Acids and Bases

Chapter 9: Acids & Bases

Introduction to Acids and Bases

Acids are hydrogen ion (H+) donors.
Bases are hydrogen ion (H+) acceptors.

Acid-Base reaction

Acids react with bases:
$HA + B \rightleftharpoons A^- + BH^+$
Where:

$HA$ is the acid.
$B$ is the base.
$A^-$ is the conjugate base.
$BH^+$ is the conjugate acid.

Common Acids

Hydrochloric acid (HCl): metal cleaning, food preparation, stomach acid.
Sulfuric acid ( $H2SO4$ ): fertilizer, explosives, dye production, batteries, electroplating.
Nitric acid ( $HNO_3$ ): fertilizer, explosives, dye production.
Acetic acid ( $HC2H3O_2$ ): plastic and rubber manufacturing, food preservative, vinegar.
Citric acid ( $H3C6H5O7$ ): present in citrus fruits, pH adjustment in foods and beverages.
Carbonic acid ( $H2CO3$ ): carbonated beverages (reaction of carbon dioxide with water).
Hydrofluoric acid (HF): metal cleaning, glass frosting and etching.
Phosphoric acid ( $H3PO4$ ): fertilizer manufacture, biological buffering, preservative in beverages.

Acidic Hydrogen

Hydrogens connected to more electronegative atoms (F, Cl, Br, I, O) are more likely to act as acidic hydrogen.

Strong vs. Weak Acids

Strong acid: An acid that completely ionizes/dissociates in water.
$HNO3(aq) + H2O(l) \longrightarrow H3O^+(aq) + NO3^-(aq)$
Weak acid: An acid that only partially ionizes in water.

Common Bases

Sodium hydroxide (NaOH): petroleum processing, soap and plastic manufacturing.
Potassium hydroxide (KOH): cotton processing, electroplating, soap production, batteries.
Sodium bicarbonate ( $NaHCO3$ ): antacid, ingredient of baking soda, source of $CO2$
Sodium carbonate ( $Na2CO3$ ): manufacture of glass and soap, general cleanser, water softener.
Ammonia ( $NH_3$ ): detergent, fertilizer and explosives manufacturing, synthetic fiber production.

Strong vs. Weak Bases

Strong bases: Bases that completely ionize/dissociate in water.
- Examples: Lithium hydroxide (LiOH), Sodium hydroxide (NaOH), Potassium hydroxide (KOH), Strontium hydroxide [Sr(OH)₂], Calcium hydroxide [Ca(OH)₂], Barium hydroxide [Ba(OH)₂].
Weak bases: Bases that only partially ionize in water.
- Examples: Ammonia ( $NH3$ ), Pyridine ( $C5H5N$ ). $NH3 + H2O \rightleftharpoons NH4^+ + OH^-$

Polyprotic Acids

Acids that can release more than one $H^+$ .
Monoprotic acid: can donate only one proton (e.g., $HCl$ ).
$HCl(aq) + H2O(l) \longrightarrow Cl(aq) + H3O^+(aq)$
Polyprotic acid: can donate more than one proton (e.g., $H2SO4$ ).
- Complete ionization of Sulfuric acid ( $H2SO4$ ):
 $H2SO4(aq) + H2O(l) \longrightarrow HSO4^-(aq) + H3O^+(aq)$ $HSO4^-(aq) + H2O(l) \longrightarrow SO4^{2-}(aq) + H_3O^+(aq)$

