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Lewis Dot Structure Geometry

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Last Updated: 13 October 2020

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Representation of Collection of Molecules Collection of increasingly complex representations of molecules and their structures is shown here. This is the Lewis Structure of benzene. B. This is a picture of the structure of taxol as organic line structure, to give a better idea of its overall shape. C. This is a representation of enzyme catalase, where color shapes and spirals are a way of representing the 3 - dimensional structure of large biological molecules. And B: Science Media Group. C: Wikimedia Commons, Public Domain. Figure 5 - 2. Valence Electrons for First 26 Main Group Elements Only Electrons in the outer, valence shell are represent. These range from one electron for alkali metals to eight electrons for noble gases. All of these Electrons were found in outermost s and p orbitals for each of these elements. Note that in bottom row, atomic number jumps from 20 to 31. The Jag line indicates where transition metals would be locate. Science Media Group. Figure 5 - 5. The Formation of Ionic Bond electron configuration of sodium and chlorine makes it possible for them to form Ionic Bond. After transfer of electrons from sodium to chlorine, compound sodium chloride, common table salt, forms. Note that it is made up of Na + and Cl -, each of which has 8 valence Electrons between its outermost s and p orbitals. For each of these, atoms and ions are written as Lewis dot formulas. Science Media Group. Figure 5 - 7. Diatomic Covalent Molecules simplest covalently - bound Molecules are diatomic molecules; these happen between only two atoms sharing at least two electrons. See how the patterns of FIgure 5 - 6 are repeated here: Fluorine makes one Covalent bond, oxygen makes two Covalent bonds, and nitrogen makes three Covalent bonds. We can make more than one Covalent Bond to the same atom in order to share enough electrons to reach octet. For example, oxygen atom has two lone pairs and two bonds, each of which represents two Electrons and four lone pair Electrons. These, plus four covalently Bond Electrons equals eight Electrons to satisfy the octet rule. Science Media Group. Figure 5 - 8. Electron Clouds and Lewis Structure H2 and HCl Molecules are represented here, showing the electron density of their electron clouds as they covalently share Electrons. Note how very electronegative chlorine atom in HCl helps to tilt the cloud of electrons on the molecule towards its side, making the chlorine end a bit more negative than the hydrogen end, which ends up a bit more positive, and thus polar. By contrast, hydrogen molecule has a perfectly symmetrical cloud of two hydrogen atoms evenly sharing electrons, and the molecule is therefore non - polar. Science Media Group. Figure 5 - 11. Sulfuryl chloride sulfur in the center of the molecule actually makes six Covalent bonds, but that allows for sulfur, which is third - row element.

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Molecular Geometries

Tetra - signifies four, and - hedral relates to the face of a solid; tetrahedral literally means having four face. This shape is found when there are four bonds all on one central atom, with no lone electron pairs. In accordance with VSEPR theory, bond angles between electron bonds are 109. 5 O. An example of a tetrahedral molecule is methane. Four equivalent bonds point in four geometrically equivalent directions in three dimensions, corresponding to four corners of the tetrahedron center of the carbon atom. A Trigonal bipyramidal shape forms when the central atom is surrounded by five atoms of molecule. In geometry, three atoms are on the same plane with bond angles of 120; other two atoms are on opposite ends of the molecule. Some elements in Group 15 of periodic table form compounds of type AX 5; examples include PCl 5 and AsF 5.

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Section 6: Advanced Lewis Structures

We also use Lewis symbols to indicate formation of covalent bonds, which are shown in Lewis Structures, drawings that describe bonding in molecules and polyatomic ions. For example, when two chlorine atoms form chlorine molecule, they share one pair of electrons: Lewis structure indicates that each Cl atom has three pairs of electrons that are not used in bonding and one share pair of electrons. Dash is sometimes used to indicate shared pair of electrons: single shared pair of electrons is called a single bond. Each Cl atom interacts with eight valence electrons: six in lone pairs and two in single bond.


