15.2 Lewis Acids and Bases

Lewis acids and bases expand our understanding of chemical reactions beyond proton transfer. These compounds interact through electron pair sharing, forming adducts and complex ions. This concept is crucial for grasping coordination chemistry and predicting molecular behavior.

The formation of Lewis acid-base complexes involves equilibrium processes, represented by formation constants. Understanding these interactions helps explain molecular geometries and bonding in various chemical systems, from simple compounds to complex metal coordination structures.

Lewis Acids and Bases

Lewis acids vs bases

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• Lewis acids act as electron pair acceptors have an incomplete octet or an empty orbital ($\text{BF}_3$, $\text{AlCl}_3$, $\text{Ag}^+$)
• Lewis bases function as electron pair donors possess a lone pair of electrons ($\text{NH}_3$, $\text{H}_2\text{O}$, $\text{OH}^-$)
• Interaction between Lewis acids and bases involves the Lewis base donating an electron pair to the Lewis acid forming a coordinate covalent bond resulting in the formation of an adduct or complex ion
• The octet rule often guides these interactions, as atoms tend to achieve a stable electron configuration

Formation of adducts and complex ions

• Adduct formation occurs when a Lewis acid and base combine to form a single species ($\text{BF}_3 + \text{NH}_3 \rightarrow \text{F}_3\text{B}:\text{NH}_3$) where $\text{BF}_3$ (Lewis acid) accepts an electron pair from $\text{NH}_3$ (Lewis base)
• Complex ion formation involves metal cations (Lewis acids) reacting with ligands (Lewis bases) to form complex ions ($\text{Ag}^+ + 2\text{NH}_3 \rightarrow [\text{Ag}(\text{NH}_3)_2]^+$) where $\text{Ag}^+$ (Lewis acid) accepts electron pairs from two $\text{NH}_3$ molecules (Lewis bases)
• Representation of adducts and complex ions utilizes Lewis structures to show the coordinate covalent bond(s) formed between the Lewis acid and base(s) with brackets and charges used to represent complex ions ($[\text{Cu}(\text{NH}_3)_4]^{2+}$, $[\text{Fe}(\text{CN})_6]^{3-}$)
• These interactions are fundamental to coordination chemistry, which studies the formation and properties of complex ions

Equilibrium in Lewis acid-base systems

• Formation constant ($K_f$) represents the equilibrium constant for the formation of an adduct or complex ion indicating the stability of the adduct or complex ion ($K_f = \frac{[\text{F}_3\text{B}:\text{NH}_3]}{[\text{BF}_3][\text{NH}_3]}$)
• Calculating equilibrium concentrations involves:
  1. Setting up an ICE table (Initial, Change, Equilibrium) using the balanced equation and initial concentrations
  2. Expressing the equilibrium concentrations in terms of the initial concentrations and the change in concentration (x)
  3. Substituting the equilibrium expressions into the $K_f$ expression and solving for x
  4. Calculating the equilibrium concentrations using the value of x
• Factors affecting the formation constant include the strength of the Lewis acid and base, steric factors (size and shape of the molecules), and solvent effects (polarity and donor/acceptor properties)

Bonding and Geometry in Lewis Acid-Base Interactions

• Bonding theory helps explain the formation of coordinate covalent bonds in Lewis acid-base reactions
• The molecular geometry of the resulting adducts or complex ions is influenced by the number and arrangement of electron pairs around the central atom
• Understanding these geometries is crucial for predicting the properties and reactivity of Lewis acid-base complexes