Formal Charge

9 Key excerpts on "Formal Charge"

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  eBook - PDF

  Basic Concepts of Chemistry

  • Leo J. Malone, Theodore O. Dolter(Authors)
  • 2012(Publication Date)
  • Wiley

    (Publisher)

  Formal Charge is the charge that each atom in a molecule would have if the electrons in the bonds were divided equally between the two atoms. In other words, all bonds are treated as if they were non- polar. Formal Charge is a method of electron bookkeeping. It is not meant to imply that it is the actual charge on the atom in question. Formal Charge is calculated by subtracting the number of lone-pair electrons on the atom in question and half of the shared electrons from the group number. (The group number represents the number of valence electrons in a neutral atom.) For example, consider the structure above, which follows the octet rule. • The Formal Charge on the S is [6 (group number) - 1/2 * 8 (four bonded pairs) = +2]. • The two O atoms bonded to both S and H have Formal Charges of zero [i.e., 6 - 4 (two lone pairs) - 1/2 * 4 (two bonded pairs) = 0]. • The two O atoms bonded only to S each have a Formal Charge of -1 [i.e., 6 - 6 (three lone pairs) - 1/2 * 2 (one bonded pair) = -1]. The Formal Charge is represented as the charge in a circle, as follows. Zero for- mal charge is not shown. All the Formal Charges add to zero for a molecule or to the charge on an ion. Now consider the alternate structure. None of the atoms have any Formal Charge; that is, for the S, 6 - 1/2 * 12 = 0, and for the O atoms with double bonds, 6 - 4 (two lone pairs) -1/2 * 4 (two bonded pairs) = 0. The other two oxygen atoms also have no Formal Charge. Many chemists consider the structure with the least amount of Formal Charge to be the more favorable structure. The point remains controversial, however. Other sulfur and phosphorus compounds are often represented by favorable Formal Charge structures at the expense of the octet rule. For example, SO 2 and the PO 4 3- ion can be represented with a structure that follows the octet rule or has a favorable Formal Charge distribution that violates the octet rule.

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  Chemistry 2e

  • Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
  • 2019(Publication Date)
  • Openstax

    (Publisher)

  As we have seen, however, in some cases, there is seemingly more than one valid structure for a molecule. We can use the concept of Formal Charges to help us predict the most appropriate Lewis structure when more than one is reasonable. Calculating Formal Charge The Formal Charge of an atom in a molecule is the hypothetical charge the atom would have if we could redistribute the electrons in the bonds evenly between the atoms. Another way of saying this is that Formal Charge results when we take the number of valence electrons of a neutral atom, subtract the nonbonding electrons, and then subtract the number of bonds connected to that atom in the Lewis structure. Thus, we calculate Formal Charge as follows: We can double-check Formal Charge calculations by determining the sum of the Formal Charges for the whole structure. The sum of the Formal Charges of all atoms in a molecule must be zero; the sum of the Formal Charges in an ion should equal the charge of the ion. We must remember that the Formal Charge calculated for an atom is not the actual charge of the atom in the molecule. Formal Charge is only a useful bookkeeping procedure; it does not indicate the presence of actual charges. 332 7 • Chemical Bonding and Molecular Geometry Access for free at openstax.org EXAMPLE 7.6 Calculating Formal Charge from Lewis Structures Assign Formal Charges to each atom in the interhalogen ion Solution Step 1. We divide the bonding electron pairs equally for all I–Cl bonds: Step 2. We assign lone pairs of electrons to their atoms. Each Cl atom now has seven electrons assigned to it, and the I atom has eight. Step 3. Subtract this number from the number of valence electrons for the neutral atom: I: 7 – 8 = –1 Cl: 7 – 7 = 0 The sum of the Formal Charges of all the atoms equals –1, which is identical to the charge of the ion (–1).

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  Chemistry

  The Molecular Nature of Matter

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

    (Publisher)

  Formal Charges. The Formal Charge assigned to an atom in a Lewis structure (which usually differs from the actual charge on the atom) is calculated as the difference between the number of valence electrons of an isolated atom of the element and the number of electrons that “belong” to the atom because of its bonds to other atoms and its unshared valence electrons. The sum of the Formal Charges always equals the net charge on the molecule or ion. The most stable (lowest energy) Lewis structure for a molecule or ion is the one with Formal Charges closest to zero. Coordinate Covalent Bonding. For bookkeeping purposes, we sometimes single out a covalent bond whose electron pair originated from one of the two bonded atoms. Once formed, a coordinate covalent bond is no different from any other covalent bond. Draw and explain resonance structures Bonds to chemically equivalent atoms must be the same; they must have the same bond length and the same bond energy, which means they must involve the sharing of the same number of electron pairs. Typically, this occurs when it is necessary to form multiple bonds during the drawing of a Lewis structure. When alternatives exist for the location of multiple bond among two or more equivalent atoms, then each possible Lewis structure is actually a resonance structure or contribut- ing structure, and we draw them all. In drawing resonance structures, the relative locations of the nuclei must be identical in all. The average bond order is calculated from the total num- ber of bonds distributed over the equivalent bond loca- tions. Remember that none of the resonance structures corresponds to a real molecule, but their composite—the reso- nance hybrid—does approximate the actual structure of the molecule or ion. Classify organic compounds and identify functional groups Organic compounds are compounds with carbon atoms that are covalently bonded to other carbon atoms, hydrogen, oxygen, nitrogen, and other nonmetals.