The Octet Rule

We will also encounter a few molecules that contain central atoms that do not have fill valence shell. Generally, these are molecules with central atoms from groups 2 and 12, outer atoms that are hydrogen, or other atoms that do not form multiple bonds. For example, in Lewis structures OF beryllium dihydride, BeH 2, and boron trifluoride, BF 3, beryllium and boron atoms each have only four and six electrons, respectively. It is possible to draw a structure with a double bond between boron atom and fluorine atom in BF 3, satisfying the octet rule, but experimental evidence indicates bond lengths are closer to that expected for B - F single bonds. This suggests the best Lewis structure has three B - F single bonds and electron deficient boron. Reactivity OF compound is also consistent with electron deficient boron. However, B - F bonds are slightly shorter than what is actually expected for B - F single bonds, indicating that some double bond characters are found in actual molecule. Atoms like boron atom in BF 3, which do not have eight electrons, are very reactive. It readily combines with molecule containing atom with a lone pair of electrons. For example, NH 3 reacts with BF 3 because lone pair of nitrogen can be shared with boron atom:

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions.

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions

Section 7: VSEPR Theory

1 Lewis Structures; Octet Rule 7. 2 Molecular Geometry Valence - Shell Electron - Pair Repulsion 7. 3 Polarity OF Molecules 7. 4 Atomic Orbitals; Hybridization 3 Steps in converting Molecular formula into Lewis Structure. Molecular formula Atom placement Sum OF Valence E - Remaining Valence E - Lewis Structure Place Atom with lowest EN in center Add - group numbers Draw single Bonds. Subtract 2e - For Each Bond. Give Each Atom 8e - Step 1 Step 2 Step 3 Step 4 4 Molecular formula Atom placement Sum OF Valence E - Remaining Valence E - Lewis Structure For NF 3 N FF F N 5e - F 7e - X 3 = 21e - Total 26e -: 5 SAMPLE PROBLEMWriting Lewis Structures For Molecules with One Central Atom SOLUTION: PROBLEM: Write Lewis Structure For CCl 2 F 2, One OF compounds responsible For depletion OF stratospheric ozone. Plan: Follow Steps outlined in Figure 10. 1 Step 1: Carbon has the lowest EN and is Central Atom. Other atoms are located around it. C Steps 2 - 4: C have 4 Valence E -, Cl and F each have 7. The sum is 4 + 4 = 32 Valence E -. Cl F F C F F Make Bonds and fill in Remaining Valence electrons placing 8e - around each atom.: 6 SAMPLE PROBLEMWriting Lewis Structure For Molecules with More Than One Central Atom PROBLEM: Write Lewis Structure For methanol, important industrial alcohol that is being used as gasoline alternative in car engines. Solution: Hydrogen can have only One Bond. So C and O must be next to each other with H filling in bonds. There are 4 + 4 + 6 = 14 Valence E -. C has 4 Bonds and O has 2. O has 2 Pair OF Nonbonding E -. Coh H H H: 7 SAMPLE PROBLEMWriting Lewis Structures For Molecules with Multiple Bonds. Plan: SOLUTION: PROBLEM: Write Lewis Structures For following: ethylene, most important reactant in manufacture of polymers, nitrogen, most abundant atmospheric gas for Molecules with Multiple Bonds, There is Step 5 which Follow other Steps in Lewis Structure construction. If Central Atom does not have 8e -, Octet, then two E - cans be moved in to form Multiple Bond.S There are 2 + 4 = 12 Valence E -. H can have only One Bond per Atom. Cc H HH H: CC H HH H N 2 has 2 = 10 Valence E -. Therefore, triple bond is required to make Octet around each N. N: N:. N: N:. N: N: 8 Resonance: delocalize Electron - Pair Bonding Resonance Structures have the same relative Atom placement but difference in locations of Bonding and Nonbonding Electron pairs. O 3 can be Draw in 2 ways - Neither structure is actually correct but can be Draw to represent a structure which is a hybrid OF Two - Resonance Structure. It is used to indicate that resonance occur. 9 SAMPLE PROBLEMWriting Resonance Structures PLAN: SOLUTION: PROBLEM: Write Resonance Structures For nitrate ion, NO 3 -.


