what is wrong with the following statement? atoms form bonds in order to satisfy the octet rule.

Chemical rule of pollex

The bonding in carbon dioxide (CO₂): all atoms are surrounded past 8 electrons, fulfilling the octet dominion.

The octet rule is a chemic rule of thumb that reflects the theory that chief-grouping elements tend to bond in such a fashion that each cantlet has viii electrons in its valence shell, giving it the same electronic configuration as a noble gas. The rule is especially applicable to carbon, nitrogen, oxygen, and the halogens, merely also to metals such equally sodium or magnesium. Other rules exist for other elements, such as the duplet rule for hydrogen and helium, or the 18-electron rule for transition metals.

The valence electrons tin can be counted using a Lewis electron dot diagram as shown at the correct for carbon dioxide. The electrons shared past the two atoms in a covalent bond are counted twice, once for each cantlet. In carbon dioxide each oxygen shares four electrons with the central carbon, two (shown in red) from the oxygen itself and 2 (shown in black) from the carbon. All iv of these electrons are counted in both the carbon octet and the oxygen octet, so that both atoms are considered to obey the octet dominion.

Example: sodium chloride (NaCl) [edit]

Ionic bonding is common between pairs of atoms, where one of the pair is a metal of depression electronegativity (such as sodium) and the 2nd a nonmetal of high electronegativity (such equally chlorine).

A chlorine cantlet has 7 electrons in its third and outer electron shell, the first and second shells being filled with 2 and eight electrons respectively. The showtime electron affinity of chlorine (the energy release when chlorine gains an electron to form Cl⁻) is 349 kJ per mole of chlorine atoms.^[1] Adding a second electron to form a hypothetical Cl^2- would require energy, energy that cannot be recovered by the formation of a chemical bond. The outcome is that chlorine volition very oft form a compound in which it has 8 electrons in its outer shell (a complete octet), every bit in Cl⁻.

A sodium cantlet has a single electron in its outermost electron beat, the start and second shells again being total with two and eight electrons respectively. To remove this outer electron requires only the kickoff ionization energy, which is +495.8 kJ per mole of sodium atoms, a small corporeality of free energy. Past contrast, the second electron resides in the deeper 2d electron beat, and the 2d ionization energy required for its removal is much larger: +4562 kJ per mole.^[2] Thus sodium will, in most cases, form a chemical compound in which it has lost a unmarried electron and have a total outer shell of 8 electrons, or octet.

The energy required to transfer an electron from a sodium atom to a chlorine atom (the difference of the 1st ionization free energy of sodium and the electron affinity of chlorine) is small-scale: +495.8 − 349 = +147 kJ mol⁻ⁱ. This energy is easily get-go by the lattice energy of sodium chloride: −783 kJ mol⁻ⁱ.^[3] This completes the explanation of the octet rule in this instance.

History [edit]

In 1864, the English pharmacist John Newlands classified the sixty-2 known elements into 8 groups, based on their concrete properties.^{[iv] [5] [6] [7]}

In the late 19th century, it was known that coordination compounds (formerly chosen "molecular compounds") were formed by the combination of atoms or molecules in such a manner that the valencies of the atoms involved obviously became satisfied. In 1893, Alfred Werner showed that the number of atoms or groups associated with a central atom (the "coordination number") is oftentimes 4 or 6; other coordination numbers up to a maximum of 8 were known, but less frequent.^[8] In 1904, Richard Abegg was one of the showtime to extend the concept of coordination number to a concept of valence in which he distinguished atoms as electron donors or acceptors, leading to positive and negative valence states that greatly resemble the modernistic concept of oxidation states. Abegg noted that the difference between the maximum positive and negative valences of an element under his model is oftentimes 8.^[9] In 1916, Gilbert N. Lewis referred to this insight equally Abegg'southward rule and used information technology to aid formulate his cubical cantlet model and the "rule of eight", which began to distinguish between valence and valence electrons.^[x] In 1919, Irving Langmuir refined these concepts further and renamed them the "cubical octet atom" and "octet theory".^[11] The "octet theory" evolved into what is at present known as the "octet dominion".

