The Octet Rule

6 Key excerpts on "The Octet Rule"

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  eBook - PDF

  Introduction to General, Organic, and Biochemistry

  • Frederick Bettelheim, William Brown, Mary Campbell, Shawn Farrell(Authors)
  • 2019(Publication Date)
  • Cengage Learning EMEA

    (Publisher)

  63 CONTENTS 3.1 The Octet Rule 3.2 Naming Anions and Cations 3.3 The Two Major Types of Chemical Bonds 3.4 An Ionic Bond 3.5 Naming Ionic Compounds 3.6 A Covalent Bond How To . . . Draw Lewis Structures 3.7 Naming Binary Covalent Compounds 3.8 Resonance How To . . . Draw Curved Arrows and Push Electrons 3.9 Predicting Bond Angles in Covalent Molecules 3.10 Determining If a Molecule Is Polar Chemical Bonds 3 3.1 The Octet Rule In 1916, Gilbert N. Lewis (Section 2.6) devised a beautifully simple model that unified many of the observations about chemical bonding and chemical reactions. He pointed out that the lack of chemical reactivity of the noble gases (Group 8A) indicates a high degree of stability of their electron config-urations: helium with a filled valence shell of two electrons (1 s 2 ), neon with a filled valence shell of eight electrons (2 s 2 2 p 6 ), argon with a valence shell of eight electrons (3 s 2 3 p 6 ), and so forth. The tendency of atoms to react in ways that achieve an outer shell of eight valence electrons is particularly common among Group 1A–7A elements and is given the special name of The Octet Rule . An atom with almost eight valence electrons tends to gain the needed electrons to have eight electrons in its valence shell and an electron configuration like that of the noble gas nearest to it in atomic number. In gaining electrons, the atom becomes a neg-atively charged ion called an anion . An atom with only one or two valence electrons tends to lose the number of electrons required to have an electron configuration like the noble gas nearest it in atomic number. In losing elec-trons, the atom becomes a positively charged ion called a cation .

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  Introductory Chemistry

  An Active Learning Approach

  • Mark Cracolice, Edward Peters, Mark Cracolice(Authors)
  • 2020(Publication Date)
  • Cengage Learning EMEA

    (Publisher)

  The bonding electrons count as valence electrons for each bonded atom. Covalent bonds form by the over- lap of half-filled electron orbitals. The stability of a noble-gas electron configuration—The Octet Rule or rule of eight—is a result of the minimization of energy associated with that configuration. Goal 5 Use Lewis symbols to show how covalent bonds are formed between two nonmetal atoms. A covalent bond is formed between two nonmetal atoms, both of which have atoms that are one, two, or even three electrons short of a noble gas electron configuration. This covalent bonding process is characterized by valence electron pairs that are shared. Goal 6 Distinguish between bonding elec- tron pairs and lone pairs. Bonding electron pairs are represented in a Lewis diagram as two dots or a straight line drawn between atoms. Both formats represent the covalent bond that holds the atoms together. Unshared pairs are also shown in Lewis diagrams. These are called lone pairs. Goal 7 Distinguish between polar and non- polar covalent bonds. In a nonpolar covalent bond, the bonding electrons are shared equally by the bonded atoms. In a polar covalent bond, the nucleus of one atom attracts the shared electrons more strongly than the other. Goal 8 Predict which end of a polar bond between identified atoms is positive and which end is negative. The relative ability of atoms of an element to attract electron pairs in covalent bonds is expressed by the electronegativity of the element. The polarity of a bond is estimated by the difference in electronegativities of the bonded atoms. The atom with the higher electronegativity is the negative end of the bond. The atom with the lower electronegativ- ity is the positive end of the bond. Goal 9 Rank bonds in order of increasing or decreasing polarity based on periodic trends in electronegativity values or actual values, if given. There is a periodic trend in the electronegativity of elements.

