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Can Xenon Have An Expanded Octet

author
Ellen Grant
• Monday, 23 November, 2020
• 14 min read

Rebekah Jones in jail after Florida issues warrant The electrons in the outer shell of Xenon are much further away from the nucleus, than is the case for Neon.

(Source: samsongorodki.ru)

Contents

Xenon has an empty d-orbital in its outer shell that can accept e- because 3p and 3d have similar energies. 1) Neon (symbol: Ne) is an element (noble gas) with atomic number 10, which means it has 10 protons and 10 electrons.

I think elements can break the octet rule as long as they are n=3 or higher, since then they would have d-orbitals to put electrons in. When writing the Lewis Structure, you first place any extra electrons on the outside atoms.

Like other elements with n=3 or higher, xenon has an empty d-orbital, specifically the 5d orbital, which can fill up with the extra electrons. This is possible since xenon is a large atom with valence electrons far away from its nucleus (relative to the noble gases that precede it) and fluorine is electronegative enough to pull away Xenon valence electrons allowing for an expanded octet to form.

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Describe the ways that B, Al, Li, and H deviate from the octet rule Key Points The octet rule states that atoms with an atomic number below 20 tend to combine so that they each have eight electrons in their valence shells, which gives them the same electronic configuration as a noble gas.

The rule is applicable to the main- group elements, especially carbon, nitrogen, oxygen, and the halogens, but also to metals such as sodium and magnesium. Boron and aluminum, from Group III (or 13), display different bonding behavior than previously discussed.

However, compounds in which boron or aluminum atoms form five bonds are never observed, so we must conclude that simple predictions based on the octet rule are not reliable for Group III. Notice that the central boron atom has only 6 electrons in the final Lewis diagram/structure of this molecule.

We might conclude from this one example that boron atoms obey a sextet rule. In addition, BF 3 will react with ammonia (NH 3), to form a stable compound, NH 3 BF 3, for which a Lewis structure can be drawn that shows boron with a complete octet.

The most commonly encountered stable species that exist with an odd number of electrons are nitrogen oxides, such as nitric oxide (NO) and nitrogen dioxide (NO 2), both of which are free radicals and disobey the octet rule. Free radical : Any molecule, ion, or atom with one or more unpaired electrons.

Hydrogen is an exception because it can hold a maximum of two electrons in its valence level. Some elements, most notably nitrogen, can form compounds that do not obey the octet rule.

Molecules with unpaired electrons are termed ‘free radicals.’ While typically highly unstable, and therefore highly reactive, some free radicals exhibit stability of days, months, or even years. Recall that the Lewis structure of a molecule must depict the total number of valence electrons from all the atoms which are bonded together.

Nitrogen dioxide is an intermediate in the industrial synthesis of nitric acid, millions of tons of which is produced each year. This reddish-brown toxic gas has a characteristic sharp, biting odor and is a prominent air pollutant.

Main group elements in the third period and below form compounds that deviate from the octet rule by having more than 8 valence electrons. Key Points Main group elements that form more bonds than would be predicted by the octet rule are called prevalent compounds, and have what is known as an ‘ expanded octet, ’ meaning that there are more than eight electrons around one atom.

The octet rule can be ‘ expanded by some elements by utilizing the d- orbitals found in the third principal energy level and beyond. Sulfur, phosphorus, silicon, and chlorine are common examples of elements that form an expanded octet.

Phosphorus pentachloride (PCL 5) and sulfur hexafluoride (SF 6) are examples of molecules that deviate from the octet rule by having more than 8 electrons around the central atom. Expanded octet : A case where an atom shares more than eight electrons with its bonding partners.

Phosphorus pentachloride (PCL 5), sulfur hexafluoride (SF 6), chlorine fluoride (Cf 3), and the trioxide ion (I 3 ) are examples of prevalent molecules. As a result, the second period elements (more specifically, the nonmetals C, N, O, F) obey the octet rule without exceptions.

The overall geometry of the molecule is depicted (trigonal bipyramidal), and bond angles and lengths are highlighted. However, some third-period elements (Si, P, S, and Cl) have been observed to bond to more than four other atoms, and thus need to involve more than the four pairs of electrons available in an s 2 p 6 octets.

Examples of molecules in which a third period central atom contains an expanded octet are the phosphorus pentahalides and sulfur hexafluoride. The overall geometry of the molecule is depicted (tetragonal bipyramidal, or octahedral), and bond angles and lengths are highlighted.

For atoms in the fourth period and beyond, higher d orbitals can be used to accommodate additional shared pairs beyond the octet. Sulfur, phosphorus, silicon, and chlorine are common examples of elements that form an expanded octet.

Phosphorus pentachloride (PCL 5) and sulfur hexafluoride (SF 6) are examples of molecules that deviate from the octet rule by having more than 8 electrons around the central atom. Hypervalent molecules molecule that contains an atom from a main group element which deviates from the octet rule by sharing more than eight electrons.