Walther Kossel^[12] and Gilbert Due north. Lewis saw that noble gases did non have the tendency of taking office in chemical reactions under ordinary atmospheric condition. On the footing of this observation, they concluded that atoms of noble gases are stable and on the basis of this decision they proposed a theory of valency known as "electronic theory of valency" in 1916:

During the formation of a chemical bond, atoms combine together by gaining, losing or sharing electrons in such a way that they larn nearest noble gas configuration. ^[13]

Caption in quantum theory [edit]

The quantum theory of the atom explains the viii electrons as a closed crush with an s^twop⁶ electron configuration. A closed-shell configuration is one in which low-lying energy levels are full and college energy levels are empty. For case, the neon cantlet ground country has a full n = two trounce (2s² 2p⁶) and an empty due north = 3 shell. According to the octet dominion, the atoms immediately before and later on neon in the periodic tabular array (i.e. C, N, O, F, Na, Mg and Al), tend to attain a similar configuration by gaining, losing, or sharing electrons.

The argon atom has an analogous 3s² 3p^six configuration. There is also an empty 3d level, just it is at considerably higher energy than 3s and 3p (dissimilar in the hydrogen atom), so that 3s² 3p⁶ is withal considered a closed crush for chemic purposes. The atoms immediately before and after argon tend to attain this configuration in compounds. There are, however, some hypervalent molecules in which the 3d level may play a part in the bonding, although this is controversial (run into below).

For helium at that place is no 1p level according to the quantum theory, so that 1s² is a airtight shell with no p electrons. The atoms before and after helium (H and Li) follow a duet rule and tend to have the same 1sⁱⁱ configuration as helium.

Exceptions [edit]

Many reactive intermediates are unstable and do non obey the octet dominion. This includes species such as carbenes, borane also as free radicals like the methyl radical (CH₃) which has an unpaired electron in a non-bonding orbital on the carbon cantlet, and no electron of opposite spin in the aforementioned orbital. Another example is the chlorine radical produced by CFCs, known to be harmful to the ozone layer. These molecules frequently react so as to complete their octet.

Although stable odd-electron molecules and hypervalent molecules are commonly taught as violating the octet rule, ab initio molecular orbital calculations show that they largely obey the octet rule (see three-electron bonds and hypervalent molecules sections below).

3-electron bonds [edit]

Lewis and MO diagrams of an private 2e bail and 3e bond

Some stable molecular radicals (east.one thousand. nitric oxide, NO) obtain octet configurations by means of a iii-electron bond which contributes one shared and ane unshared electron to the octet of each bonded cantlet.^[14] In NO, the octet on each atom consists of two electrons from the three-electron bond, plus four electrons from two two-electron bonds and ii electrons from a solitary pair of not-bonding electrons on that atom alone. The bond order is 2.5, since each 2-electron bond counts every bit one bond while the three-electron bond has only one shared electron and therefore corresponds to a half-bond.

Dioxygen is sometimes represented as obeying the octet dominion with a double bond (O=O) containing ii pairs of shared electrons.^[15] Yet the ground state of this molecule is paramagnetic, indicating the presence of unpaired electrons. Pauling proposed that this molecule really contains two iii-electron bonds and one normal covalent (two-electron) bond.^[16] The octet on each atom and so consists of ii electrons from each three-electron bond, plus the ii electrons of the covalent bond, plus 1 solitary pair of non-bonding electrons. The bond order is 1+0.five+0.v=2.

Nitric oxide

Dioxygen

Hypervalent molecules [edit]

Chief-group elements in the third and later rows of the periodic tabular array can course hypercoordinate or hypervalent molecules in which the central master-group atom is bonded to more than 4 other atoms, such equally phosphorus pentafluoride, PF₅, and sulfur hexafluoride, SF₆. For example, in PF₅, if it is supposed that there are five truthful covalent bonds in which five distinct electron pairs are shared, then the phosphorus would be surrounded by 10 valence electrons in violation of the octet dominion. In the early days of breakthrough mechanics, Pauling proposed that tertiary-row atoms can form 5 bonds by using one due south, three p and one d orbitals, or six bonds by using one s, three p and ii d orbitals.^[17] To form five bonds, the one s, three p and one d orbitals combine to grade five sp^threed hybrid orbitals which each share an electron pair with a halogen atom, for a total of 10 shared electrons, 2 more than the octet dominion predicts. Similarly to form six bonds, the 6 spⁱⁱⁱd² hybrid orbitals grade vi bonds with 12 shared electrons.^[18] In this model the availability of empty d orbitals is used to explain the fact that third-row atoms such as phosphorus and sulfur tin can form more than four covalent bonds, whereas second-row atoms such as nitrogen and oxygen are strictly limited by the octet rule.^[xix]

five resonance structures of phosphorus pentafluoride

However other models describe the bonding using only south and p orbitals in agreement with the octet rule. A valence bond description of PF₅ uses resonance between different PF₄ ⁺ F⁻ structures, so that each F is bonded by a covalent bond in four structures and an ionic bond in one structure. Each resonance structure has eight valence electrons on P.^[xx] A molecular orbital theory description considers the highest occupied molecular orbital to be a non-bonding orbital localized on the five fluorine atoms, in improver to 4 occupied bonding orbitals, so over again at that place are only eight valence electrons on the phosphorus.^{[ citation needed ]} The validity of the octet rule for hypervalent molecules is further supported by ab initio molecular orbital calculations, which testify that the contribution of d functions to the bonding orbitals is pocket-size.^{[21] [22]}

