Chapter 2: Polar Covalent Bonds; Acids and Bases – Study Notes

Polar Covalent Bonds; Acids and Bases

Introduction

This chapter explores the nature of polar covalent bonds, the concept of electronegativity, the inductive effect, formal charges, resonance, and the fundamental theories of acids and bases (Brønsted-Lowry and Lewis). Understanding these concepts is essential for predicting molecular behavior and reactivity in organic chemistry.

Polar Covalent Bonds

Definition and Characteristics

• Polar covalent bonds occur when bonding electrons are attracted more strongly by one atom than by the other, resulting in an unequal sharing of electrons.

• This unequal sharing leads to bond polarity, where one atom acquires a partial negative charge (δ−) and the other a partial positive charge (δ+).

• Electron distribution between atoms in a polar covalent bond is not symmetrical.

• Bonds can be classified as:

  • Nonpolar covalent: Electronegativity difference < 0.5 (e.g., C–H bond)

  • Polar covalent: Electronegativity difference between 0.5 and 2.0 (e.g., C–O, C–X bonds)

  • Ionic: Electronegativity difference > 2.0

• Example: In H2C–Cl, the Cl atom is more electronegative and pulls electron density toward itself, making the bond polar.

Electronegativity (EN)

• Electronegativity is the intrinsic ability of an atom to attract shared electrons in a covalent bond.

• Differences in EN between atoms produce bond polarity.

• Fluorine (F) is the most electronegative element (EN = 4.0), while cesium (Cs) is among the least (EN = 0.7).

• Metals (left side of the periodic table) attract electrons weakly; nonmetals (right side, especially halogens) attract electrons strongly.

• EN of carbon (C) is 2.5.

Electronegativity Trends

• Electronegativity increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

• Color coding in periodic tables:

  • Red = Most electronegative

  • Green = Least electronegative

  • Yellow = Intermediate electronegativity

Bond Polarity and Dipole Moments

Bond Dipoles and Molecular Polarity

• Individual bonds can be polar, and the overall molecule can also be polar if the bond dipoles do not cancel out.

• Dipole moment (μ) is a measure of molecular polarity, resulting from the vector sum of individual bond dipoles and lone-pair contributions.

• Measured in Debye units (D).

• Example: Water (H2O) has a dipole moment of 1.85 D due to its bent shape and polar O–H bonds.

Symmetry and Dipole Moments

• In symmetrical molecules, the dipole moments of individual bonds may cancel each other, resulting in a nonpolar molecule (e.g., CO2, CH4).

• Asymmetrical molecules (e.g., H2O, NH3) have net dipole moments.

Formal Charge

Definition and Calculation

• Formal charge is the charge assigned to an atom in a molecule, assuming equal sharing of electrons in bonds.

• Formula: $\text{Formal charge} = (\text{Number of valence electrons in free atom}) - (\text{Number of bonds}) - (\text{Number of nonbonding electrons})$

• Helps identify reactive sites and resonance contributors.

• Example: In dimethyl sulfoxide (CH3SOCH3), oxygen has a formal negative charge, and sulfur has a formal positive charge.

Resonance

Resonance Structures and Hybrids

• Some molecules cannot be represented by a single Lewis structure; instead, they are described by resonance forms that differ only in the placement of π or nonbonding electrons.

• The actual structure is a resonance hybrid, a weighted average of all valid resonance forms.

• Resonance increases molecular stability by delocalizing electrons.

• Example: The acetate ion (CH3CO2−) has two resonance forms, with the negative charge delocalized over two oxygen atoms.

Rules for Drawing Resonance Forms

• Resonance forms must obey normal rules of valence.

• Only electrons (not atoms) are moved, typically shown with curved arrows.

• Resonance forms do not have to be equivalent.

• Resonance is possible in any system with a conjugated π system or adjacent lone pairs and π bonds.

Acids and Bases

Brønsted-Lowry Definition

• Brønsted-Lowry acid: Substance that donates a proton (H+).

• Brønsted-Lowry base: Substance that accepts a proton (H+).

• When an acid donates a proton, it forms its conjugate base; when a base accepts a proton, it forms its conjugate acid.

