Lewis Dot Diagrams

6 Key excerpts on "Lewis Dot Diagrams"

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  eBook - ePub

  Survival Guide to General Chemistry

  • Patrick E. McMahon, Rosemary McMahon, Bohdan Khomtchouk(Authors)
  • 2019(Publication Date)
  • CRC Press

    (Publisher)

  14 Alternate Methods for Visualizing and Constructing

  Lewis Structures of Covalent Molecules

  I	INTRODUCTION TO INTERPRETATION OF LEWIS STRUCTURES

  Lewis structures are pictorial representations, not exact theoretical analyses, of the role of all valence electrons (outer shell electrons for non-metals) in a covalent molecule. All valence electrons for all atoms in the molecule are accounted for as either covalent bonding electron pairs, shown as a line, or as electrons not used in bonding, termed lone electron pairs. A lone electron pair is usually shown as a pair of dots; occasionally, non-octet atoms will have an unpaired electron, shown as a single dot.

  The Lewis structure for an individual atom shows all valence electrons as dots around the symbol of the atom. The generally accepted method is to show the electrons as pairs or to show them symmetrically distributed around the symbol, especially when indicating bonding in molecules.

  Example: Acceptable for carbon with four valence electrons:

  Example: Acceptable for oxygen with six valence electrons:

  The standard method for constructing Lewis structures involves counting up all valence electrons for all atoms in the molecule, then distributing them into a bonding pattern according to a set of fixed rules. The methods described in this chapter adapts and extends the general bonding pattern information generated from this approach. The purpose is to generate techniques for understanding the electron behaviors responsible for bonding in typical covalent molecules. Corresponding techniques are then produced for determining valid atom connection patterns for larger covalent molecules without counting all valence electrons and in the absence of a provided bonding pattern.

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  Quantum Chemistry

  A Unified Approach

  • David B Cook(Author)
  • 2012(Publication Date)
  • ICP

    (Publisher)

  So we have to try to look at the idea of the chemical bond with fresh eyes. The first theory of the chemical bond which said something about the ‘mechanism’ of bonding — the Lewis ‘dots and crosses’ theory — involved the idea that the chemical bond was associated with the sharing of electrons between atoms, each of the lines being thought of as repre-senting a pair of electrons. Modern theories still retain this central idea, but are now able to explain the mechanism and details of this electron sharing. It is important to have a clear idea of what is meant by ‘explanation’ in science and, in particular, in the theory of the chemical bond. It does not mean simply: • A convenient way to remember some facts: e.g. ‘the valency of carbon is 4’; • A rule of thumb to get the numbers right: e.g. ‘octet (8-electron) or 18-electron rules’; • A pictorial way of representing some ideas: e.g. ‘dots and crosses’; • A numerical rule which (sometimes) gives the right answers: e.g. ‘count-ing bonding and anti-bonding electrons’, although a good explanation may well do some or all of these things. What it does mean is, for the time being anyway, A way of describing the phenomenon of chemical bond for-mation, which uses the properties of the particles of which the system is composed (electrons and nuclei), and the laws of interaction amongst these particles. 48 What We Know About Atoms and Molecules Notice the fact that, as emphasised in Chapter 1, this does not mean that chemistry is just the dynamics of electrons and nuclei. We must make a chemical choice of what the ‘system’ to study is. Going back to the start, a chemical bond is simply represented by a line drawn between two chemical symbols between atoms. This symbolism is conveniently ambiguous since: • The line may simply be a convenient way of indicating ‘what is joined to what’ in the molecule, or it may indicate some details of the structure of the molecule.

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  Electron Flow in Organic Chemistry

  A Decision-Based Guide to Organic Mechanisms

  • Paul H. Scudder(Author)
  • 2023(Publication Date)
  • Wiley

    (Publisher)

  You must be able to draw a proper Lewis structure complete with formal charges accurately and quickly. Your command of curved arrows must also be automatic. These two points cannot be overemphasized, since all explanations of reactions will be expressed in the language of Lewis structures and curved arrows. A Lewis structure contains the proper number of valence electrons, the correct distribution of those electrons over the atoms, and the correct formal charge. We will show all valence electrons; lone pairs are shown as darkened dots and bonds by lines. An atom in a molecule is most stable if it can achieve the electronic configuration of the nearest noble gas, thus having a completely filled valence shell. Hydrogen with two electrons around it, a duet, achieves the configuration of helium. Second-row elements achieve the configuration of neon with an octet of valence electrons. Third-row elements like sulfur achieve an octet but may also expand their valence shell; for example SF 6 is a stable molecule with six single bonds to sulfur (12 bonding electrons total). 1.3.1 Procedure for Drawing Lewis Structures Use the periodic table to find the valence electrons contributed from each atom. Add an additional electron for a negative charge, or subtract one to account for a positive charge to get the total number of valence electrons. Then draw single bonds between all connected atoms to establish a skeleton, or preliminary structure. You need to know the pattern in which the atoms are connected. If you have to guess at the connectivity, the most symmetrical structure is often correct. Sulfuric acid, for example, has the one sulfur in the center surrounded by the four oxygens, two of which have attached hydrogens. Since hydrogen forms one covalent bond, it is always on the outside of the structure. Place any additional bonds between adjacent atoms that both have incomplete octets to satisfy the following general bonding trends.

