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Chemistry LibreTexts

1: Chapter 1 Structure Determines Properties

  • Page ID
    • 1.1: Development of Chemical Bonding Theory
      Lewis Dot Symbols are a way of indicating the number of valence electrons in an atom. They are useful for predicting the number and types of covalent bonds within organic molecules. The molecular shape of molecules is predicted by Valence Shell Electron Pair Repulsion (VSEPR) theory. The shapes of common organic molecules are based on tetrahedral, trigonal planar or linear arrangements of electron groups.
    • 1.2: The Nature of Chemical Bonds - Valence Bond Theory
      Covalent bonds form as valence electrons are shared between two atoms. Lewis Structures and structural formulas are common ways of showing the covalent bonding in organic molecules. Formal charge describes the changes in the number of valence electrons as an atom becomes bonded into a molecule. If the atom has a net loss of valence electrons it will have a positive formal charge. If the atom has a net gain of valence electrons it will have a negative formal charge.
    • 1.3: 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.
    • 1.4: 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.
    • 1.5: 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.
    • 1.6: 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.
    • 1.7: 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 on one terminal oxygen can be delocalized by resonance through the other terminal oxygen.
    • 1.8: 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.
    • 1.9: 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.
    • 1.10: 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.
    • 1.11: Predicting Acid-Base Reactions from pKa Values
      pKa values can be used to predict the equilibrium of an acid-base reaction.  The equilibrium will favor the side with the weaker acid.
    • 1.12: Organic Acids and Organic Bases
      In the absence of pKa values, the relative strength of an organic acid can be predicted based on the stability of the conjugate base that it forms.  The acid that forms the more stable conjugate base will be the stronger acid.  The common factors that affect the conjugate base's stability are 1) the size and electronegativity of the the atom that has lost the proton, 2) resonance effects, 3) inductive effects, and 4) solvation effects.
    • 1.13: 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.
    • 1.14: Noncovalent Interactions Between Molecules
      In contrast to intramolecular forces, such as the covalent bonds that hold atoms together in molecules and polyatomic ions, intermolecular forces hold molecules together in a liquid or solid. Intermolecular forces are generally much weaker than covalent bonds.  The most common intermolecular forces in organic chemistry are from strongest to weakest are hydrogen bonds, dipole-dipole interactions, and London Dispersion (van der Waals) forces.
    • 1.15: Molecular Models
    • 1.1 Atoms, Electrons, and Orbitals