Lewis Acids and Bases (5.6 - 5.11) A Lewis acid is an electron-pair acceptor Gilbert N. Lewis 1875-1946 • A Lewis acid has an empty orbital (or can rearrange electrons to create an empty orbital) that can “accept” an electron pair to form a new bond A Lewis base is an electron-pair donor • A Lewis base has an accessible lone pair of electrons that can be “donated” to a Lewis acid to form a new bond The product of a reaction between a LA and LB is a complex called an adduct, held together by a newly formed covalent (dative) bond From a MO perspective, we can consider the reactivity of frontier orbitals: Empty (acceptor) orbital of a Lewis acid → LUMO Filled (donor) orbital of a Lewis base → HOMO Energy → Adduct formation occurs via overlap: • The stability of adducts comes from a net lowering of energy of filled orbitals LUMO HOMO ⥯ ⥯ LA [LA←LB] complex or adduct LB • Creates new HOMO/LUMO pair (bonding/antibonding MOs) • The lower the energy of the empty orbital on the LA (LUMO of LA), the stronger the Lewis acid • The higher the energy of the filled orbital on the LB (HOMO of LB), the stronger the Lewis base F B F The resulting overlap creates a new covalent bond and a Lewis Acid-Base Adduct Lewis Acid F F Lewis Base H N H H The lone pair in the N sp3 orbital can donate into the empty 2p orbital of the B H F B N F F or H H H F B N F H H Bonds formed in a Lewis acid/base reaction are referred to as coordinate or dative bonds, and are sometimes written as arrows showing where the electrons came from (the Lewis base) Lewis Acid and Base Theory Explains Brønsted-Lowry Lewis Acid H+ The resulting overlap creates a new covalent bond and a Lewis AcidBase Adduct H H N H H Lewis Base The lone pair in the N sp3 orbital can donate into the empty 1s orbital of a proton (H+) H N H H • Note connection to nucleophiles (e-pair donor) and electrophiles (e-pair acceptors) which are used to discuss reaction rates (kinetics) • LA/LB highlight the equilibrium (thermodynamic) properties of reactions Examples of Lewis Acid-Base Chemistry Identify the Lewis acid and Lewis base. Determine the resulting LA-LB adduct. (a) AlCl3 + NCl3 → ? AlCl3, like BH3, has an incomplete octet and a vacant (3p) orbital that can accept a pair of electrons ← LEWIS ACID NCl3 has a full octet and a lone pair it can donate to form a new covalent bond ← LEWIS BASE Cl Cl Al Cl Cl + Cl N Cl Cl Cl Al N Cl Cl Cl Cl Cl Cl Cl Cl Al N Cl Cl Examples of Lewis Acid-Base Chemistry Identify the Lewis acid and Lewis base. Determine the resulting LA-LB adduct. (b) BCl3 + CH3OCH3 → ? BCl3 has an incomplete octet and a vacant (2p) orbital that can accept a pair of electrons ← LEWIS ACID The O atom in dimethyl ether has a full octet and two lone pairs it can donate to form a new covalent bond ← LEWIS BASE Cl Cl B Cl Cl + H 3C O CH3 H 3C B Cl Cl Cl O CH3 H 3C B Cl Cl O CH3 Examples of Lewis Acid-Base Chemistry Identify the Lewis acid and Lewis base. Determine the resulting LA-LB adduct. (c) Fe3+ + 6H2O → ? Fe3+ has vacant orbitals that can accept a pair of electrons → LEWIS ACID The O atom in water has a full octet and two lone pairs it can donate to form a new covalent bond ← LEWIS BASE OH2 Fe3+ + 6 H O H 2O H H 2O Fe OH2 3+ OH2 OH2 OH2 H 2O H 2O Fe H 2O 3+ OH2 OH2 Equilibria in LA-LB Reactions and Adduct Formation • Reactions between free Lewis acids and Lewis bases involve the formation of an adduct held together by a coordinate (= dative = covalent) bond (5.9a) • However, LA-LB reactions more typically involve equilibria with adducts “trading partners”, and equilibrium constants depending on the thermodynamics of the reaction For example, displacement reactions (5.9b): H 3C Cl O H 3C B Cl Cl + Cl N N B Cl Cl + H 3C O CH3 And “double displacement” (a.k.a. metathesis) reactions (5.9c): LA1-LB2 + LA2-LB2 ⇄ LA1-LB2 + LA2-LB2 “swapping partners…” (metathesis