Acids and Bases Introduction

### Introduction to Acids and Bases #### Definitions of Acids and Bases - Three definitions: - Arrhenius - Bronsted-Lowry - Lewis - Definitions progressed historically and in completeness. - Lewis's definition is the most complete. ##### Arrhenius Definition - Deals only with aqueous solutions. - Acid: Increases the $H_3O^+$ concentration in water. - $HCl$ (hydrochloric acid) is a classic example. - $HCl + H_2O \rightarrow H_3O^+ + Cl^-$ - Base: Increases the $OH^-$ concentration. - $NaOH$ (sodium hydroxide) is a classic example. - $NaOH \rightarrow Na^+ + OH^-$ - In aqueous solution, equal concentrations of $H_3O^+$ and $OH^-$ exist. ##### Bronsted-Lowry Definition - Can use any solvent or even the gas phase. - Acid: $H^+$ donor (proton donor). - Base: $H^+$ acceptor (proton acceptor). - Example in water: - $HCl + H_2O \rightarrow H_3O^+ + Cl^-$ - $HCl$ is the acid (proton donor). - $H_2O$ is the base (proton acceptor). - Example in the gas phase: - $HCl + NH_3 \rightarrow NH_4^+ + Cl^-$ - $HCl$ is the acid. - $NH_3$ is the base. - Most useful definition for identifying acids and bases. ##### Lewis Definition - Deals with electrons rather than protons. - Acid: Electron acceptor (electron pair acceptor). - Base: Electron donor (electron pair donor). - Making a new bond is key in an acid-base reaction. - If this bond is made to a hydrogen ion ($H^+$), it qualifies as both a Lewis and Bronsted-Lowry acid-base reaction. - If the new bond is made to something other than hydrogen, only Lewis's definition applies. - Example:
$HCl + NH_3 \rightarrow NH_4^+ + Cl^-$
\* $NH_3$ donates electrons (Lewis base).
\* $HCl$ accepts electrons (Lewis acid). - Boron compounds (e.g., $BF_3$) often act as Lewis acids due to incomplete octet. - $BF_3 + NH_3 \rightarrow F_3B-NH_3$ - Boron has an empty orbital and can accept electrons. - Metal ions (cations) with empty orbitals can also act as Lewis acids. - Lewis bases must have a lone pair of electrons. #### Conjugate Acids and Bases - Concept is relevant when discussing acidity and basicity trends. - Acid-Base reactions involving conjugate pairs are often in equilibrium. - Based on Bronsted-Lowry definition. - To find the conjugate base, have the acid act as an $H^+$ donor. - To find the conjugate acid, have the base act as an $H^+$ acceptor. - Examples: - $HCl$ (conjugate acid) has $Cl^-$ as its conjugate base. - $NH_3$ (conjugate base) has $NH_4^+$ as its conjugate acid. - Amphiprotic species can act as both proton donors and acceptors. - $HCO_3^-$ (bicarbonate) can act as an acid or a base. - Conjugate base: $CO_3^{2-}$ - Conjugate acid: $H_2CO_3$ - $OH^-$ Hydroxide - Conjugate acid: $H_2O$ Water - Conjugate base: $O^{-2}$ Oxide - In a reversible acid-base reaction, identify acids and bases on both sides of the reaction. - $CH_3COOH + NH_3 \rightleftharpoons CH_3COO^- + NH_4^+$ in the forward direction, $CH_3COOH$ is the acid and $NH_3$ is the base. In the reverse direction, $CH_3COO^-$ is the base and $NH_4^+$ is the acid. #### Strong vs. Weak Acids and Bases - Strength is based on the degree of dissociation. - Strong acids dissociate 100% in water. - Seven common strong acids: - $HCl$ - $HBr$ - $HI$ - $HNO_3$ - $HClO_4$ - $H_2SO_4$ - $HClO_3$ - In aqueous solutions of strong acids, the only acid present is $H_3O^+$ (leveling effect). - Weak acids dissociate partially (less than 5%). Concentration-dependent. - HF is a weak acid, not a strong acid. - Acetic acid ($CH_3COOH$ or $CH_3CO_2H$) is another common weak acid. - Carboxylic acids contain the COOH group. - Strong bases are typically metal hydroxides. - Group 1 metal hydroxides (LiOH, NaOH, KOH, RbOH, CsOH). - Most Group 2 metal hydroxides (Mg(OH)2, Ca(OH)2, Sr(OH)2, Ba(OH)2). - Solubility matters; Group 2 hydroxides are not highly soluble. - The amount that does dissolve, dissociates completely. - Weak bases are more complex to identify. - Ammonia ($NH_3$) is a common weak base. - Organic amines (derivatives of ammonia) are also weak bases. - $CH_3NH_2$ (methylamine) is an example.