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  General Chemistry I as a Second Language

  Mastering the Fundamental Skills

  • David R. Klein(Author)
  • 2015(Publication Date)
  • Wiley

    (Publisher)

  196 CHAPTER 7 DRAWING LEWIS STRUCTURES For many types of problems in this course, you will need to convert a molecular formula into a Lewis structure: C 2 H 5 OH C C O H H H H H H Molecular Formula Lewis Structure The Lewis structure of a compound shows useful information: it tells us which atoms are connected together, and more importantly, where the electrons are. In this section, we will focus on the skills that you need to draw Lewis structures (fre- quently called Lewis diagrams or Lewis formulas). 7.1 A DEEPER UNDERSTANDING OF THE OCTET RULE You probably remember learning about the octet rule in high school. You probably also remember that there are exceptions to the rule. The truth is that there are MANY exceptions to the rule. This should prompt us to ask several questions: Why do we have a rule that gets broken so often? When does it get broken, and when is it fol- lowed? And what exactly is this rule in the first place? Where did the rule come from? In this section, we will take a closer look at the octet rule. We will try to un- derstand it better, so that we can predict how, why, and when it can be broken. This will be critical for drawing Lewis structures. But first, we need to quickly review some information that we saw in the previous chapter. We saw that electrons exist in predefined regions of space, called orbitals. Every orbital can hold a maximum of two electrons. So, any orbital can have one of the following possibilities: the orbital can contain (a) no electrons, (b) one elec- tron, or (c) two electrons. These are the only possibilities, since an orbital can never have more than two electrons. Let’s quickly consider each of these situations, be- cause they will prove relevant for drawing and appreciating Lewis structures. • If an orbital contains no electrons, then there is nothing to talk about. Re- member that an orbital is just a predefined region of space where 95% of the

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  Chemistry: Atoms First

  • William R. Robinson, Edward J. Neth, Paul Flowers, Klaus Theopold, Richard Langley(Authors)
  • 2016(Publication Date)
  • Openstax

    (Publisher)

  Oxyacids are named by changing the ending of the anion to –ic, and adding “acid;” H 2 CO 3 is carbonic acid. 4.4 Lewis Symbols and Structures Valence electronic structures can be visualized by drawing Lewis symbols (for atoms and monatomic ions) and Lewis structures (for molecules and polyatomic ions). Lone pairs, unpaired electrons, and single, double, or triple bonds are used to indicate where the valence electrons are located around each atom in a Lewis structure. Most structures—especially those containing second row elements—obey the octet rule, in which every atom (except H) is surrounded by eight electrons. Exceptions to the octet rule occur for odd-electron molecules (free radicals), electron- deficient molecules, and hypervalent molecules. 4.5 Formal Charges and Resonance In a Lewis structure, formal charges can be assigned to each atom by treating each bond as if one-half of the electrons are assigned to each atom. These hypothetical formal charges are a guide to determining the most appropriate Lewis structure. A structure in which the formal charges are as close to zero as possible is preferred. Resonance occurs in cases where two or more Lewis structures with identical arrangements of atoms but different distributions of electrons can be written. The actual distribution of electrons (the resonance hybrid) is an average of the distribution indicated by the individual Lewis structures (the resonance forms). 4.6 Molecular Structure and Polarity VSEPR theory predicts the three-dimensional arrangement of atoms in a molecule. It states that valence electrons will assume an electron-pair geometry that minimizes repulsions between areas of high electron density (bonds and/ or lone pairs). Molecular structure, which refers only to the placement of atoms in a molecule and not the electrons, is equivalent to electron-pair geometry only when there are no lone electron pairs around the central atom.

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  BIOS Instant Notes in Physical Chemistry

  • Gavin Whittaker, Andy Mount, Matthew Heal(Authors)
  • 2000(Publication Date)
  • Taylor & Francis

    (Publisher)

  The Lewis theory of covalent bonding may be regarded as an elementary form of valence bond theory. It is nonetheless useful for describing covalent molecules with simple covalent bonds, and works successfully in describing the majority of, for example, organic compounds. Lewis theory recognizes both the free energy gains made in the formation of complete atomic electron shells, and the ability of atoms to achieve this state by sharing electrons. The sharing process is used as a description of covalent bonds.

  The atoms are firstly drawn so as to represent their relative arrangement, with electron pairs (marked as pairs of dots) between neighboring atoms to indicate a shared bonding electron pair. No attempt is made to describe the three-dimensional geometric shape of the molecule. Multiple bonds are represented by two or three electron pairs as appropriate (Fig. 1a ). Further electrons are added to each atom, so as to represent the non-bonding electrons and so complete the electron configuration of all the atoms (Fig. 1b ). It is customary to replace the bonding pairs of shared electrons with one line for each pair—each line representing a bond—with multiple lines representing multiple bonds (Fig. 1c ).

  Fig. 1. (a)–(c) Development of a Lewis bonding scheme for HCONH2 . Examples of (d) hyper-valency, and (e) resonance hybridization.

  Although main group elements tend to adopt inert gas configurations, which may be represented by eight valence electrons (an octet ), or two in the case of helium, a number of elements are energetically stable with incomplete octets . The most commonly cited example is boron, which is stable with six valence electrons as in BF3 , or Be with four as in BeCl2 . Larger elements are capable of hypervalency , where it is energetically favorable for more than eight valence electrons to be held in an expanded octet. Examples of this are PF5 (ten valence electrons) and XeF4 (twelve valence electrons) (Fig. 1d ).

  In many compounds, it is possible to devise two or more equivalent bonding schemes (canonical forms

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