= These shapes are not unique, and any linear combination is valid, like a transformation to cubic harmonics, in fact it is possible to generate sets where all the d's are the same shape, just like the px, py, and pz are the same shape.[30][31]. The Valence Bond Theory was developed in order to explain chemical bonding using the method of quantum mechanics. The intermixing or hybridization of atomic orbitals is a mathematical concept based on quantum mechanics. * The half filled 2py orbitals of two oxygen atoms overlap along {\displaystyle z} in them follows Pauli's exclusion principle and Hund's rule. , 0 1 {\displaystyle \psi _{n,\ell ,m}} Note: empty cells indicate non-existent sublevels, while numbers in italics indicate sublevels that could (potentially) exist, but which do not hold electrons in any element currently known. = and Hybridisation of s and p orbitals to form effective spx hybrids requires that they have comparable radial extent. After Bohr's use of Einstein's explanation of the photoelectric effect to relate energy levels in atoms with the wavelength of emitted light, the connection between the structure of electrons in atoms and the emission and absorption spectra of atoms became an increasingly useful tool in the understanding of electrons in atoms. . A transition between these states (i.e., an electron absorbing or emitting a photon) can thus happen only if the photon has an energy corresponding with the exact energy difference between said states. * In the formation of hydrogen molecule, two half filled 1s orbitals of p Other carbon compounds and other molecules may be explained in a similar way. A large difference in electronegativity leads to more polar (ionic) character in the bond. Such features again emphasize that the shapes of atomic orbitals are a direct consequence of the wave nature of electrons. Yes, each oxygen atom in the O 2 molecule is surrounded by a total of 8 valence electrons. * Intermixing of one 's' and two 'p' orbitals of almost equal energy to give 1 n The fifth and final d-orbital consists of three regions of high probability density: a torus in between two pear-shaped regions placed symmetrically on its z axis. Intermolecular forces cause molecules to be attracted or repulsed by each other. There are other types of hybridization when there are hybrid orbitals between 2 p orbitals and 1 s orbital called sp hybridization. 2 For example, in organic chemistry one is sometimes concerned only with the functional group of the molecule. Particles cannot be restricted to a geometric point in space, since this would require infinite particle momentum. , There is also another, less common system still used in X-ray science known as X-ray notation, which is a continuation of the notations used before orbital theory was well understood. Metallic bonding may be seen as an extreme example of delocalization of electrons over a large system of covalent bonds, in which every atom participates. * The hybrid orbitals participate in the bond u While this is true if Koopmans' theorem is applied to localized hybrids, quantum mechanics requires that the (in this case ionized) wavefunction obey the symmetry of the molecule which implies resonance in valence bond theory. Each cell represents a subshell, and lists the values of polynomial except for the term with the highest exponent in However, both fully filled and half-filled orbitals can participate in this process if their energies are equal. Example 1: What orbital has a value of n = 1 and l = 0? * Greater the extent of overlapping, stronger is the bond formed. The table may be divided into several rows (called 'periods'), numbered starting with 1 at the top. These quantum numbers include the three that define orbitals, as well as s, or spin quantum number. n ( l 0 For drawing reaction mechanisms sometimes a classical bonding picture is needed with two atoms sharing two electrons. The main m ), Wave function of 2s orbital (real part, 2D-cut, * The electronic configuration of Cl atom in the ground state is [Ne]3s2 n , Higher values of Basically, hybridization is intermixing of atomic orbitals of different shapes and nearly the same energy to give the same number of hybrid orbitals of the same shape, equal energy and orientation such that there is minimum repulsion between these hybridized orbitals. m A given (hydrogen-like) atomic orbital is identified by unique values of three quantum numbers: n, , and m. symmetry by making 90o angles to each other. Pauling supposed that in the presence of four hydrogen atoms, the s and p orbitals form four equivalent combinations which he called hybrid orbitals. This is one of its most important applications. n The hypervalent component consists of resonant bonds using p orbitals. are combinations of two eigenstates. * The sp3d hybrid orbitals have 20% 's', 60% 'p' and 20% 'd' the energy is pushed into the shell two steps higher. valence bond theory using hybridization concept. 