Withal, for historical reasons, structures implying more than than eight electrons around elements like P, S, Se, or I are nonetheless mutual in textbooks and research manufactures. In spite of the unimportance of d beat expansion in chemical bonding, this exercise allows structures to be shown without using a large number of formal charges or using partial bonds and is recommended by the IUPAC equally a user-friendly formalism in preference to depictions that improve reverberate the bonding. On the other hand, showing more than eight electrons around Exist, B, C, N, O, or F (or more than two effectually H, He, or Li) is considered an fault by most government.

Other rules [edit]

The octet rule is only applicable to main-group elements. Other elements follow other electron counting rules as their valence electron configurations are different from main-group elements. These other rules are shown below:

Element type	Offset shell	p-cake (Principal group)	d-block (Transition metal)
Electron counting rules	Duet/Duplet rule	Octet rule	18-electron dominion
Total valence configuration	due south2	southward2psix	d10s2phalf-dozen

• The duet dominion or duplet rule of the first shell applies to H, He and Li—the noble gas helium has two electrons in its outer beat, which is very stable. (Since at that place is no 1p subshell, idue south is followed immediately by 2southward, and thus shell 1 tin can only have at virtually 2 valence electrons). Hydrogen just needs one additional electron to accomplish this stable configuration, while lithium needs to lose ane.
• For transition metals, molecules tend to obey the 18-electron rule which corresponds to the utilization of valence d, s and p orbitals to form bonding and non-bonding orbitals. Nevertheless, unlike the octet dominion for main-grouping elements, transition metals do not strictly obey the 18-electron dominion and the valence electron count can vary between 12 to 18.^{[23] [24] [25] [26]}

Run into also [edit]

• Lewis structure
• Electron counting

References [edit]