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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)

  Formal Charge can help identify electron-rich and electron-deficient areas. However, a molecule may have more than one Lewis structure to consider (Section 1.3.2). Example: Methoxide anion, CH 3 O – The total number of valence electrons is 14; we get 6 from O, 4 from C, 1 from each of 3 H’s, and 1 for the minus charge. It took four bonds to connect the atoms, so the number of shared electrons is 8. We have used eight electrons, and there are six electrons remaining to be added as lone pairs to complete oxygen’s octet. All that is left to do is assign Formal Charge. Oxygen started with six valence electrons, and in this structure has one bond to it and six unshared electrons, so –1 must be its Formal Charge. A check shows that the shells are correct for all the atoms, all the valence electrons have been used, and that the sum of the Formal Charges equals the total charge. The final structure is: We can’t have a bottle of just anions. All compounds are charge balanced, so the methoxide anion above would have a cation to balance the negative charge. This is often a group 1A cation like Li + , Na + , or K + . These cations have the stable configuration of a noble gas, and therefore are just spectator ions. To simplify things, organic chemists often don’t bother writing the spectator ions because they don’t participate in the reaction. Example: Acetaldehyde, CH 3 CHO The total number of valence electrons is 18; we get 6 from the O, 4 from each of 2 C's, 1 from each of 4 H's. The six-bond skeleton shares 12 electrons. Another bond must be made because two adjacent atoms, C and O, have less than expected from the general bonding trends. This bond should go between those two atoms giving the structure: Seven bonds used 14 electrons; the 4 unshared valence electrons are added as lone pairs to complete the octet of oxygen.

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  Chemistry

  • Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson(Authors)
  • 2015(Publication Date)
  • Openstax

    (Publisher)

  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). 7.5 Strengths of Ionic and Covalent Bonds The strength of a covalent bond is measured by its bond dissociation energy, that is, the amount of energy required to break that particular bond in a mole of molecules. Multiple bonds are stronger than single bonds between the same atoms. The enthalpy of a reaction can be estimated based on the energy input required to break bonds and the energy released when new bonds are formed. For ionic bonds, the lattice energy is the energy required to separate one mole of a compound into its gas phase ions. Lattice energy increases for ions with higher charges and shorter distances between ions. Lattice energies are often calculated using the Born-Haber cycle, a thermochemical cycle including all 388 Chapter 7 | Chemical Bonding and Molecular Geometry This OpenStax book is available for free at http://cnx.org/content/col11760/1.9 of the energetic steps involved in converting elements into an ionic compound. 7.6 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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  eBook - ePub

  BIOS Instant Notes in Inorganic Chemistry

  • Tony Cox(Author)
  • 2004(Publication Date)
  • Taylor & Francis

    (Publisher)

  − isoelectronic to N, with a nonbonding pair). They are not always written on inorganic valence structures, but the idea is useful in judging the viability of a proposed structure. Some general principles are:

  •  structures without Formal Charges are preferred if possible; •  structures with Formal Charges outside the range 1 to 1 are generally unfavorable; •  negative Formal Charges should preferably be assigned to more electronegative atoms, positive charges to more electropositive atoms.

  Thus in N2 O (14), the structure with O− is probably more significant than that with N− . The BF molecule (15) is isoelectronic with CO but the corresponding triple-bonded structure appears very unlikely because it requires Formal Charges B2− and F2+ . The single-bonded form without charges may best describe the bonding.

  Formal Charge is very different from oxidation state, which is assigned by apportioning electrons in a bond to the more electronegative atom rather than equally (Topic B4 ). Both are artificial assignments, useful in their respective ways, but neither is intended as a realistic judgment of the charges on atoms.

  Limitations

  The model described in this section can be justified theoretically using the quantum mechanical valence bond theory. Nevertheless, there are many molecules where bonding cannot be described simply in terms of electron pairs localized between two atoms. Diborane is an example. The structure (16) as often drawn appears to have eight bonds and would therefore seem to need 16 valence electrons. In fact, there are only 12 and the molecule is sometimes described as electron deficient. Two pairs of electrons form three-center bonds each linking two boron atoms and a bridging hydrogen, as illustrated in the preferable way of drawing the valence structure in 16

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  Chemistry, 5th Edition

  • Allan Blackman, Steven E. Bottle, Siegbert Schmid, Mauro Mocerino, Uta Wille(Authors)
  • 2022(Publication Date)
  • Wiley

    (Publisher)