Key Concepts and Summary

Lewis electron structures give no information about molecular geometry, arrangement of bond atoms in molecules or polyatomic ion, which is crucial to understanding chemistry of molecule. The Valence - shell electron - pair repulsion model allows us to predict which of possible structures is actually observed in most cases. It is based on the assumption that pairs of electrons occupy space, and the lowest - energy structure is one that minimizes electron pair - electron pair repulsions. In the VSEPR model, molecule or polyatomic ion is given AX m E n designation, where is central atom, X is bond atom, E is nonbonding valence electron group, and m and n are integers. Each group around central atom is designated as a bonding pair or lone pair. From BP and LP interactions we can predict both relative positions of atoms and angles between bonds, called bond angles. From this we can describe molecular geometry. The VSEPR model can be used to predict shapes of many molecules and polyatomic ions, but it gives no information about bond lengths and presence of multiple bonds. The combination of VSEPR and bonding model, such as Lewis electron structures, is necessary to understand the presence of multiple bonds. Molecules with polar covalent bonds can have dipole moment, asymmetrical distribution of charge that results in a tendency for molecules to align themselves in applied electric field. Any diatomic molecule with polar covalent bond has a dipole moment, but in polyatomic molecules, presence or absence of net dipole moment depends on structure. For some highly symmetrical structures, individual bond dipole moments cancel one another, giving dipole moment of zero.


Two Electron Groups

The central atom, carbon, contributes four valence electrons, and each oxygen atom contributes six. Lewis electron structure is 2. Carbon atoms form two double bonds. Each double bond is group, so there are two electron groups around the central atom. Like BeH 2, arrangement that minimizes repulsions places groups 180 apart. 3. Once again, both groups around central atom are bonding pairs, so CO 2 is designated as AX 2. 4. Vsepr only recognizes groups around central atom. Thus, lone pairs of oxygen atoms do not influence molecular geometry. With two bonding pairs on the central atom and no lone pairs, molecular geometry of CO 2 is linear. Figure: structure of {matheq}{CO2}{endmatheq} is shown in Figure:

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions.

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions

Section 8: Hybrid Orbitals

Lewis Structures provide NO information about shapes of molecules, being limited to which atoms are connected to each other and whether bonds are single, double or triple. More sophisticated treatment of bonding is needed to begin to understand shapes of molecules. In this section, we present a quantum mechanical description of Bonding, in which Bonding electrons are viewed as being localized between nuclei of Bond Atoms. Overlap of Bonding Orbitals is substantially increased through a process called Hybridization, which results in the formation of stronger Bonds. Using this model, we will be able to predict shapes of molecules. Figure 6. 25 Hypothetical Stepwise Process for Formation of BeH 2 from Gaseous Be Atom and Two Gaseous H Atoms promotion of Electron from 2 S Orbital of beryllium To One of 2 P Orbitals is energetically uphill. The overall process of forming BeH 2 molecule from Be Atom and Two H Atoms will therefore be energetically favorable only IF the amount of energy released by formation of two Be - H Bonds is greater than the amount of energy required for promotion and Hybridization. The Promotion and Hybridization Process is exactly the same as shown for CH 4 in Chapter. The only difference is that C Atom uses four singly occupied sp 3 hybrid Orbitals to form Electron - pair Bonds With Only Three H Atoms, and an electron is added to the fourth hybrid Orbital to give charge of 1 -. Electron - pair geometry is tetrahedral, but molecular geometry is pyramidal, as in NH 3.