1. ^ Housecroft, Catherine East.; Sharpe, Alan G. (2005). Inorganic Chemistry (second ed.). Pearson Education Limited. p. 883. ISBN0130-39913-2. Source gives enthalpy change -349 kJ corresponding to energy release +349 kJ
2. ^ Housecroft, Catherine E.; Sharpe, Alan G. (2005). Inorganic Chemistry (second ed.). Pearson Education Limited. p. 880. ISBN0130-39913-2.
3. ^ Housecroft, Catherine East.; Sharpe, Alan One thousand. (2005). Inorganic Chemistry (2d ed.). Pearson Education Limited. p. 156. ISBN0130-39913-2.
4. ^ Come across:
   • Newlands, John A. R. (7 Feb 1863). "On relations among the equivalents". The Chemical News. 7: lxx–72.
   • Newlands, John A. R. (20 August 1864). "On relations among the equivalents". The Chemical News. 10: 94–95.
   • Newlands, John A. R. (18 August 1865). "On the law of octaves". The Chemical News. 12: 83.
   • (Editorial staff) (ix March 1866). "Proceedings of Societies: Chemical Lodge: Thursday, March 1". The Chemical News. 13: 113–114.
   • Newlands, John A.R. (1884). On the Discovery of the Periodic Law and on Relations amid the Atomic Weights. E. & F.N. Spon: London, England.
5. ^ in a alphabetic character published in Chemistry News in February 1863, according to the Notable Names Data Base of operations
6. ^ Newlands on classification of elements
7. ^ Ley, Willy (Oct 1966). "For Your Information: The Delayed Discovery". Galaxy Science Fiction. 25 (1): 116–127.
8. ^ Run across:
   • Werner, Alfred (1893). "Beitrag zur Konstitution anorganischer Verbindungen" [Contribution to the constitution of inorganic compounds]. Zeitschrift für anorganische und allgemeine Chemie (in High german). 3: 267–330. doi:10.1002/zaac.18930030136.
   • English translation: Werner, Alfred; Kauffman, G.B., trans. & ed. (1968). Classics in Coordination Chemistry, Office I: The selected papers of Alfred Werner. New York Urban center, New York, USA: Dover Publications. pp. 5–88.
9. ^ Abegg, R. (1904). "Die Valenz und das periodische Arrangement. Versuch einer Theorie der Molekularverbindungen" [Valency and the periodic system. Attempt at a theory of molecular compounds]. Zeitschrift für Anorganische Chemie. 39 (1): 330–380. doi:10.1002/zaac.19040390125.
10. ^ Lewis, Gilbert Due north. (1916). "The Atom and the Molecule". Periodical of the American Chemic Society. 38 (4): 762–785. doi:10.1021/ja02261a002.
11. ^ Langmuir, Irving (1919). "The Arrangement of Electrons in Atoms and Molecules". Journal of the American Chemical Order. 41 (6): 868–934. doi:x.1021/ja02227a002.
12. ^ Kossel, W. (1916). "Über Molekülbildung als Frage des Atombaus" [On the germination of molecules as a question of atomic structure]. Annalen der Physik (in German). 354 (3): 229–362. Bibcode:1916AnP...354..229K. doi:10.1002/andp.19163540302.
13. ^ "The Atom and the Molecule. April 1916. - Published Papers and Official Documents - Linus Pauling and The Nature of the Chemical Bond: A Documentary History". Osulibrary.oregonstate.edu. Archived from the original on November 25, 2013. Retrieved 2014-01-03 .
14. ^ Harcourt, Richard D., ed. (2015). "Chapter 2: Pauling "3-Electron Bonds", four-Electron three-Eye Bonding, and the Demand for an "Increased-Valence" Theory". Bonding in Electron-Rich Molecules: Qualitative Valence-Bond Approach via Increased-Valence Structures. Springer. ISBN9783319166766.
15. ^ For example, General chemical science past R.H.Petrucci, W.S.Harwood and F.G.Herring (8th ed., Prentice-Hall 2002, ISBN 0-xiii-014329-4, p.395) writes the Lewis structure with a double bond, but adds a question marking with the caption that there is some doubt about the validity of this construction because information technology fails to account for the observed paramagnetism.
16. ^ 50. Pauling The Nature of the Chemical Bail (3rd ed., Oxford University Printing 1960) affiliate ten.
17. ^ L. Pauling The Nature of the Chemical Bond (third ed., Oxford University Press 1960) p.63. In this source Pauling considers as examples PCl₅ and the PF₆ ⁻ ion. ISBN 0-8014-0333-2
18. ^ R.H. Petrucci, W.S. Harwood and F.G. Herring, Full general Chemistry (8th ed., Prentice-Hall 2002) p.408 and p.445 ISBN 0-13-014329-4
19. ^ Douglas B.E., McDaniel D.H. and Alexander J.J. Concepts and Models of Inorganic Chemical science (2nd ed., John Wiley 1983) pp.45-47 ISBN 0-471-21984-three
20. ^ Housecroft C.E. and Sharpe A.G., Inorganic Chemistry, 2nd ed. (Pearson Teaching Ltd. 2005), p.390-1
21. ^ Miessler D.L. and Tarr G.A., Inorganic Chemistry, second ed. (Prentice-Hall 1999), p.48
22. ^ Magnusson, Eastward., J.Am.Chem.Soc. (1990), v.112, p.7940-51 Hypercoordinate Molecules of 2nd-Row Elements: d Functions or d Orbitals?
23. ^ Frenking, Gernot; Shaik, Sason, eds. (May 2014). "Chapter 7: Chemic bonding in Transition Metal Compounds". The Chemical Bail: Chemical Bonding Across the Periodic Table. Wiley -VCH. ISBN978-iii-527-33315-8.
24. ^ Frenking, Gernot; Fröhlich, Nikolaus (2000). "The Nature of the Bonding in Transition-Metal Compounds". Chem. Rev. 100 (2): 717–774. doi:10.1021/cr980401l. PMID 11749249.
25. ^ Bayse, Craig; Hall, Michael (1999). "Prediction of the Geometries of Unproblematic Transition Metal Polyhydride Complexes by Symmetry Assay". J. Am. Chem. Soc. 121 (6): 1348–1358. doi:10.1021/ja981965+.
26. ^ King, R.B. (2000). "Structure and bonding in homoleptic transition metal hydride anions". Coordination Chemistry Reviews. 200–202: 813–829. doi:ten.1016/S0010-8545(00)00263-0.

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Source: https://en.wikipedia.org/wiki/Octet_rule

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