Acids & Bases

There are three primary touchstones with which chemists define acids and bases:

Arrhenius Definition

An Arrhenius acid dissociates in aqueous solution to produce hydrogen (H+) ions.

Ex: HCl → Cl- + H+

An Arrhenius base dissociates in aqueous solution to produce hydroxide (OH-) ions.

Ex: NaOH → OH- + Na+

Brønsted-Lowry Definition

A Brønsted acid is a proton donor.

Ex: HI + F- → HF + I-

A Brønsted base is a proton acceptor.

Ex: HI + F- → HF + I-

Lewis Definition

A Lewis acid is an electron pair acceptor.

Ex: BF3 + F- → BF4-

A Lewis base is an electron pair donor.

Ex: 2NH3 + Ag+ → Ag(NH3) 2+

Conjugate Acids and Bases

CH3COOH + NH3 → NH4+ + CH3COO-

In the above reaction, ammonium is the conjugate acid (formed when a base gains a proton) of ammonia. Likewise, the product acetate ion is the conjugate base (formed when an acid loses a proton) of acetic acid. Clearly, the concept of conjugate bases and acids stems from the Brønsted-Lowry definition (donation and acceptance of protons). The stronger an acid, the weaker is its conjugate base. The stronger a base, the weaker is its conjugate acid.

Acid-Base Nomenclature

  • The “ic” ending is applied to the acid with the higher oxidation state; the “ous” ending is applied to the acid with the lower oxidation state.
  • The “hypo” prefix is used when the oxidation state is lowest; the “per” prefix is used when the oxidation state is highest.
  • Here is an example:

HClO  (hypochlorous acid)       Chlorine has an oxidation state of +1.

HClO2 (chlorous acid)              Chlorine has an oxidation state of +3.

HClO3 (chloric acid)               Chlorine has an oxidation state of +5.

HClO4 (perchloric acid)           Chlorine has an oxidation state of +7.

Weak and Strong Acids and Bases

Weak acids such as cyanic acid (HOCN), hydrofluoric acid (HF), and phosphoric acid (H3PO4) only partially dissociate to form H+ ions, while strong acids (You should memorize these: HClO4, HI, HBr, HCl, H2SO4, HNO3) completely dissociate.

Likewise, weak bases such as ammonia (NH3), pyridine (C5H5N), and fluoride (F-) only partially dissociate (or react with water) to form OH- ions, while strong bases such as NaOH, KOH, and LiOH completely dissociate.

The Concept of pH

pH is a measure of the acidity of a solution. It is equal to the negative base-10 logarithm of the molarity of dissolved hydrogen ions (H3O+ or H+).  pOH is a measure of the basicity of a solution. It is equal to the negative base-10 logarithm of the molarity of dissolved hydroxide ions (OH-). Here are some important equations to know for manipulating between these important quantities:

pH = -log[H+]        pOH = -log[OH-]        10-pH = [H+]        10-pOH = [OH-]

Self-ionization of Water

Water self-ionizes, forming hydronium and hydroxide:

2H2O(l) ↔ H3O+(aq) + OH(aq)                          At 298.15K & 1atm, Kw = 1.0×10−14 = [H3O+][OH]

∆H > 0

Because this process is endothermic, the self-ionization constant of water, Kw, increases with temperature.

Dissociation of Weak Acids

If “HA” is a general weak monoprotic acid, we can write the equilibrium equation for its ionization as follows:

HA ↔ H+ +  A-

The form of the equilibrium-constant expression is:

Ka is known as the acid dissociation constant. Strong acids have very high (essentially infinitely high) acid dissociation constants.

Dissociation of Weak Bases

If B is a general weak base, then the equilibrium for the weak base reaction is:

B(aq)  +  H2O(l) ↔ BH+(aq)  +  OH-(aq)

The equilibrium constant, symbolized by Kb, for such a reaction is called the base ionization constant, and the equilibrium law expression is:

Acidic and Basic Salts

Salts formed

  • from strong acids and strong bases are neutral in solution.
  • from strong acids and weak bases are acidic in solution.
  • from weak acids and strong bases are basic in solution.
  • from weak acids and weak bases may be acidic, basic, or neutral. It depends on the ions’ relative strengths as bases/acids.

Buffer Solutions

Buffers resist pH changes. Buffer solutions comprise of either a solution of a weak acid in the presence of one of its salts or a solution of a weak base in the presence of one of its salts. Buffer capacity is a measure of the resistance of a buffer solution to pH change. The Henderson-Hasselbalch equation  describes the pH in terms of pKa (-log(Ka)) of a buffer solution:

Acid-Base Titration

An acid-base titration is the determination of the concentration of an acid or base by exactly neutralizing the unknown concentration of acid/base with an acid or base of known concentration. This allows for quantitative analysis of the concentration of an unknown acid or base solution. It makes use of the neutralization reaction that occurs between acids and bases and the knowledge of how acids and bases will react if their formulas are known. A typical acid-base titration requires a:

  • Burette
  • Pipette
  • Indicator – changes color based on pH
  • Erlenmeyer flask
  • Analyte – “unknown” solution to be titrated
  • Titrant – solution of known concentration

A titration curve results from an acid-base titration that looks somewhat like this:

A generic titration curve of a weak acid titrated with a strong base

 

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