• General reaction: $\text{HA} + \text{B} \rightleftharpoons \text{A}^- + \text{HB}^+$

Acid and Base Strength

• Acidity constant (Ka): Measures the strength of an acid in water.

• Formula: $K_a = \frac{[\text{H}_3\text{O}^+][\text{A}^-]}{[\text{HA}]}$

• pKa: Negative logarithm of Ka: $pK_a = -\log K_a$

• Stronger acids have smaller pKa values; weaker acids have larger pKa values.

• There is an inverse relationship between the strength of an acid and its conjugate base.

Water as Acid and Base

• Water can act as both an acid and a base (amphoteric).

• Ion product constant of water: $K_w = [\text{H}_3\text{O}^+][\text{OH}^-]$ At 25°C, $K_w = 1.0 \times 10^{-14}$ and $pK_w = 14$.

Relative Strengths of Acids and Bases

• Common organic acids: Alcohols, carboxylic acids, phenols.

• Common organic bases: Amines, oxygen- and nitrogen-containing compounds with lone pairs.

• Table: Relative pKa Values (approximate)

  Compound Type	pKa Range
  Alkanes (R–H)	~50–60
  Alkenes (R=H)	~40–50
  Amines (R–NH2)	~35
  Alcohols (R–OH)	~15–20
  Carboxylic acids (R–COOH)	~5

Predicting Acid-Base Reactions Using pKa Values

• Acid-base reactions favor the formation of the weaker acid (higher pKa).

• To predict the direction of a reaction, compare the pKa values of the acids on both sides; the equilibrium favors the side with the weaker acid.

Organic Acids and Bases

Organic Acids

• Characterized by the presence of a positively polarized hydrogen atom.

• Two main types:

  • Hydrogen bonded to an electronegative oxygen atom (e.g., alcohols, carboxylic acids).

  • Hydrogen bonded to a carbon atom adjacent to a carbonyl group (O=C–C–H).

• Examples: Methanol (pKa = 15.56), Acetic acid (pKa = 4.76), Acetone (pKa = 19.3).

Organic Bases

• Contain an atom with a lone pair of electrons that can bond to H+.

• Amines (nitrogen-containing compounds derived from ammonia) are the most common organic bases.

• Oxygen-containing compounds can act as bases with strong acids or as acids with strong bases.

Mechanism of Acid-Base Reactions

• Acid-base reactions involve the transfer of a proton from the acid to the base.

• Curved arrows are used to show the movement of electron pairs during the reaction.

• Example: $\text{H}^+ + \text{NaC} \equiv \text{N} \rightarrow \text{HC} \equiv \text{N} + \text{Na}^+$

Lewis Acids and Bases

Lewis Definition

• Lewis acid: Electron pair acceptor.

• Lewis base: Electron pair donor.

• This definition encompasses the Brønsted-Lowry definition and includes many metal cations and electron-deficient compounds.

• Examples of Lewis acids: BF3, AlCl3, transition metal compounds (e.g., TiCl4, FeCl3).

• Examples of Lewis bases: Compounds with lone pairs such as alcohols, ethers, amines, and carbonyl compounds.

Curved Arrow Formalism

• Curved arrows indicate the movement of electron pairs from the Lewis base (donor) to the Lewis acid (acceptor).

• Used to depict the formation of acid-base complexes and other electron-pair transfer reactions.

Examples of Lewis Acids and Bases

• Neutral proton donors: H2O, HCl, HBr, HNO3, H2SO4

• Carboxylic acids, phenols, alcohols (as acids)

• Cations: Metal ions such as Al3+, Fe3+, Zn2+

• Metal compounds: AlCl3, TiCl4, FeCl3, ZnCl2

Compounds Acting as Both Acids and Bases

• Some compounds, such as water, alcohols, and amides, can act as both acids and bases depending on the reaction conditions.

Summary Table: Common Lewis Acids and Bases

Lewis Acids	Lewis Bases
BF3, AlCl3, TiCl4, FeCl3	H2O, NH3, R–OH, R–NH2, ethers, carbonyls
Metal cations (e.g., Zn2+, Fe3+)	Carboxylates, alkoxides, amides

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