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  Chemistry

  The Molecular Nature of Matter

  • Neil D. Jespersen, Alison Hyslop(Authors)
  • 2014(Publication Date)
  • Wiley

    (Publisher)

  In this section we will develop a simple method for drawing Lewis structures for both molecules and polyatomic ions (which are also held together by covalent bonds). Although The Octet Rule is important in covalent bonding, it is not always obeyed. For instance, there are some molecules in which one or more atoms must have more than an octet in the valence shell. Examples are PCl 5 and SF 6 , whose Lewis structures are P Cl S Cl Cl Cl Cl F F F F F F In these molecules the formation of more than four bonds to the central atom requires that the central atom have a share of more than eight electrons. Reactivity of nonmetals and the periodic trends ■ Fluorine: F 2 + 2Cl - h 2F - + Cl 2 F 2 + 2Br - h 2F - + Br 2 F 2 + 2I - h 2F - + I 2 Chlorine: Cl 2 + 2Br - h 2Cl - + Br 2 Cl 2 + 2I - h 2Cl - + I 2 Bromine: Br 2 + 2I - h 2Br - + I 2 This statue of Prometheus overlooking the skating rink in Rockefeller Center in New York City is covered in a thin layer of gold, providing both beauty and weather resistance. philipus/Alamy 374 Chapter 8 | The Basics of Chemical Bonding There are also some molecules (but not many) in which the central atom behaves as though it has less than an octet. The most common examples involve compounds of beryl- lium and boron. Cl 9 Be 9 Cl Cl Be + 2 four electrons around Be Cl Cl 9 B 9 Cl Cl B + 3 six electrons around B Although Be and B sometimes have less than an octet, the elements in Period 2 never exceed an octet. The reason is because their valence shells, having n = 2, can hold a maxi- mum of only 8 electrons. (This explains why The Octet Rule works so well for atoms of carbon, nitrogen, and oxygen.) However, elements in periods below Period 2, such as phosphorus and sulfur, sometimes do exceed an octet, because their valence shells can hold more than 8 electrons.

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  eBook - PDF

  Structure of Matter

  With Contributions in Memoriam Including a Complete Bibliography of His Works

  • C. Guy Suits(Author)
  • 2013(Publication Date)
  • Pergamon

    (Publisher)

  For example, two fluorine atoms, each having seven electrons in its outside shell, would not be able to form octets at all except by sharing electrons. By sharing a single pair of electrons, however, two octets can be formed since two octets holding a pair in common require only 14 electrons. This is clear if we consider two cubes with electrons at each of the eight corners. When the cubes are placed so that an edge of one is in contact with an edge of the other a single pair of electrons at the ends of the common edge will take the place of four electrons in the original cubes. For each pair of electrons held in common between two octets there is a decrease of two in the total number of electrons needed to form the octets. Let e represent the number of electrons in the outside shell of the atoms that combine to form a molecule. Let n be the number of octets that are formed from these e electrons, and let p be the number of pairs of electrons which the 1 The theories of Kossel, Lacomblé, Teudt, etc., which have recently been proposed in Ger-many, have not advanced beyond this point and are therefore very unsatisfactory as a general theory of valence. The Structure of Atoms and Its Bearing on Chemical Valence 99 octets share with one another. Since every pair of electrons thus shared reduces by two the number of electrons required to form the molecule it follows that e = Sn—2p or p = l(8n— e). This simple equation tells us in each case how many pairs of electrons or chemical bonds must exist in any given molecule between the octets formed. Hydrogen nuclei, however, may attach themselves to pairs of electrons in the octets which are not already shared. For example, in the formation of hydrogen fluoride from a hydrogen atom and a fluorine atom there are 8 electrons in the shells (e = 8). We place n — 1 in the above equation and find p = 0. In other words, the fluorine atoms do not share electrons with each other.