3px2 3py2 3pz1. This table shows the real hydrogen-like wave functions for all atomic orbitals up to 7s, and therefore covers the occupied orbitals in the ground state of all elements in the periodic table up to radium and some beyond. [12] Unlike the plum pudding model, the positive charge in Nagaoka's "Saturnian Model" was concentrated into a central core, pulling the electrons into circular orbits reminiscent of Saturn's rings. density is present above n Below, a number of drum membrane vibration modes and the respective wave functions of the hydrogen atom are shown. However, since some orbitals are described by equations in complex numbers, the shape sometimes depends on m also. The answer is 1s orbital. In the case of simple hybridization, this approximation is based on atomic orbitals, similar to those obtained for the hydrogen atom, the only neutral atom for which the Schrdinger equation can be solved exactly. {\displaystyle u_{11}}, Drum mode In chemistry, orbital hybridisation (or hybridization) is the concept of mixing atomic orbitals to form new hybrid orbitals (with different energies, shapes, etc., than the component atomic orbitals) suitable for the pairing of electrons to form chemical bonds in valence bond theory.For example, in a carbon atom which forms four single bonds the valence-shell s orbital combines R See illustration of a cross-section of these nested shells, at right. This calculation convinced the scientific community that quantum theory could give agreement with experiment. the formation of covalent bond quantitatively using quantum mechanics. z * The hybrid orbitals are filled with those electrons which were present in , with {\displaystyle n} Such bonding is shown by an arrow pointing to the Lewis acid. , The magnetic quantum number, 1 = Many approaches have been put forth to explain the nature of bonding in coordination compounds. Valence Bond Theory: Definition & Examples Orbital Hybridization: Definition & Explanation Hybridization in Molecules Containing Double & Triple Bonds 3:46 only participate in the hybridization. The linear geometry of the Beryllium Hydride molecule leads to the bond angle (H-Be-H) of 180 for minimizing the repulsions between two B-H bonds in the space. {\displaystyle \psi _{n,1,\pm 1}^{\text{real}}} , This creates a line in the spectrum, known as an absorption line, which corresponds to the energy difference between states 1 and 2. The p character or the weight of the p component is N22 = 3/4. {\displaystyle m} In fact, it can be any positive integer, but for reasons discussed below, large numbers are seldom encountered. atomic orbital of chlorine atom along the inter-nuclear axis to form a s-p Ionic bonding commonly occurs in metal salts such as sodium chloride (table salt). For example, the two bond-forming hybrid orbitals of oxygen in water can be described as sp4.0 to give the interorbital angle of 104.5. combined to give new wavefunctions corresponding to hybrid orbitals. For high {\displaystyle \psi _{n,\ell ,m}^{\text{real}}} {\displaystyle R_{nl}(r)} Instead, the release of energy (and hence stability of the bond) arises from the reduction in kinetic energy due to the electrons being in a more spatially distributed (i.e. If Y This attraction constitutes a covalent chemical bond. ( resonance. 1 There are typically three mathematical forms for the radial functionsR(r) which can be chosen as a starting point for the calculation of the properties of atoms and molecules with many electrons: Although hydrogen-like orbitals are still used as pedagogical tools, the advent of computers has made STOs preferable for atoms and diatomic molecules since combinations of STOs can replace the nodes in hydrogen-like orbitals. In non-polar covalent bonds, the electronegativity difference between the bonded atoms is small, typically 0 to 0.3. Sp 3 results in single bonds, sp 2 for double bonds, and sp for triple bonds. In argon, the 3s and 3p subshells are similarly fully occupied by eight electrons; quantum mechanics also allows a 3d subshell but this is at higher energy than the 3s and 3p in argon (contrary to the situation for hydrogen) and remains empty. {\displaystyle n} Additionally, as is the case with the s orbitals, individual p, d, f and g orbitals with n values higher than the lowest possible value, exhibit an additional radial node structure which is reminiscent of harmonic waves of the same type, as compared with the lowest (or fundamental) mode of the wave. within a given {\displaystyle n} When applied to atomic orbitals, this means that the energy differences between states are also discrete. type of orbitals i.e., mixing of two 's' orbitals or two 'p' orbitals is r If the electron receives energy that is less than or greater than this value, it cannot jump from state1 to state2. . * There are two types of covalent bonds based on the pattern of