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2.S: Polar Covalent Bonds; Acids and Bases (Summary)

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    2.1 Polar Covalent Bonding

    • The difference in electronegativity values of two atoms determines whether the bond between those atoms is classified as either ionic, polar covalent, or non-polar covalent.
    • Ionic bonds result from large differences in electronegativity values, such as that between a metal and non-metal atom (Na and Cl).
    • Covalent bonding generally results when both atoms are non-metals, like C, H, O, N and the halides.
    • When both atoms the same and/or have the same electronegativity value, then the bonding electrons are shared equally and the bond is classified as non-polar covalent.
    • Polar covalent bonds occur when the difference in electronegativity values is small, and the bonding electrons are not shared equally.

    2.2 Polar Covalent Bonding - Dipole Moments

    • The molecular dipole moment is the sum of all the bond dipoles within a molecule and depends on both the molecular geometry and the bond polarity.
    • Molecules that contain no polar bonds, like CH4, and/or completely symmetrical molecules, like CO2, generally have no net dipole moment.
    • Asymmetrical molecules that contain bonds of different polarities or non-bonding lone pairs typically have a molecular dipole moment.

    2.3 Formal Charges

    • Formal Charge compares how many valence electrons surround a free atom versus how many surround that same type of atom bonded with a molecule or ion.
    • Formal charge can be calculated using the equation

      Formal Charge = (# of valence electrons in free atom) - (# of lone-pair electrons) - (1/2 # of bond pair electrons) Eqn. 2.3.1

    • Formal charges of zero generally represent the most stable structures.
    • These bonding patterns for the atoms commonly found in organic molecules result in a formal charge of zero
      • Carbon - 4 bonds, no lone pairs
      • Hydrogen - 1 bond, no lone pairs
      • Nitrogen - 3 bonds, 1 lone pair
      • Oxygen - 2 bonds, 2 lone pairs
      • Halogens - 1 bond, 3 lone pairs.

    2.4 Resonance

    • Resonance Theory is often used when the observed chemical and physical properties of a molecule or ion cannot be adequately described by a single Lewis Structure. A classic example is the benzene molecule, C6H6. The Lewis Structure of benzene could be drawn in two different ways. Both structures have alternating double bond and single bonds between the carbons. The only difference is the location of the pi bonds.

    If these structures are correct, then the benzene molecule should have two different C-C bond lengths and bond energies, corresponding to a C-C single bond and to a C=C double bond. However, analysis shows that benzene contains only one type of carbon-carbon bond and it's bond length and energy are half between those of a single bond and double bond. Resonance theory states that benzene exists as the "average" of the two structures called a resonance hybrid, in which the six pi electrons delocalized over all six carbon atoms. Each C-C bond in benzene would be the average of a single bond and double bond or a "bond and a half". Dashed lines are often used to show type of "partial" bonding in a resonance hybrid of benzene

    2.5 Rules for Resonance Forms

    • The rules for estimating stability of resonance structures are
      • The resonance form in which all atoms have complete valence shells is more stable.
      • The greater the number of covalent bonds, the greater the stability since more atoms will have complete octets
      • The structure with the least number of formal charges is more stable
      • The structure with the least separation of formal charges is more stable
      • A structure with a negative charge on the more electronegative atom will be more stable
      • Positive charges on the least electronegative atom (most electropositive) is more stable
      • Resonance forms that are equivalent have no difference in stability and contribute equally. (eg. benzene)
    • If these rules are applied to the two Lewis Structures of benzene, the result would be that both structures will have the same relative stability and will both contribute equally to the character of the resonance hybrid.

    2.6 Drawing Resonance Forms

    • In resonance structures, the electrons are able to move to help stabilize the molecule. This movement of the electrons is called delocalization.
    • The rules for drawing resonance structures are:
      • Resonance structures should have the same number of electrons, do not add or subtract any electrons. (You can check the number of electrons by counting them)
      • All resonance structures must follow the rules of writing Lewis Structures.
      • The hybridization of the structure must stay the same.
      • The skeleton of the structure can not be changed (only the electrons move).
      • Resonance structures must also have the same amount of lone pairs.

    2.7 Acids and Bases - The Brønsted-Lowry Definition

    • A Brønsted-Lowry acid is a proton (H+) donor and a Brønsted-Lowry base is a proton acceptor.

    2.8 Acid and Base Strength

    • The strength of Brønsted-Lowry acids is measured indicated by its pKa value. The lower the pKa - the stronger the acid.
    • A strong acid will have a weak conjugate base. A strong base will have a weak conjugate acid.

    2.9 Predicting Acid-Base Reactions from pKa Values

    • The equilibrium of an acid-base reaction favors the formation of weaker acids from stronger acids. To predict the direction of the equilibrium, identify Brønsted-Lowry acid on each side of the reaction. Assign/look up pKa values for each acid. The equilibrium will favor the side that has the weakest acid (the highest pKa).

    2.10 Organic Acids and Organic Bases

    • Organic acids are stronger when the conjugate base that is formed upon loss of a proton is more stable.
    • Some factors that effect the stability of the conjugate base (often an anion) are the anionic atom's size and electronegativity, resonance effects, inductive effects, and solvation.

    2.11: Acids and Bases - The Lewis Definition

    • A Lewis acid is a lone pair acceptor and a Lewis base is a lone pair donor.

    2.12: Non-covalent Interactions between Molecules

    • Non-covalent Interactions, also known as Intermolecular Forces, significantly effect the physical properties of organic molecules. Hydrogen bonding is the most important of these interactions, but others include ion-dipole, dipole-dipole, and London Dispersion Forces.

    2.13: Molecular Models

    Skills to Master

    • Skill 2.1 Predict whether a bond is ionic, polar covalent, or non-polar covalent based on the position of the atoms in the periodic table.
    • Skill 2.2 Identify the partial positive and partial negative atoms of a polar covalent bond based on relative electronegativity.
    • Skill 2.3 Determine the dipole moment of a molecule based on molecular geometry and bond polarity.
    • Skill 2.4 Identify the chemicals in a reaction as Brønsted-Lowry acids or bases, and conjugate acids and bases.
    • Skill 2.5 Predict the products of an acid-base reaction.
    • Skill 2.6 Use pKa values to predict the equilibrium direction of an acid-base reaction.
    • Skill 2.7 Predict the relative strength of an organic acid by examining the stability of the conjugate base.
    • Skill 2.8 Use molecular structure and analysis of intermolecular forces to rank a series of organic molecules with respect to physical properties like melting point and boiling point.
    • Skill 2.9 Identify the chemicals in a reaction as Lewis acids or bases.

    Memorization Tasks (MT)

    MT 2.1 Memorize that the C-H bond is considered to be non-polar.

    MT 2.2 Memorize the common bonding patterns for C, H, N, O and the halogens that have a zero formal charge.

    MT 2.3 Memorize the factors that affect the relativity stability of conjugate bases.


    • Dr. Kelly Matthews (Professor of Chemistry, Harrisburg Area Community College)