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  Chemistry

  Principles and Reactions

  • William Masterton, Cecile Hurley(Authors)
  • 2020(Publication Date)
  • Cengage Learning EMEA

    (Publisher)

  Cengage Learning reserves the right to remove additional content at any time if subsequent rights restrictions require it. 156 CHAPTER 7 Covalent Bonding ▼ new bonding orbital. At any rate, calculations suggest that this is the principal factor accounting for the stability of the H 2 molecule. This chapter is devoted to the covalent bond as it exists in molecules and poly-atomic ions. We consider ■ ■ the distribution of outer level (valence) electrons in species in which atoms are joined by covalent bonds. These distributions are most simply described by Lewis structures (Section 7-1). ■ ■ molecular geometries. The so-called valence shell electron-pair repulsion (VSEPR) model can be used to predict the angles between covalent bonds formed by a central atom (Section 7-2). ■ ■ the polarity of covalent bonds and the molecules they form (Section 7-3). Most bonds and many molecules are polar in the sense that they have a positive and a negative pole. ■ ■ the distribution of valence electrons among atomic orbitals, using the valence bond approach (Section 7-4). 7-1 Lewis Structures; The Octet Rule The idea of the covalent bond was first suggested by the American physical chemist Gilbert Newton Lewis (1875–1946) in 1916. He pointed out that the electron configuration of the noble gases appears to be a particularly stable one. Noble-gas atoms are themselves extremely unreactive. Moreover, as pointed out in Chapter 6, a great many monatomic ions have noble-gas structures. Lewis suggested that non-metal atoms, by sharing electrons to form an electron-pair bond, can acquire a stable noble-gas structure. ▼ Consider, for example, two hydrogen atoms, each with one electron. The process by which they combine to form an H 2 molecule can be shown as H 1 H H H using dots to represent electrons; the circles emphasize that the pair of electrons in the covalent bond can be considered to occupy the 1s orbital of either hydrogen atom.

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  Chemistry

  An Atoms First Approach

  • Steven Zumdahl, Susan Zumdahl, Donald J. DeCoste, , Steven Zumdahl, Steven Zumdahl, Susan Zumdahl, Donald J. DeCoste(Authors)
  • 2020(Publication Date)
  • Cengage Learning EMEA

    (Publisher)

  See Exercises 3.59 and 3.60 LET’S REVIEW Lewis Structures: Comments About The Octet Rule » The second-row elements C, N, O, and F should always be assumed to obey The Octet Rule. » The second-row elements B and Be often have fewer than eight electrons around them in their compounds. These electron-deficient compounds are very reactive. » The second-row elements never exceed The Octet Rule, since their valence orbitals (2s and 2p) can accommodate only eight electrons. » Third-row and heavier elements often satisfy The Octet Rule but can exceed The Octet Rule. » When writing the Lewis structure for a molecule, satisfy The Octet Rule for the atoms first. If electrons remain after The Octet Rule has been satisfied, then place them on any element from Period 3 or beyond. In the PCl 5 and SF 6 molecules, the central atoms (P and S, respectively) must have the extra electrons. However, in molecules having more than one atom that can exceed The Octet Rule, it is not always clear which atom should have the extra electrons. Con- sider the Lewis structure for the triiodide ion (I 3 2 ), which has 3s7d 1 1 5 22 valence electrons h h I 21 charge Indicating the single bonds gives IOIOI. At this point, 18 electrons (22 2 4) remain. Trial and error will convince you that one of the iodine atoms must exceed The Octet Rule, but which one? The rule we will follow is that when it is necessary to exceed The Octet Rule for one of several third-row (or higher) elements, assume that the extra electrons should be placed on the central atom. Thus for I 3 2 the Lewis structure is where the central iodine exceeds The Octet Rule. This structure agrees with known prop- erties of I 3 2 . 121 3.8 Exceptions to The Octet Rule Copyright 2021 Cengage Learning. All Rights Reserved. May not be copied, scanned, or duplicated, in whole or in part. Due to electronic rights, some third party content may be suppressed from the eBook and/or eChapter(s).

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