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2: Polar Covalent Bonds; Acids and Bases

  • Page ID
    162000
  • Chapter Objectives

    This chapter provides a review of the more advanced material covered in a standard introductory chemistry course through a discussion of the following topics:

    • the use of electronegativity to determine bond polarity, and the application of this knowledge to determine whether a given molecule possesses a dipole moment.
    • the drawing and interpretation of organic chemical structures.
    • the concept and determination of formal charge.
    • resonance and drawing of resonance forms
    • the Brønsted-Lowry and Lewis definitions of acids and bases, acidity constants and acid-base reactions.
    • the use of ball-and-stick molecular models.

    • 2.1: Polar Covalent Bonds - Electronegativity
      Because the tendency of an element to gain or lose electrons is so important in determining its chemistry, various methods have been developed to quantitatively describe this tendency. The most important method uses a measurement called electronegativity, defined as the relative ability of an atom to attract electrons to itself in a chemical compound.
    • 2.2: Polar Covalent Bonds - Dipole Moments
      Mathematically, dipole moments are vectors; they possess both a magnitude and a direction. The dipole moment of a molecule is therefore the vector sum of the dipole moments of the individual bonds in the molecule. If the individual bond dipole moments cancel one another, there is no net dipole moment.
    • 2.3: Formal Charges
      A formal charge is the charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity.
    • 2.4: Resonance
      Resonance structures are a set of two or more Lewis Structures that collectively describe the electronic bonding a single polyatomic species including fractional bonds and fractional charges. Resonance structure are capable of describing delocalized electrons that cannot be expressed by a single Lewis formula with an integer number of covalent bonds.
    • 2.5: Rules for Resonance Forms
      The above resonance structures show that the electrons are delocalized within the molecule and through this process the molecule gains extra stability. Ozone with both of its opposite formal charges creates a neutral molecule and through resonance it is a stable molecule. The extra electron that created the negative charge one terminal oxygen can be delocalized by resonance through the other terminal oxygen.
    • 2.6: Drawing Resonance Forms
      Resonance structures are used when one Lewis structure for a single molecule cannot fully describe the bonding that takes place between neighboring atoms relative to the empirical data for the actual bond lengths between those atoms. The net sum of valid resonance structures is defined as a resonance hybrid, which represents the overall delocalization of electrons within the molecule. A molecule that has several resonance structures is more stable than one with fewer.
    • 2.7: Acids and Bases - The Brønsted-Lowry Definition
      In 1923, chemists Johannes Brønsted and Martin Lowry independently developed definitions of acids and bases based on compounds abilities to either donate or accept protons (H+ ions). Here, acids are defined as being able to donate protons in the form of hydrogen ions; whereas bases are defined as being able to accept protons. This took the Arrhenius definition one step further as water is no longer required to be present in the solution for acid and base reactions to occur.
    • 2.8: Acid and Base Strength
      The relative acidity of different compounds or functional groups – in other words, their relative capacity to donate a proton to a common base under identical conditions – is quantified by a number called the dissociation constant, abbreviated Ka. The common base chosen for comparison is water.
    • 2.9: Predicting Acid-Base Reactions from pKa Values
    • 2.10: Organic Acids and Organic Bases
    • 2.11: Acids and Bases - The Lewis Definition
      A broader definition is provided by the Lewis theory of acids and bases, in which a Lewis acid is an electron-pair acceptor and a Lewis base is an electron-pair donor. This definition covers Brønsted-Lowry proton transfer reactions, but also includes reactions in which no proton transfer is involved.
    • 2.12: Molecular Models
    • 2.13: Polar Covalent Bonds; Acids and Bases (Summary)

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