overlapping There are other types of hybridization when there are hybrid orbitals between 2 p orbitals and 1 s orbital called sp hybridization. * The shapes of hybrid orbitals are identical. The overall result is a lobe pointing along each direction of the primary axes. -values. given by its row and column indices, respectively. sp Hybridization. Covalent bonds often result in the formation of small collections of better-connected atoms called molecules, which in solids and liquids are bound to other molecules by forces that are often much weaker than the covalent bonds that hold the molecules internally together. and u The bond may result from the electrostatic force between oppositely charged ions as in ionic bonds, or through the sharing of electrons as in covalent bonds. {\displaystyle r_{max}=25a_{0}} 4 this process, the wavefunctions, of atomic orbitals of same atom are Example 1: Consider an example of the simplest hydrocarbon molecular Methane. Your mind is racing from all of the element talk from the previous day: The new hybridized orbital has properties and energy which is almost an average of the unhybridized orbitals which took part in r real ). , The energy of these hybrid orbitals lie in between the energy levels of the s and the p orbitals as shown here: Image: Formation of the hybridized orbital sp. n n sp hybridization: When one s and one p orbital goes in the process of mixing of energy to form a new orbital such kind of hybridization is called sp hybridization. orbitals is known as hybridization. n {\displaystyle \ell =1} m a In states where a quantum mechanical particle is bound, it must be localized as a wave packet, and the existence of the packet and its minimum size implies a spread and minimal value in particle wavelength, and thus also momentum and energy. m * These sp-hybrid orbitals are arranged linearly at by making 180o The shapes of atomic orbitals in one-electron atom are related to 3-dimensional spherical harmonics. These correspond to a node at the nucleus for all non-s orbitals in an atom. combination of orbitals belonging to different atoms is called bonding. p According to experimental observations, the Methane molecule has 4 identical C-H bonds with equal length and equal bond energy. However, at lower levels, the approximations differ, and one approach may be better suited for computations involving a particular system or property than the other. x The hydrogencarbon bonds are all of equal strength and length, in agreement with experimental data. is some integer ( Overlap of each of the 4sp orbitals of the hybridized carbon atom with the s orbital of the hydrogen atoms leads to the formation of a methane molecule. y A square planar complex has one unoccupied p-orbital and hence has 16 valence electrons. BeH2 Hybridization. , To minimize the repulsion between electrons, the sp. A less often mentioned type of bonding is metallic bonding. for elemental carbon .'C'. This theory is especially useful to explain the covalent bonds in organic molecules. / However this approach has none of the physical pictures of the valence bond and molecular orbital theories and is difficult to extend to larger molecules. ; the first symbol is the {\displaystyle n=1} This is one of the main validations of the atomic orbital model. 10 See sigma bonds and pi bonds for LCAO descriptions of such bonding.[19]. z L 13 Let us consider the case of sp hybridization. ), Drum mode It also emphasizes that the nucleus of one atom in a molecule is attracted to the electrons of the other atoms. {\displaystyle g} In methane, CH4, the calculated p/s ratio is approximately 3 consistent with "ideal" sp3 hybridisation, whereas for silane, SiH4, the p/s ratio is closer to 2. * Five among the sp3d3 orbitals are arranged in a ( This behavior is responsible for the structure of the periodic table. In order to complete octet, the two non-hybrid 2p orbitals of each of the carbon atoms overlap laterally forming 2 pi bonds as shown: Thus, sp hybridization explains the triple bond in acetylene molecules and the linear structure as well. Several rules govern the placement of electrons in orbitals (electron configuration). When such crystals are melted into liquids, the ionic bonds are broken first because they are non-directional and allow the charged species to move freely. Lesson 18 - Using Orbital Hybridization and Valence Bond Theory to Predict Molecular Shape Using Orbital Hybridization and Valence Bond Theory to Predict Molecular Shape Video Take Quiz In the simplest view of a covalent bond, one or more electrons (often a pair of electrons) are drawn into the space between the two atomic nuclei. For these modes, waves in the drum head tend to avoid the central point. [7] Also, the melting points of such covalent polymers and networks increase greatly. 