  In this case, each atom has a Formal Charge of zero, and so no redistribu- tion of electrons is required. Is our answer reasonable? Step 4 results in 10 electrons around the chlorine atom. However, chlorine is a third-row element and is therefore able to accommodate more than eight electrons. We have followed the rules and obtained the structure with Formal Charges of zero on every atom, so our answer is reasonable. As mentioned in relation to the Lewis structure of NOCl, it is also possible to include ionic interactions, which avoid the necessity to expand the octet. In the case of ClF 3 , this would add a positive charge to the Cl and a negative charge to one of the F atoms (with three possible resonance forms). PRACTICE EXERCISE 5.1 Determine the Lewis structure of sulfur tetrafuoride, SF 4 . FIGURE 5.12 There are three possible ways to minimise the Formal Charges in the nitrate ion; any of the three oxygen atoms can supply a pair of electrons. or O N ‒ N ‒ or O N N ‒ O N N O O O O O O O O O O O O O O O Note that some molecules and all ions will have preferred Lewis structures in which the for- mal charges are not all zero. Recall that in these cases, Formal Charges are more plausible if posi- tive charges are on the less electronegative atoms, while negative charges are on the more elec- tronegative atoms (e.g. the nitrate ion, NO 3 − , in figure 5.12, where the more electronegative O atoms, rather than the N atom, have a negative Formal Charge). It is also important to realise that Formal Charges are not the same as the partial charges induced as a result of bond polarity. For example, in ClF 3 , the chlorine atom would have a + charge, despite the fact that its Formal Charge is zero. Resonance structures In completing step 5 of the Lewis structure pro- cedure, there might be more than one way to minimise the Formal Charges on the atoms; that is, more than one Lewis structure is possible for a particular molecule or ion.

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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)

  with 120° angles between each pair and the central atom, and the other two form the apex of two pyramids, one above and one below the triangular plane shape in which three outside groups are placed in a flat triangle around a central atom with 120° angles between each pair and the central atom bond in which three pairs of electrons are shared between two atoms theory used to predict the bond angles in a molecule based on positioning regions of high electron density as far apart as possible to minimize electrostatic repulsion quantity having magnitude and direction Key Equations • Formal Charge = # valence shell electrons (free atom) − # one pair electrons − 1 2 # bonding electrons Summary 4.1 Ionic Bonding Atoms gain or lose electrons to form ions with particularly stable electron configurations. The charges of cations formed by the representative metals may be determined readily because, with few exceptions, the electronic structures Chapter 4 | Chemical Bonding and Molecular Geometry 241 of these ions have either a noble gas configuration or a completely filled electron shell. The charges of anions formed by the nonmetals may also be readily determined because these ions form when nonmetal atoms gain enough electrons to fill their valence shells. 4.2 Covalent Bonding Covalent bonds form when electrons are shared between atoms and are attracted by the nuclei of both atoms. In pure covalent bonds, the electrons are shared equally. In polar covalent bonds, the electrons are shared unequally, as one atom exerts a stronger force of attraction on the electrons than the other. The ability of an atom to attract a pair of electrons in a chemical bond is called its electronegativity. The difference in electronegativity between two atoms determines how polar a bond will be. In a diatomic molecule with two identical atoms, there is no difference in electronegativity, so the bond is nonpolar or pure covalent.

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  eBook - ePub

  Organic Reaction Mechanisms, Selected Problems, and Solutions

  Second Edition

  • William C. Groutas, Athri D. Rathnayake(Authors)
  • 2023(Publication Date)
  • CRC Press

    (Publisher)

  PART A

  Passage contains an image

  Minireview 1 Lewis structures

  A sound understanding of mechanistic organic chemistry requires a proficiency in writing Lewis structures. Without the ability to draw Lewis structures correctly and with facility, a student is so severely handicapped that he or she will ultimately resort to learning organic chemistry by rote (a tedious, frustrating, and minimally-successful endeavor). The importance of this will become apparent momentarily.

  A Lewis structure is a type of structural formula that shows the way in which the atoms are bonded together and depicts the bonding between atoms using pairs of non-bonded electrons (shown as dots) and bonded electrons (shown as dashes). In writing Lewis structures, the general approach outlined below should be followed.

  1. Determine the total number of valence electrons.

  For neutral molecules, this is simply accomplished by adding up the valence electrons of the individual atoms. In the case of ions, an electron is added for each negative charge (anions), and an electron is subtracted for each positive charge (cations). Recall that the number of valence electrons for an element corresponds to the group number of that element in the periodic table. For example, nitrogen has five valence electrons (nitrogen is located in group five of the periodic table), fluorine has seven valence electrons (fluorine is located in group seven). etc.

  Formal Charge = X -Y- Z

  where, X= number of valence electrons (of atom under Y= number of non-bonded electrons, and Z= half the number of bonded electrons

  CCI4

   1 C 1 X 4 = 4  4 CI 4 X 7= 28  32 Total number of valence electrons

  1. Connect the atoms in the given molecular formula using single lines (dashes).
     It’s helpful to remember that in the case of polyatomic molecules or ions, the atom of lower electronegativity is typically the central atom. Recall that electronegativity follows the order F > O > CI, N > Br > C, H.

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