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3: Chemical Formulas and Bonding

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    • 3.1 An Atomic-Level Perspective of Ionic and Covalent Compounds
      Metals (particularly those in groups IA and IIA) tend to lose the number of electrons that would leave them with the same number of electrons as in the preceding noble gas in the periodic table. By this means, a positively charged ion is formed. Similarly, nonmetals (especially those in groups VIA and VIIA, and, to a lesser extent, those in Group VA) can gain the number of electrons needed to provide atoms with the same number of electrons as in the next noble gas in the periodic table.
    • 3.2 Composition of Compounds
      A chemical formula is used to express the structure of a molecule. The formula tells which elements and how many of each element are present in a compound. Formulas are written using the elemental symbol of each atom and a subscript to denote the number of elements. Parentheses are used in chemical formulas to provide information about specifically bonded groups (polyatomic ions, alkyl groups, etc.) in the compound. A mid-line dot is used for compounds that contain water in their crystals.
    • 3.3 Chemical Bonds
      Chemical bonds form when electrons can be simultaneously close to two or more nuclei, but beyond this, there is no simple, easily understood theory that would not only explain why atoms bind together to form molecules, but would also predict the three-dimensional structures of the resulting compounds as well as the energies and other properties of the bonds themselves.
    • 3.4 Ionic Compounds: Formulas and Names
      Chemists use nomenclature rules to clearly name compounds. Ionic and molecular compounds are named using somewhat-different methods. Binary ionic compounds typically consist of a metal and a nonmetal. The name of the metal is written first, followed by the name of the nonmetal with its ending changed to –ide. For example, K2O is called potassium oxide. If the metal can form ions with different charges, a Roman numeral in parentheses follows the name of the metal to specify its charge.
    • 3.5 Covalently-Bonded Species: Formulas and Names
      Covalently-bonded compounds are named by stating the name of the less electronegative element first, followed by the name of the more electronegative element, which has been modified by replacing the last few letters of the element name withe the letters"-ide". In addition, the number of atoms of each element is noted using Greek numerical prefixes (although "mono-" is never used for the first element.) Binary compounds containing hydrogen are named as if they were ionically-bonded compounds.
    • 3.6 Electronegativity and Bond Polarity
      Bond polarity and ionic character increase with an increasing difference in electronegativity. The electronegativity (χ) of an element is the relative ability of an atom to attract electrons to itself in a chemical compound and increases diagonally from the lower left of the periodic table to the upper right. The Pauling electronegativity scale is based on measurements of the strengths of covalent bonds between different atoms, whereas the Mulliken electronegativity of an element is the average
    • 3.7 Lewis Structures
      Lewis dot symbols provide a simple rationalization of why elements form compounds with the observed stoichiometries. A plot of the overall energy of a covalent bond as a function of internuclear distance is identical to a plot of an ionic pair because both result from attractive and repulsive forces between charged entities. Lewis structures are an attempt to rationalize why certain stoichiometries are commonly observed for the elements of particular families.
    • 3.8 Resonance and Formal Charge Revisited
      Some molecules have two or more chemically equivalent Lewis electron structures, called resonance structures. Resonance is a mental exercise and method within the Valence Bond Theory of bonding that describes the delocalization of electrons within molecules. These structures are written with a double-headed arrow between them, indicating that none of the Lewis structures accurately describes the bonding but that the actual structure is an average of the individual resonance structures.
    • 3.9 Exceptions to the Octet Rule
      Following the Octet Rule for Lewis Dot Structures leads to the most accurate depictions of stable molecular and atomic structures and because of this we always want to use the octet rule when drawing Lewis Dot Structures. There are three exceptions: (1) When there are an odd number of valence electrons, (2) When there are too few valence electrons, and (3) when there are too many valence electrons
    • 3.10 Shapes of Molecules - VSEPR Theory and Valence Bond Theory
      Molecular shapes can be predicted from the number of electron pairs attached to each central atom in a Lewis dot structure. VSEPR Theory gives a simple description of the shapes to expect, but offers little in explanation for how the molecular shapes come about. Valence bonding theory uses a process called hybridization, in which atomic orbitals on each atom are combined mathematically to produce  hybrid atomic orbitals. These hybrid orbitals are for forming sigma bonds and holding lone pairs.
    • 3.11 Practice Problems

    3: Chemical Formulas and Bonding is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by LibreTexts.

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