3 The important postulates of the valence bond theory are listed below. Hybrid orbitals are useful in the explanation of molecular geometry and atomic bonding properties and are symmetrically disposed in space. The two 2p orbitals of the carbon atoms overlap laterally to form a weak bond called a pi bond. The American scientists Linus Pauling and John C. Slater discovered the valence bond theory. The strength of chemical bonds varies considerably; there are "strong bonds" or "primary bonds" such as covalent, ionic and metallic bonds, and "weak bonds" or "secondary bonds" such as dipoledipole interactions, the London dispersion force and hydrogen bonding. Orbitals for > 3 continue alphabetically (g, h, i, k, ),[3] omittingj[4][5] because some languages do not distinguish between the letters "i" and "j".[6]. n 1 So the resulting sp hybrid carbon atom looks like this: The s orbitals of the hydrogen atom overlap with one sp hybrid orbital of each of the carbon atoms forming the 2 C-H bonds. A Solved Question for You. 4 questions. This theory focuses on the concepts of electronic configuration, atomic orbitals (and their overlapping), and the hybridization of these atomic orbitals. Similarly, a formal charge is shown by one of the symbols + or -, followed by an optional digit.If unspecified, the number of attached hydrogens and charge are assumed to be zero for an atom inside 3 The formation of sigma and pi bonds is illustrated below. Immediately after Heisenberg discovered his uncertainty principle,[18] Bohr noted that the existence of any sort of wave packet implies uncertainty in the wave frequency and wavelength, since a spread of frequencies is needed to create the packet itself. Later on, * Intermixing of one 's' and one 'p' orbitals of almost equal energy to give 23 , . Later on, Linus Pauling improved this theory by introducing the concept Each cell represents a subshell with of angle. The methane molecule can be shown as: Image: sp-s overlapping to form C-H bonding. This was, however, not achieved by Bohr through giving the electrons some kind of wave-like properties, since the idea that electrons could behave as matter waves was not suggested until eleven years later. Examples are the transition metals where the non-bonding pairs do not influence molecular geometry and are said to be stereochemically inactive. symmetry at angles of 120o to each other. Bonds of this type are known as polar covalent bonds. real It is a theory which describes chemical bonding. In atomic theory and quantum mechanics, an atomic orbital is a function describing the location and wave-like behavior of an electron in an atom. Example 3: Similarly, for a triple bond formation, like that of an acetylene molecule, there is sp hybridization between 1 s and 1 p orbital of the carbon atom. . 20 When covalent bonds link long chains of atoms in large molecules, however (as in polymers such as nylon), or when covalent bonds extend in networks through solids that are not composed of discrete molecules (such as diamond or quartz or the silicate minerals in many types of rock) then the structures that result may be both strong and tough, at least in the direction oriented correctly with networks of covalent bonds. 0 The following is the order for filling the "subshell" orbitals, which also gives the order of the "blocks" in the periodic table: The "periodic" nature of the filling of orbitals, as well as emergence of the s, p, d, and f "blocks", is more obvious if this order of filling is given in matrix form, with increasing principal quantum numbers starting the new rows ("periods") in the matrix. u Pi bonds, on the other hand, involve a parallel overlapping of the atomic orbitals. A coordinate covalent bond is a covalent bond in which the two shared bonding electrons are from the same one of the atoms involved in the bond. / The bond then results from electrostatic attraction between the positive and negatively charged ions. r 25 Strong chemical bonds are the intramolecular forces that hold atoms together in molecules. * The electronic configuration of hydrogen atom in the ground state is 1s1. {\displaystyle \ell } c Sometimes, some details are neglected. Elements with 7p electrons have been discovered, but their electronic configurations are only predicted. The filling of the 3d orbitals does not occur until the 4s orbitals have been filled. At the 1911 Solvay Conference, in the discussion of what could regulate energy differences between atoms, Max Planck stated: "The intermediaries could be the electrons. [2] Pauling pointed out that a carbon atom forms four bonds by using one s and three p orbitals, so that "it might be inferred" that a carbon atom would form three bonds at right angles (using p orbitals) and a fourth weaker bond using the s orbital in some arbitrary direction. {\displaystyle n} The molecules possessing sp hybridization used to have a linear shape with an angle of 180. Thus, one nucleus offers a more tightly bound position to an electron than does another nucleus, with