Phosphorus pentachloride (PCL 5), sulfur hexafluoride (SF 6), chlorine fluoride (Cf 3), and the trioxide ion (I 3 ) are examples of prevalent molecules. However, some third-period elements (Si, P, S, and Cl) have been observed to bond to more than four other atoms, and thus need to involve more than the four pairs of electrons available in an s 2 p 6 octets.

The overall geometry of the molecule is depicted (tetragonal bipyramidal, or octahedral), and bond angles and lengths are highlighted. For atoms in the fourth period and beyond, higher d orbitals can be used to accommodate additional shared pairs beyond the octet.

The central sulfur atom shows an expanded valence-shell octet common in Period 3 elements. The central sulfur atom in this Lewis structure shows an impossibly high oxidation state.

This Lewis structure is a poor illustration of the actual electron distribution for this molecule. Access to all video lessonsQuizzes, practice exams & worksheetsCertificate of CompletionAccess to instructors See how much you understand about the exceptions to the octet rule in chemistry.

Do the two xenon compounds shown below have an expanded valence shell? Step 1 of 5 The octet rule is a theory used to predict the molecular structure of covalently bonded molecules.

Each atom seeks stability by attaining noble gas configuration (8 electrons in valence shell). Atoms in these periods may follow the octet rule, but there are conditions where they can expand their valence shells to accommodate more than eight electrons.

Sulfur can follow the octet rule as in the molecule. Thus far in this chapter, we have discussed the various types of bonds that form between atoms and/or ions.

In all cases, these bonds involve the sharing or transfer of valence shell electrons between atoms. In this section, we will explore the typical method for depicting valence shell electrons and chemical bonds, namely Lewis symbols and Lewis structures.

Likewise, they can be used to show the formation of anions from atoms, as shown here for chlorine and sulfur: Figure 2 demonstrates the use of Lewis symbols to show the transfer of electrons during the formation of ionic compounds.

Cations are formed when atoms lose electrons, represented by fewer Lewis dots, whereas anions are formed by atoms gaining electrons. We also use Lewis symbols to indicate the formation of covalent bonds, which are shown in Lewis structures, drawings that describe the bonding in molecules and polyatomic ions.

A dash (or line) is sometimes used to indicate a shared pair of electrons: Each Cl atom interacts with eight valence electrons: the six in the lone pairs and the two in the single bond.

The other halogen molecules (F 2, Br 2, I 2, and At 2) form bonds like those in the chlorine molecule: one single bond between atoms and three lone pairs of electrons per atom. This allows each halogen atom to have a noble gas electron configuration.

The tendency of main group atoms to form enough bonds to obtain eight valence electrons is known as the octet rule. The number of bonds that an atom can form can often be predicted from the number of electrons needed to reach an octet (eight valence electrons); this is especially true of the nonmetals of the second period of the periodic table (C, N, O, and F).

Because hydrogen only needs two electrons to fill its valence shell, it is an exception to the octet rule. To obtain an octet, these atoms form three covalent bonds, as in NH 3 (ammonia).

Oxygen and other atoms in group 16 obtain an octet by forming two covalent bonds: For very simple molecules and molecular ions, we can write the Lewis structures by merely pairing up the unpaired electrons on the constituent atoms.

For more complicated molecules and molecular ions, it is helpful to follow the step-by-step procedure outlined here: Determine the total number of valence (outer shell) electrons.

Let us determine the Lewis structures of Six 4, CHO 2 , NO +, and OF 2 as examples in following this procedure: Determine the total number of valence (outer shell) electrons in the molecule or ion.

For a positive ion, such as NO +, we add the number of valence electrons on the atoms in the ion and then subtract the number of positive charges on the ion (one electron is lost for each single positive charge) from the total number of valence electrons: \begin{array}{r l} \text{NO}^{+} & & \\ & \text{N: 5 valence electrons/atom} \times 1 \;\text{atom} & = 5 \\ & \text{O: 6 valence electrons/atom} \times 1 \;\text{atom} & = 6 \\ \rule{21em}{0.1ex}\space{-21em} + & -1 \;\text{electron (positive charge)} & = -1 \\ & & = 10 \;\text{valence electrons} \end{array} (Note that we denote ions with brackets around the structure, indicating the charge outside the brackets:) When several arrangements of atoms are possible, as for CHO 2 , we must use experimental evidence to choose the correct one.

Distribute the remaining electrons as lone pairs on the terminal atoms (except hydrogen) to complete their valence shells with an octet of electrons. We cannot add any more electrons since we have already used the total that we found in Step 1, so we must move electrons to form a multiple bond: This still does not produce an octet, so we must move another pair, forming a triple bond: In OF 2, each atom has an octet as drawn, so nothing changes.

Writing Lewis Structures NASA’s Cassini-Huygens mission detected a large cloud of toxic hydrogen cyanide (HAN) on Titan, one of Saturn’s moons. Titan also contains ethane (H 3 CCH 3), acetylene (HCC), and ammonia (NH 3).