9.3 Delocalized Bonding and Molecular Orbitals

Molecular Orbital theory is also able to explain the presence of lone pairs of electrons. Consider, for example, HCl molecule, whose Lewis electron structure has three lone pairs of electrons on chlorine atom. Using the molecular Orbital approach to describe Bonding in HCl, we can see from Figure 9. 31 molecular Orbital Energy - Level Diagram for HCl that 1 s Orbital of atomic hydrogen is closest in energy to 3 p orbitals of chlorine. Consequently, Cl 3 s atomic Orbital is not involved in bonding to any appreciable extent, and the only important interactions are those between H 1 s and Cl 3 p orbitals. Of the three p orbitals, only one, designated as 3 p z, can interact with H 1 s Orbital. 3 p x and 3 p y atomic orbitals have no net overlap with 1 s Orbital on hydrogen, so they are not involved in Bonding. Because energies of Cl 3 s, 3 p x, and 3 p y orbitals do not change when HCl forms, they are called nonbonding molecular orbitals molecular Orbital that form when atomic orbitals or Orbital lobes interact only very weakly, creating essentially no change in electron probability density between nuclei. Nonbonding molecular Orbital occupied by a pair of electrons is the molecular Orbital equivalent of a lone pair of electrons. By definition, electrons in nonbonding orbitals have no effect on bond order, so they are not counted in calculation of bond order. Thus, predict bond order of HCl is 2 = 1. Because Bonding molecular Orbital is closer in energy to Cl 3 p z than to H 1 s atomic Orbital, electrons in Orbital concentrate closer to chlorine atom than to hydrogen. The Molecular Orbital approach to Bonding can therefore be used to describe polarization of H–Cl bond to give {matheq}{matheq}{matheq}{endmatheq}{endmatheq}{endmatheq} as described in Chapter 8 ionic versus Covalent Bonding. Although the molecular Orbital approach reveals a great deal about bonding in give molecule, procedure quickly becomes computationally intensive for molecules of even moderate complexity. Furthermore, because computed molecular orbitals extend over entire molecule, they are often difficult to represent in a way that is easy to visualize. Therefore, we do not use the pure molecular Orbital approach to describe bonding in molecules or ions with more than two atoms. Instead, we use the valence bond approach and molecular Orbital approach to explain, among other things, concept of resonance, which cannot adequately be explained using other methods.


AX 2 E: SO 2

In ammonia, Central Atom, nitrogen, has five valence electrons and each hydrogen donates one valence Electron, producing Lewis Electron structure 2. There are four Electron Groups around nitrogen, three Bonding pairs and one Lone Pair. Repulsions are minimized by directing each hydrogen Atom and Lone Pair to corners of the tetrahedron. 3. With three Bonding pairs and one Lone Pair, structure is designated as AX 3 E. This designation has a total of four Electron pairs, three X and one E. We expect LP - BP interactions to cause Bonding Pair angles to deviate significantly from angles of perfect tetrahedron. 4. There are three nuclei and one Lone Pair, SO Molecular geometry is trigonal pyramidal. In essence, this is a tetrahedron with vertex missing. However, H - NH bond angles are less than the ideal angle of 109. 5 because of LP - BP repulsions.


Hybridization Using d Orbitals

What is the hybridization of central atom in each species? Describe bonding in each species. Xef 4 SO 4 2 SF 4 Determine geometry of molecule using the strategy in Example 1. From valence of electron configuration of central atom and number of electron pairs, determine hybridization. B Place total number of electrons around central atom in hybrid Orbitals and describe bonding. Using the VSEPR model, we find that Xe in XeF 4 forms four bonds and has two lone pairs, So its structure is square planar and it has six electron pairs. Six electron pairs form an octahedral arrangement, SO Xe must be sp 3 d 2 hybridize. B With 12 electrons around Xe, four of six sp 3 d 2 hybrid Orbitals form Xe - F bonds, and two are occupied by lone pairs of electrons. S in SO 4 2 ion has four electron pairs and has four bond atoms, SO structure is tetrahedral. Sulfur must be sp 3 hybridized to generate four S - O bonds. B Filling sp 3 hybrid Orbitals With eight electrons from four bonds produces four fill sp 3 hybrid Orbitals. S atom in SF 4 contains five electron pairs and four bond atoms. The molecule has a seesaw structure with one lone pair: To accommodate five electron pairs, sulfur atom must be sp 3 d hybridize. B Filling these Orbitals with 10 electrons gives four sp 3 d hybrid Orbitals forming S - F bonds and one with lone pair of electrons. What is the hybridization of central atom in each species? Describe bonding. Pcl 4 + BrF 3 SiF 6 2 sp 3 With four P - Cl bonds sp 3 d With three Br - F bonds and two lone pairs sp 3 d 2 With six Si - F bonds