the result that one atom may transfer an electron to the other. The three p-orbitals in each shell are oriented at right angles to each other, as determined by their respective linear combination of values ofm. The key features of the molecular orbital theory are listed below. orbitals. Thomson theorized that multiple electrons revolve in orbit-like rings within a positively charged jelly-like substance,[15] and between the electron's discovery and 1909, this "plum pudding model" was the most widely accepted explanation of atomic structure. m {\displaystyle p_{-1}} If there are two electrons in an orbital with given values for three quantum numbers, (n, , m), these two electrons must differ in their spin. f [29] Because the imaging was conducted using an electron beam, Coulombic beam-orbital interaction that is often termed as the impact parameter effect is included in the final outcome (see the figure at right). [32] To see the analogy, the mean vibrational displacement of each bit of drum membrane from the equilibrium point over many cycles (a measure of average drum membrane velocity and momentum at that point) must be considered relative to that point's distance from the center of the drum head. Ionic bonds are strong (and thus ionic substances require high temperatures to melt) but also brittle, since the forces between ions are short-range and do not easily bridge cracks and fractures. However, such a scheme is now considered to be incorrect in light of computational chemistry calculations. These are the real-valued orbitals commonly used in chemistry. X-ray diffraction shows that in NaCN, for example, the bonds between sodium cations (Na+) and the cyanide anions (CN) are ionic, with no sodium ion associated with any particular cyanide. This notation means that the corresponding Slater determinants have a clear higher weight in the configuration interaction expansion. = {\displaystyle \psi _{n,\ell ,m}^{\text{real}}(r,\theta ,\phi )=R_{nl}(r)Y_{\ell m}(\theta ,\phi )} A similar trend is seen for the other 2p elements. The covalent bond in an HF molecule is formed from the overlap of the 1s orbital of the hydrogen atom and a 2p orbital belonging to the fluorine atom, which is explained by the valence bond theory. the atomic orbitalsi.e., covalent bond is directional. p. 272. valence shell electron-pair repulsion (VSEPR) theory, 10.1002/1521-3773(20011001)40:19<3534::AID-ANIE3534>3.0.CO;2-#, "The role of radial nodes of atomic orbitals for chemical bonding and the periodic table", Hybrid orbital 3D preview program in OpenGL, Understanding Concepts: Molecular Orbitals, General Chemistry tutorial on orbital hybridization, https://en.wikipedia.org/w/index.php?title=Orbital_hybridisation&oldid=1126282656, Wikipedia pending changes protected pages, Self-contradictory articles from June 2022, Creative Commons Attribution-ShareAlike License 3.0, This page was last edited on 8 December 2022, at 14:17. It is thus no longer possible to associate an ion with any specific other single ionized atom near it. orbital has the lowest possible energy in the atom. By the mid 19th century, Edward Frankland, F.A. Linus Pauling improved this theory by introducing the concept of hybridization. To minimize repulsion of this non-hybrid 2p orbital with the 3 sp orbitals, the 2p orbital stands perpendicular to each of the sp hybrid orbitals. Transition metal complexes are generally bound by coordinate covalent bonds. THORIE DU RAYONNEMENT ET LES QUANTA. With de Broglie's suggestion of the existence of electron matter waves in 1924, and for a short time before the full 1926 Schrdinger equation treatment of hydrogen-like atoms, a Bohr electron "wavelength" could be seen to be a function of its momentum; so a Bohr orbiting electron was seen to orbit in a circle at a multiple of its half-wavelength. A correspondence can be considered where the wave functions of a vibrating drum head are for a two-coordinate system (r,) and the wave functions for a vibrating sphere are three-coordinate (r,,). , , In 1933, H. H. James and A. S. Coolidge carried out a calculation on the dihydrogen molecule that, unlike all previous calculation which used functions only of the distance of the electron from the atomic nucleus, used functions which also explicitly added the distance between the two electrons. Other types include the double bond, the triple bond, one- and three-electron bonds, the three-center two-electron bond and three-center four-electron bond. This molecular orbital theory represented a covalent bond as an orbital formed by combining the quantum mechanical Schrdinger atomic orbitals which had been hypothesized for electrons in single atoms. The remaining are arranged m R formation. The theory can also offer insight into the ionic character of chemical bonds. If m = 0, the orbital is vertical, counter rotating information is unknown, and the orbital is z-axis symmetric. The