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.