metallic form on Earth and have long been valued. More common "base" metals can only be protected by more temporary means; some metals have slow reaction kinetics though their corrosion is thermodynamically favorable. These include such metals as zinc and cadmium. While corrosion of these metals is continuous and ongoing, it happens at an acceptably slow rate.
An extreme example is graphite, which releases large amounts of energy upon oxidation, but has such slow kinetics that it is immune to electrochemical corrosion under normal conditions. Passivation refers to the spontaneous formation of an ultrathin film of corrosion products, known as a passive film, on the metal's surface that act as a barrier to further oxidation; the chemical composition and microstructure of a passive film are different from the underlying metal. Typical passive film thickness on aluminium, stainless steels, alloys is within 10 nanometers; the passive film is different from oxide layers that are formed upon heating and are in the micrometer thickness range – the passive film recovers if removed or damaged whereas the oxide layer does not. Passivation in natural environments such as air and soil at moderate pH is seen in such materials as aluminium, stainless steel and silicon. Passivation is determined by metallurgical and environmental factors; the effect of pH is summarized using Pourbaix diagrams.
Some conditions that inhibit passivation include high pH for aluminium and zinc, low pH or the p