Key Concepts and Summary

Lewis electron structures give NO information about molecular geometry, arrangement of bond atoms in molecules or polyatomic ion,s which is crucial to understanding chemistry of molecule. The Valence - shell electron - pair repulsion model allows us to predict which of possible structures is actually observed in most cases. It is based on the assumption that pairs of electrons occupy space, and the lowest - energy structure is one that minimizes electron pair - electron pair repulsions. In the VSEPR model, molecule or polyatomic ion is given AX m E N designation, where is central atom, X is bond atom, E is nonbonding valence electron group, and m and N are integers. Each group around central atom is designated as a bonding pair or lone pair. From BP and LP interactions we can predict both relative positions of atoms and angles between bonds, called bond angles. From this we can describe molecular geometry. The combination of VSEPR and bonding model, such as Lewis electron structures, however, is necessary to understand the presence of multiple bonds. Molecules with polar covalent bonds can have dipole moment, asymmetrical distribution of charge that results in a tendency for molecules to align themselves in applied electric field. Any diatomic molecule with polar covalent bond has a dipole moment, but in polyatomic molecules, presence or absence of net dipole moment depends on structure. For some highly symmetrical structures, individual bond dipole moments cancel one another, giving dipole moment of zero. The Localize bonding model assumes that covalent bonds are formed when atomic orbitals overlap and that the strength of covalent bond is proportional to the amount of overlap. It also assumes that atoms use combinations of atomic orbitals to maximize overlap with adjacent atoms. Formation of hybrid atomic orbitals can be viewed as occurring via promotion of electron from fill ns 2 subshell to empty np or d valence orbital, followed by hybridization, combination of orbitals to give new set of equivalent orbitals that are oriented properly to form bonds. The combination of ns and np orbital gives rise to two equivalent sp hybrids oriented at 180, whereas the combination of ns and two or three np orbitals produces three equivalent sp 2 hybrids or four equivalent sp 3 hybrids, respectively. Bonding in molecules with more than octet of electrons around the central atom can be explained by invoking participation of one or two d orbitals to give sets of five sp 3 d or six sp 3 d 2 hybrid orbitals, capable of forming five or six bonds, respectively. Spatial orientation of hybrid atomic orbitals is consistent with geometries predicted using the VSEPR model. Molecular orbital is allows spatial distribution of electrons in molecule that is associated with particular orbital energy.


sp Hybridization

The Beryllium atom in gaseous BeCl 2 molecule is an example of a central atom with no lone pair of electrons in a linear arrangement of three atoms. There are two regions of valence electron density in BeCl 2 molecule that correspond to two covalent Be - Cl bonds. To accommodate these two electron domains, two of Be atoms four valence orbitals will mix to yield two hybrid orbitals. This hybridization process involves mixing of valence s orbital with one of valence p orbitals to yield two equivalent sp hybrid orbitals that are oriented in linear geometry. In this figure, set of sp orbitals appear similar in shape to the original p orbital, but there is an important difference. The number of atomic orbitals combined always equals the number of hybrid orbitals form. A p orbital is one orbital that can hold up to two electrons. Sp set is two equivalent orbitals that point 180 from each other. Two electrons that were originally in s orbital are now distributed to two sp orbitals, which are half filled. In gaseous BeCl 2, these half - fill hybrid orbitals will overlap with orbitals from chlorine atoms to form two identical bonds.

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions.