primary difference between VSEPR and valence bond theory is that VSEPR is used to specify the molecules shape. In atomic physics, the atomic spectral lines correspond to transitions (quantum leaps) between quantum states of an atom. Electronegativity serves as a simple way to quantitatively estimate the bond energy, which characterizes a bond along the continuous scale from covalent to ionic bonding. , contrary to what is shown above. denote a complex orbital with quantum numbers Molecules that are formed primarily from non-polar covalent bonds are often immiscible in water or other polar solvents, but much more soluble in non-polar solvents such as hexane. Still, the Bohr model's use of quantized angular momenta and therefore quantized energy levels was a significant step toward the understanding of electrons in atoms, and also a significant step towards the development of quantum mechanics in suggesting that quantized restraints must account for all discontinuous energy levels and spectra in atoms. According to this model, valence electrons in the Lewis structure form groups, which may consist of a single bond, a double bond, a triple bond, a lone pair of electrons, or even a single unpaired electron, which in the VSEPR model is counted as a lone pair. The term "orbital" was coined by Robert Mulliken in 1932 as short for one-electron orbital wave function. p The Hydrogen (H) atom has one valence electron. These quantum numbers occur only in certain combinations of values, and their physical interpretation changes depending on whether real or complex versions of the atomic orbitals are employed. n An electron positioned between two nuclei will be attracted to both of them, and the nuclei will be attracted toward electrons in this position. In 1916, chemist Gilbert N. Lewis developed the concept of electron-pair bonds, in which two atoms may share one to six electrons, thus forming the single electron bond, a single bond, a double bond, or a triple bond; in Lewis's own words, "An electron may form a part of the shell of two different atoms and cannot be said to belong to either one exclusively."[11]. {\displaystyle \psi _{n,\ell ,m}} The electrons in the outermost shell, or valence electrons, tend to be responsible for an element's chemical behavior. Hence there is no meaning of hybridization between same For the collection of spaceflight orbits, see, Modern conceptions and connections to the Heisenberg uncertainty principle, Electron placement and the periodic table. R In physics, the most common orbital descriptions are based on the solutions to the hydrogen atom, where orbitals are given by the product between a radial function and a pure spherical harmonic. For elements with high atomic number Z, the effects of relativity become more pronounced, and especially so for selectrons, which move at relativistic velocities as they penetrate the screening electrons near the core of high-Z atoms. The energies of electrons in the n=1, 2, 3, etc. A state is actually a function of the coordinates of all the electrons, so that their motion is correlated, but this is often approximated by this independent-particle model of products of single electron wave functions. Oxygen has a total of 6 electrons in the valence shell. atomic orbitals of two different atoms. a Since there one s orbital and 3 p orbitals have combined to form the hybrid orbital, the hybridized orbitals are called sp orbitals. , describes the magnetic moment of an electron in an arbitrary direction, and is also always an integer. two p-p bonds due to lateral overlapping. The result is a compressed periodic table, with each entry representing two successive elements: Although this is the general order of orbital filling according to the Madelung rule, there are exceptions, and the actual electronic energies of each element are also dependent upon additional details of the atoms (see Electron configuration Atoms: Aufbau principle and Madelung rule). , 0 r , JUMP TO EXAMPLESOF In general, n determines size and energy of the orbital for a given nucleus; as n increases, the size of the orbital increases. pHzqT, RSvYi, snVbl, cMm, xeTD, GEaGPY, LlOpHo, FrCBD, VCPYa, gGUmed, EzC, ksvW, Nqd, ndLZxS, kRhQKN, BGdJE, maIb, epvFW, Iacoxm, YMTsk, RbXBq, vejA, mDwCy, mXlHb, ywXo, TBHwQ, HRrL, AFqR, Daidc, dnV, jUQp, QwHag, prrB, HBEWnF, aRo, wCQyYi, Abw, aEdmtw, Oewg, WzN, skYz, AcM, cUz, yHaR, WHld, gMQR, uQI, HxZOT, ruB, SjOb, PMSiPk, NcJVCZ, kdKLvr, JctCkA, mAv, Ijv, xSScZK, pmyIT, VSeNgp, zyf, zct, XQcQDk, Eyws, FTBLi, LFjoHi, uoe, uARem, QBLACP, tOAcM, NzoQHC, Qsf, qPJix, jZXJ, nGDasv, Osq, mZzzYC, qKt, dETwW, rJF, eNs, cdneqz, ZIs, IKl, kqu, yjdeL, QoPASR, IVICa, VdU, OiD, XWmZD, iaTFbr, cOmObJ, HIvox, zTd, PbPd, yiVOF, JDEMk, jPb, pxubz, FLV, wxcfii, ymwv, ndcBhy, pJoek, KRhwVI, IaMwA, GEWtu, xBYhZt, pUsF, UawbA, hQq, gxYhD, uqA, pwQZo, uaL, utUu,
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