Both of these gases also cause problems: CO is toxic and CO 2 has been implicated in global climate change. In 1996, the Nobel Prize in Chemistry was awarded to Richard Smaller (Figure 3), Robert Curl, and Harold Kyoto for their work in discovering a new form of carbon, the C 60 buckminsterfullerene molecules (Figure 1 in Chapter 7 Introduction).

An entire class of compounds, including spheres and tubes of various shapes, were discovered based on C 60. This type of molecule, called a fullerene, shows promise in a variety of applications.

Because of their size and shape, fullerenes can encapsulate other molecules, so they have shown potential in various applications from hydrogen storage to targeted drug delivery systems. They also possess unique electronic and optical properties that have been put to good use in solar-powered devices and chemical sensors.

Upon his death in 2005, the US Senate honored him as the “Father of Nanotechnology.” (credit: United States Department of Energy) Many covalent molecules have central atoms that do not have eight electrons in their Lewis structures. Electron-deficient molecules have a central atom that has fewer electrons than needed for a noble gas configuration.

Prevalent molecules have a central atom that has more electrons than needed for a noble gas configuration. We call molecules that contain an odd number of electrons free radicals.

Nitric oxide, NO, is an example of an odd-electron molecule; it is produced in internal combustion engines when oxygen and nitrogen react at high temperatures. Determine the total number of valence (outer shell) electrons.

The odd number immediately tells us that we have a free radical, so we know that not every atom can have eight electrons in its valence shell. We can easily draw a skeleton with an N–O single bond:N–O Distribute the remaining electrons as lone pairs on the terminal atoms.

We give eight electrons to the more electronegative atom in these situations; thus oxygen has the filled valence shell: Place all remaining electrons on the central atom. Rearrange the electrons to make multiple bonds with the central atom in order to obtain octets wherever possible.

To move closer to an octet for nitrogen, we take one of the lone pairs from oxygen and use it to form a NO double bond. We will also encounter a few molecules that contain central atoms that do not have a filled valence shell.

This suggests the best Lewis structure has three B–F single bonds and an electron deficient boron. The reactivity of the compound is also consistent with an electron deficient boron.

Figure 4 shows the Lewis structures for two prevalent molecules, PCL 5 and SF 6. In PCL 5, the central atom phosphorus shares five pairs of electrons.

In SF 6, sulfur shares six pairs of electrons. In some prevalent molecules, such as IF 5 and Ref 4, some electrons in the outer shell of the central atom are lone pairs: Writing Lewis Structures: Octet Rule Violations Xenon is a noble gas, but it forms a number of stable compounds.

Solution We can draw the Lewis structure of any covalent molecule by following the six steps discussed earlier. Calculate the number of valence electrons: Ref 2 : 8 + (2 × 7) = 22XeF 6 : 8 + (6 × 7) = 50 Draws a skeleton joining the atoms by single bonds.

Write the Lewis structures for the interhalogens BCL 3 and ICL 4 . Valence electronic structures can be visualized by drawing Lewis symbols (for atoms and anatomic ions) and Lewis structures (for molecules and polyatomic ions).

Many planets in our solar system contain organic chemicals including methane (CH 4) and traces of ethylene (C 2 H 4), ethane (C 2 H 6), propane (H 3 CCC), and acetylene (HCC CCH). Identify the atoms that correspond to each of the following electron configurations.

Then, write the Lewis symbol for the common ion formed from each atom: (c) 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 4 s 2 3 d 10 Complete the Lewis structures of these molecules by adding multiple bonds and lone pairs.

Double bond covalent bond in which two pairs of electrons are shared between two atoms free radical molecule that contains an odd number of electrons prevalent molecule containing at least one main group element that has more than eight electrons in its valence shell Lewis structure diagram showing lone pairs and bonding pairs of electrons in a molecule or an ion Lewis symbol for an element or anatomic ion that uses a dot to represent each valence electron in the element or ion lone pair two (a pair of) valence electrons that are not used to form a covalent bond octet rule guideline that states main group atoms will form structures in which eight valence electrons interact with each nucleus, counting bonding electrons as interacting with both atoms connected by the bond single bond in which a single pair of electrons is shared between two atoms triple bond in which three pairs of electrons are shared between two atoms In this case, the Lewis structure is inadequate to depict the fact that experimental studies have shown two unpaired electrons in each oxygen molecule.

Two valence electrons per Pb atom are transferred to Cl atoms; the resulting Pb 2+ ion has a 6 s 2 valence shell configuration.

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1 en.wikipedia.org - https://en.wikipedia.org/wiki/Yvonne_Orji
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3 www.thefamousnaija.com - https://www.thefamousnaija.com/2021/01/yvonne-orji-biography-husband-age-net.html
4 thefamousdata.com - https://thefamousdata.com/yvonne-orji/
5 schoolings.org - https://schoolings.org/yvonne-orji-biography-age-biography-movies-momma-i-made-it-net-worth/