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions

Section 11: Conclusion

Writing Lewis Structures, NASA's Cassini - Huygens mission detected a large cloud of toxic hydrogen cyanide on Titan, one of Saturn's moons. Titan also contains ethane, acetylene, and ammonia. What are Lewis structures of these molecules? Calculate the number of valence electrons. Hcn: + = 10H 3 CCH 3: + = 14HCCH: + = 10NH 3: + = 8 Draw skeleton and connect atoms with single bonds. Remember that H is never central atom: Where needed to distribute electrons to terminal atoms: HCN: six electrons placed on NH 3 CCH 3: no electrons remainHCCH: no terminal atoms capable of accepting electrons. Nh 3: no terminal atoms capable of accepting electrons Where needed to place remaining electrons on the central atom: HCN: no electrons remainH 3 CCH 3: no electrons remainHCCH: four electrons placed on carbon NH 3: two electrons placed on nitrogen Where needed to rearrange electrons to form multiple bonds in order to to obtain octet on each atom: HCN: form two more C - N bondsH 3 CCH 3: all atoms have correct number of electronsHCCH: form triple bond between two carbon atomsNH 3: all atoms have correct number of electrons check Your Learning Both carbon monoxide, CO, and carbon dioxide, CO 2, are products of combustion of fossil fuels. Both of these gases also cause problems: CO is toxic and CO 2 has been implicated in global climate change. What are Lewis structures of these two molecules?


Lewis Structures

For very simple molecules and molecular ions, we can write Lewis structures by merely pairing up unpaired electrons on constituent atoms. See these examples: For more complicated molecules and molecular ions, it is helpful to follow the step - by - step procedure outlined here: determining total number of valence electrons. For cations, subtract one electron for each positive charge. For anions, add one electron for each negative charge. Draw skeleton structure of a molecule or ion, arranging atoms around the central atom. Connect each atom to the central atom with a single bond. Distribute remaining electrons as lone pairs on terminal atoms, completing octet around each atom. Place all remaining electrons on the central atom. Rearrange electrons OF outer atoms to make multiple bonds with central atom in order to obtain octets wherever possible. Let us determine Lewis structures OF SiH 4, {matheq}{endmatheq} NO +, and OF 2 as examples in following this procedure: determine total number OF valence electrons in molecule or ion. For molecule, we add the number OF valence electrons on each atom in molecule: {matheq}SiH4 Si: 4 valence electrons/atom×1 atom=4 ¯ =8 valence electrons{endmatheq} For negative ion, such as {matheq}{endmatheq} we add the number OF valence electrons on atoms to the number of negative charges on ion: {matheq}CHO2−C: 4 valence electrons/atom×1 atom=4H: 1 valence electron/atom×1 atom=1O: 6 valence electrons/atom×2 atoms=12¯=18 valence electrons{endmatheq} For positive ion, such as NO +, we add the number OF valence electrons on atoms in ion and then subtract number OF positive charges on ion from total number OF valence electrons: {matheq}NO+N: 5 valence electrons/atom×1 atom=5O: 6 valence electron/atom×1 atom=6¯=10 valence electrons{endmatheq} since OF 2 is neutral molecule, We simply add number OF valence electrons: {matheq}OF2O: 6 valence electrons/atom×1 atom=6¯=20 valence electrons{endmatheq} draw skeleton structure OF molecule or ion, arranging atoms around central atom and connecting each atom to central atom with single bond. When several arrangements OF atoms are possible, as for {matheq}{endmatheq} we must use experimental evidence to choose the correct one. In general, less electronegative elements are more likely to be central atoms. In {matheq}{endmatheq} less electronegative carbon atom occupies central position with oxygen and hydrogen atoms surrounding it. Other examples include P in POCl 3, S in SO 2, and Cl. In {matheq}{endmatheq} exception is that hydrogen is almost never the central atom. Like most electronegative element,ss fluorine also cannot be central atom. Distribute remaining electrons as lone pairs on terminal atoms to complete their valence shells with octet OF electrons. There are NO remaining electrons on SiH 4, SO it is unchanged: Place all remaining electrons on the central atom. For SiH 4, {matheq}{endmatheq} and NO +, there are NO remaining electrons; We already place all OF electrons determined in Step 1. For OF 2, we had 16 electrons remaining in Step 3, and we Place 12, leaving 4 to be Place on the central atom: Rearrange electrons OF outer atoms to make multiple bonds with the central atom in order to obtain octets wherever possible.

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions.

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions

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* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions.

* Please keep in mind that all text is machine-generated, we do not bear any responsibility, and you should always get advice from professionals before taking any actions

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