is Greek for “exchange”) Examples of Lewis Acids (electron-pair acceptors) (5.6) Need a vacant or “vacatable” energetically accessible orbital… 1. Molecules with incomplete octets of valence electrons F (a) trivalent Group 13 compounds: F B F + FF F B → good Lewis acids, so the adducts they form tend to be relatively stable: AlX4-, BX4-… → used as Lewis-acid catalysts (e.g., Friedel-Crafts, 5.7b) (b) carbocations or “carbenium” cations e.g., triphenylmethyl, [C(C6H5)3]+ C F F Examples of Lewis Acids (electron-pair acceptors) (5.6) Need a vacant or “vacatable” energetically accesible orbital… 1. Molecules with incomplete octets of valence electrons (c) a proton (H+) is a Lewis acid! → Lewis acid/base theory subsumes Brønsted-Lowry acid/base theory! + H Cl → a Brønsted-Lowry acid is Lewis acid-base adduct of the acid’s conjugate base and a proton → dissociation of a Brønsted-Lowry acid in water is a displacement reaction between two Lewis bases (conjugate base of the acid and water) competing for a single Lewis acid (H+) HA(aq) + H2O(l) ⇄ H3O+(aq) + A-(aq) Examples of Lewis Acids (electron-pair acceptors) (5.6) Need a vacant or “vacatable” energetically accessible orbital… 2. Metal cations can accept electron pairs from Lewis bases (e.g., solvents) to form coordination complexes → aqua acids, solvated cations OH2 Fe3+ + 6 H O H 2O H H 2O Fe OH2 3+ OH2 OH2 OH2 H 2O H 2O Fe H 2O 3+ OH2 OH2 Examples of Lewis Acids (electron-pair acceptors) (5.6) Need a vacant or “vacatable” energetically accessible orbital… 3. Molecules with complete octets that can rearrange their valence electrons to accept an electron pair → such molecules are said to be Lewis acidic O C O + OH- Similarly, SO2, SO3... O O C OH Examples of Lewis Acids (electron-pair acceptors) (5.7) Need a vacant or “vacatable” energetically accessible orbital… 4. Molecules that can “expand” their valence shell (or is simply large enough!) to accept another electron pair F F F Si F F + 2 F- F F Si F 2- F F → tendency of Group 14 compounds (e.g., Group 14 tetrahalides) to act as Lewis acids by becoming 5/6 coordinate e.g., SnCl4 + 2 Cl- → [SnCl6]2Note: lower valent Sn(II) halides can act as Lewis bases as well as Lewis acids Examples of Lewis Acids (electron-pair acceptors) (5.7) Need a vacant or “vacatable” energetically accessible orbital… 4. Molecules that can “expand” their valence shell (or is simply large enough!) to accept another electron pair Group 14 C Si Ge Sn 15 N P As Sb 16 O S Se Te 17 F Cl Br I → Group 15 pentahalides (PnX5, Pn = N, P, As, Sb…) of the heavier elements F Sb X X X Sb(III) F Sb F + 2HF F F Sb(V) F [H2F]+ F Sb F F F F - Examples of Lewis Acids (electron-pair acceptors) (5.7) Need a vacant or “vacatable” energetically accessible orbital… 5. Group 16 Lewis Acids Sulfur dioxide (like CO2) can act as a Lewis acid (or a LB!) O O S O O S O + N R R R R S N R R O Sulfur trioxide is a stronger Lewis acid (and much weaker LB) O O S O O O S O O + N R R R R O S N R R O Examples of Lewis Acids (electron-pair acceptors) (5.7) Need a vacant or “vacatable” energetically accessible orbital… 6. Group 17 diatomics • I2 and Br2 are mildly Lewis acidic • vacant σ* orbital is the “acceptor orbital” • violet I2 turns brown in Lewis-basic solvents I I + I I I I triiodide anion (I3-) I I I Polyhalide anions (17.10) • An I3- anion can similarly interact with other I2 molecules and form larger mononegative polyiodides of composition [(I2)nI-] • The triiodide anion is the most stable of the series • Cl3-, Br3-, BrI2- also known Figure 17.12 Focusing on the 3 x pz orbitals from three I atoms, we can look at the formation of the Lewis acid/base adduct: • 3 AOs form 3 MOs • Bonding, anti-bonding and non-bonding • Bond order: 2 electrons in a bonding MO shared by two bonds (1/2 bond per pair of atoms) I Figure 17.11 I I I I I I I I