## 9.2 The Lewis Concept of Acids and Bases The Arrhenius and Brønsted-Lowry theories of acids and bases, while fundamental, have limitations in explaining the acid-base behavior of certain substances. For instance, compounds like $\mathrm{SO_2}$, $\mathrm{CO_2}$, $\mathrm{CaO}$, and $\mathrm{BF_3}$ act as acids or bases despite not having the ability to donate or accept protons. To address these limitations, G.N. Lewis proposed a more general acid-base theory in 1923, focusing on electron pair transfer. ### Definitions **Lewis Acid:** A substance that can *accept a pair of electrons* to form a coordinate covalent bond. Lewis acids are often referred to as **electron pair acceptors**. **Lewis Base:** A substance that can *donate a pair of electrons* to form a coordinate covalent bond. Lewis bases are often referred to as **electron pair donors**. In a Lewis acid-base reaction, a **coordinate covalent bond** (also known as a dative bond) is formed. This is a covalent bond in which both shared electrons originate from one atom (the Lewis base). The resulting product is often called an **Adduct**. ### Illustrative Example: Boron Trifluoride and Ammonia Consider the reaction between boron trifluoride ($\mathrm{BF_3}$) and ammonia ($\mathrm{NH_3}$): $$ \mathrm{BF_3} + \mathrm{NH_3} \rightarrow \mathrm{F_3B \leftarrow NH_3} $$ The bond formed between boron and nitrogen is a coordinate covalent bond, as the electron pair is donated by nitrogen. Let's analyze this reaction based on Lewis's definitions: 1. **Which species is donating an electron pair?** $\mathrm{NH_3}$ (specifically, the nitrogen atom in ammonia) 2. **Which species is accepting an electron pair?** $\mathrm{BF_3}$ (specifically, the boron atom in boron trifluoride) 3. **Which species is a Lewis acid?** $\mathrm{BF_3}$ (electron pair acceptor) 4. **Which species is a Lewis base?** $\mathrm{NH_3}$ (electron pair donor) ### Gas-Phase Neutralization Reactions The Lewis theory can effectively explain gas-phase neutralization reactions, which are not typically covered by Arrhenius or Brønsted-Lowry theories. For example, the reaction between hydrogen chloride ($\mathrm{HCl}$) gas and ammonia ($\mathrm{NH_3}$) gas forms ammonium chloride ($\mathrm{NH_4Cl}$): $$ \mathrm{HCl(g)} + \mathrm{NH_3(g)} \rightarrow \mathrm{NH_4Cl(s)} $$ In this reaction, the nitrogen atom in ammonia donates an electron pair to the hydrogen atom in $\mathrm{HCl}$ (which then leaves as $\mathrm{Cl^-}$). Therefore, $\mathrm{NH_3}$ is the Lewis base (electron pair donor), and $\mathrm{HCl}$ (specifically, the $\mathrm{H}$ atom acting as $\mathrm{H^+}$) is the Lewis acid (electron pair acceptor). This theory extends the concept by focusing on electron pair transfer rather than proton transfer. ### Characteristics of Lewis Acids and Bases **Lewis Acids:** - The central atom or atom of the Lewis acid typically has a *deficiency of an electron pair* (an incomplete octet). - It must have an empty orbital to accommodate an unshared electron pair. - Examples include species like $\mathrm{BF_3}$, $\mathrm{AlCl_3}$, $\mathrm{H^+}$, metal ions ($\mathrm{Fe^{3+}}$, $\mathrm{Cu^{2+}}$), and molecules with multiple bonds where atoms can expand their octet ($\mathrm{CO_2}$, $\mathrm{SO_2}$). **Lewis Bases:** - The central atom of a Lewis base usually has a *complete octet* and possesses *one or more unshared electron pairs* (lone pairs). - Hence, the base has the ability to donate an unshared electron pair. - Examples include species like $\mathrm{NH_3}$, $\mathrm{H_2O}$, $\mathrm{OH^-}$, $\mathrm{Cl^-}$, and organic molecules with lone pairs on heteroatoms (e.g., ethers, amines). ## Worked Examples ### Example 9.1: Classifying Substances as Lewis Acids or Lewis Bases Identify the Lewis acid and Lewis base in the following reaction: $$ \mathrm{BF_3} + \mathrm{NH_3} \rightarrow \mathrm{F_3B \leftarrow NH_3} $$ #### Problem Solving Strategy 1. **Draw Electronic Structures:** Sketch the Lewis structures of both species involved in the reaction. For help with electronic arrangements, see . 2. **Identify Electron Pair Donor (Lewis Base):** Look for the species that has a lone pair of electrons or a negative charge. Such a species typically has a complete octet and thus can donate an electron pair. This species is the Lewis base. 3. **Identify Electron Pair Acceptor (Lewis Acid):** Look for the species that can accommodate an electron pair (e.g., has an incomplete octet or an empty orbital). This species can accept an electron pair and is thus the Lewis acid. #### Solution 1. **Electronic Structures:** For $\mathrm{NH_3}$: Nitrogen has one lone pair of electrons and forms three single bonds with hydrogen atoms. For $\mathrm{BF_3}$: Boron forms three single bonds with fluorine atoms and has only six valence electrons (an incomplete octet). 2. **Identify Lewis Base:** $\mathrm{NH_3}$ has a lone pair on the nitrogen atom. This lone pair can be donated. Therefore, $\mathrm{NH_3}$ is an electron pair donor, making it a **Lewis base**. 3. **Identify Lewis Acid:** Boron in $\mathrm{BF_3}$ has only six valence electrons (three electron pairs), meaning it has an incomplete octet. It needs an electron pair to complete its octet. Hence, $\mathrm{BF_3}$ is an electron pair acceptor, making it a **Lewis acid**. ## Possible Questions/Answers ### Concept Assessment Exercise 9.1 Identify the Lewis acid and the Lewis base in the following examples. 1. $$ \mathrm{Cl^-} + \mathrm{AlCl_3} \rightarrow [\mathrm{AlCl_4}]^- $$ **Q:** Which is the Lewis acid and which is the Lewis base? **A:** $\mathrm{Cl^-}$ has lone pairs and a negative charge, making it an electron pair donor, thus a **Lewis base**. $\mathrm{AlCl_3}$ has aluminum with an incomplete octet (6 electrons), making it an electron pair acceptor, thus a **Lewis acid**. 2. $$ \mathrm{H^+} + \mathrm{OH^-} \rightarrow \mathrm{H_2O} $$ **Q:** Which is the Lewis acid and which is the Lewis base? **A:** $\mathrm{H^+}$ is electron-deficient (no electrons), making it an electron pair acceptor, thus a **Lewis acid**. $\mathrm{OH^-}$ has lone pairs and a negative charge, making it an electron pair donor, thus a **Lewis base**.