🥼Organic Chemistry Unit 2 – Polar Covalent Bonds; Acids and Bases

Key Concepts and Definitions

• Polar covalent bonds form when electrons are unequally shared between atoms due to differences in electronegativity
• Electronegativity measures an atom's ability to attract electrons in a chemical bond
• Dipole moment quantifies the degree of polarity in a bond or molecule, represented by the Greek letter mu ($\mu$)
• Acids are proton ($H^+$) donors, while bases are proton acceptors in chemical reactions
  • Arrhenius acids dissociate in water to produce $H^+$ ions (hydrochloric acid, $HCl$)
  • Arrhenius bases produce $OH^-$ ions in aqueous solution (sodium hydroxide, $NaOH$)
• Conjugate acid-base pairs consist of a species and its corresponding proton-transfer product ($NH_4^+$ and $NH_3$)
• Amphoteric substances can act as both acids and bases depending on the reaction environment (water, $H_2O$)

Electronegativity and Bond Polarity

• Electronegativity values are assigned to elements based on the Pauling scale, ranging from 0.7 (cesium) to 4.0 (fluorine)
• The greater the electronegativity difference between bonded atoms, the more polar the covalent bond
  • Nonpolar covalent bonds have an electronegativity difference of less than 0.4 ($C-H$ bond)
  • Polar covalent bonds have an electronegativity difference between 0.4 and 1.7 ($C-O$ bond)
• Partial charges ($\delta+$ and $\delta-$) develop on atoms in polar covalent bonds due to unequal electron sharing
• Bond dipole moments are vector quantities with magnitude and direction, pointing from the positive to the negative end of the bond
• Molecular polarity depends on the geometry and arrangement of bond dipoles within the molecule ($CO_2$ is nonpolar, $H_2O$ is polar)

Formation of Polar Covalent Bonds

• Polar covalent bonds form when valence electrons are attracted more strongly by one atom than the other
• The more electronegative atom gains a partial negative charge ($\delta-$), while the less electronegative atom acquires a partial positive charge ($\delta+$)
• Electron density is shifted towards the more electronegative atom, resulting in an unequal distribution of charge
  • In the $H-Cl$ bond, the electron density is shifted towards the more electronegative chlorine atom
• The extent of electron sharing depends on the electronegativity difference between the bonded atoms
• Polar covalent bonds are intermediate between nonpolar covalent bonds and ionic bonds

Properties of Polar Molecules

• Polar molecules have an uneven distribution of charge, with distinct positive and negative regions
• The dipole moment of a molecule is the sum of its bond dipole moments, taking into account molecular geometry
  • Molecules with symmetric geometries ($CO_2$, $BF_3$) can be nonpolar despite having polar bonds
• Polar molecules interact through dipole-dipole forces, aligning their positive and negative ends
• Polar molecules are attracted to charged species and other polar molecules ($H_2O$ molecules in liquid water)
• Polar molecules tend to have higher melting and boiling points compared to nonpolar molecules of similar size due to stronger intermolecular forces
• Polar molecules are often soluble in polar solvents (ethanol in water) and insoluble in nonpolar solvents (oil in water)

Introduction to Acids and Bases

• Acids and bases are crucial in many chemical reactions and biological processes (digestion, blood pH regulation)
• Acids taste sour (citric acid in lemons), while bases taste bitter and feel slippery (baking soda, $NaHCO_3$)
• Acids and bases can be classified as strong or weak depending on their extent of ionization in aqueous solution
  • Strong acids and bases completely ionize ($HCl$, $NaOH$), while weak acids and bases partially ionize ($CH_3COOH$, $NH_3$)
• The pH scale measures the acidity or basicity of a solution, ranging from 0 (strongly acidic) to 14 (strongly basic)
  • Neutral solutions have a pH of 7 (pure water at 25°C)
• Acid-base indicators change color depending on the pH of the solution (phenolphthalein, litmus paper)

Brønsted-Lowry Theory

• The Brønsted-Lowry theory defines acids as proton donors and bases as proton acceptors
• Proton transfer occurs from the acid to the base, forming a conjugate acid-base pair
  • In the reaction $HCl + H_2O \rightarrow H_3O^+ + Cl^-$, $HCl$ is the acid, $H_2O$ is the base, $H_3O^+$ is the conjugate acid, and $Cl^-$ is the conjugate base
• The strength of an acid is related to the strength of its conjugate base, and vice versa (strong acid $\rightarrow$ weak conjugate base)
• Amphoteric substances can act as either acids or bases, depending on the reaction partner ($H_2O$ with $HCl$ or $NH_3$)
• The leveling effect occurs when a strong acid or base is dissolved in a solvent that is a weaker acid or base ($HCl$ in $H_2O$)

Lewis Acid-Base Theory

• The Lewis theory defines acids as electron pair acceptors and bases as electron pair donors
• Lewis acids have vacant orbitals that can accept electron pairs ($BF_3$, $AlCl_3$), while Lewis bases have lone pairs that can be donated ($NH_3$, $H_2O$)
• Lewis acid-base reactions involve the formation of a coordinate covalent bond, with both electrons coming from the Lewis base
  • In the reaction $BF_3 + NH_3 \rightarrow F_3B\leftarrow NH_3$, $BF_3$ is the Lewis acid and $NH_3$ is the Lewis base
• Lewis acids and bases can be neutral molecules or ions ($Cu^{2+}$, $OH^-$)
• The Lewis theory encompasses a wider range of species compared to the Brønsted-Lowry theory (metal cations, $CO_2$)

Acid-Base Reactions and Equilibria

• Acid-base reactions involve the transfer of protons between species, reaching a state of equilibrium
• The acid dissociation constant ($K_a$) and base dissociation constant ($K_b$) quantify the strength of acids and bases, respectively
  • Larger $K_a$ values indicate stronger acids, while larger $K_b$ values indicate stronger bases
• The acid and base dissociation constants are related by $K_w = K_a \times K_b$, where $K_w$ is the ion product constant of water ($1.0 \times 10^{-14}$ at 25°C)
• The pH of a solution can be calculated using the negative logarithm of the hydrogen ion concentration: $pH = -\log[H^+]$
• Titration is a technique used to determine the concentration of an acid or base by reacting it with a standard solution of known concentration
  • The equivalence point is reached when the acid and base have completely reacted, often indicated by a color change in the presence of an indicator (phenolphthalein)
• Buffer solutions resist changes in pH upon the addition of small amounts of acid or base, maintaining a relatively constant pH
  • Buffers contain a weak acid and its conjugate base (acetic acid and sodium acetate) or a weak base and its conjugate acid (ammonia and ammonium chloride)