Common Polyprotic Acids and Ionization Constants

Name (Formula)	Structure	$K_{a1}$	$K_{a2}$	$K_{a3}$
Sulfuric Acid ( $H<em>2SO</em>4$ )	O=S(OH)₂	Strong	1.2 × 10⁻²
Oxalic Acid ( $H<em>2C</em>2O_4$ )	HOOC-COOH	6.0 × 10⁻²	6.1 × 10⁻⁵
Sulfurous Acid ( $H<em>2SO</em>3$ )	HOS(O)OH	1.6 × 10⁻²	6.4 × 10⁻⁸
Phosphoric Acid ( $H<em>3PO</em>4$ )	OP(OH)₃	7.5 × 10⁻³	6.2 × 10⁻⁸	4.2 × 10⁻¹³
Citric Acid ( $H<em>3C</em>6H<em>5O</em>7$ )	HOOC-CH₂-C(OH)(COOH)-CH₂-COOH	7.4 × 10⁻⁴	1.7 × 10⁻⁵	4.0 × 10⁻⁷
Ascorbic Acid ( $H<em>2C</em>6H<em>6O</em>6$ )	Structure not provided, but it's a 6-carbon ring with hydroxyl groups and a lactone	8.0 × 10⁻⁵	1.6 × 10⁻¹²
Carbonic Acid ( $H<em>2CO</em>3$ )	OC(OH)₂	4.3 × 10⁻⁷	5.6 × 10⁻¹¹

Conjugate Acid-Base Pairs

Each acid has a conjugate base and every base has a conjugate acid.
$HA + B \rightleftharpoons A^- + BH^+$
- $HA$ is the acid, and $A^-$ is its conjugate base.
- $B$ is the base, and $BH^+$ is its conjugate acid.

Acid Ionization Constant

Acid Ionization Constant ( $K_a$ ): the equilibrium constant for the ionization reaction of an acid with water.
$HA + H2O \rightleftharpoons A^- + H3O^+$
$K_a = \frac{[H^+][A^-]}{[HA]}$
Large $K_a$ = Strong acid.
Small $K_a$ = Weak acid.

Base Ionization Constant

Base Ionization Constant ( $K_b$ ): the equilibrium constant for the ionization reaction of a base with water.
$B + H_2O \rightleftharpoons BH^+ + OH^-$
$K_b = \frac{[BH^+][OH^-]}{[B]}$
Large $K_b$ = Strong base.
Small $K_b$ = Weak base.

Autoionization of Water

Also called "Self Ionization".
About 1 out of every 10 million water molecules form ions through self ionization.
$2 H2O(l) \rightleftharpoons H3O^+(aq) + OH^-(aq)$
All aqueous solutions contain both $H_3O^+$ and $OH^-$ .

Ion Product Constant for Water

Ion Product Constant for Water ( $K_w$ ): the numerical value obtained by multiplying the molar concentrations for hydronium and hydroxide ions present in pure water at 25°C.
$Kw = [H3O^+][OH^-] = 1.00 \times 10^{-14}$ at 25 °C
The concentration of $H_3O^+$ and $OH^-$ are equal in pure water.
$[H_3O^+] = [OH^-] = 10^{-7} M$ @ 25°C

Acidic and Basic Solutions

Neutral solutions have equal $[H3O^+]$ and $[OH^-]$ . $[H3O^+] = [OH^-] = 1 \times 10^{-7}$
Acidic solutions have a larger $[H3O^+]$ than $[OH^-]$ . [H3O^+] > [OH^-]
[H_3O^+] > 1 \times 10^{-7}; [OH^-] < 1 \times 10^{-7}
Basic solutions have a larger $[OH^-]$ than $[H3O^+]$ . [H3O^+] < [OH^-] [H_3O^+] < 1 \times 10^{-7}; [OH^-] > 1 \times 10^{-7}

pH and pOH

pH is a measure of the concentration of $H^+$ or $H3O^+$ in solution. $pH = -log[H^+]$ OR $pH = -log[H3O^+]$
$[H_3O^+] = 10^{-pH}$
pOH is a measure of the concentration of $OH^-$ in solution.
$pOH = -log[OH^-]$
$[OH^-] = 10^{-pOH}$
$pH + pOH = 14$

Classification of Water Soluble Substances

Electrolytes: Solutes that separate into ions when dissolved in water (they're soluble) and have the ability to conduct electricity.
- Strong electrolytes: Solutes that completely dissociate into ions when dissolved in water (e.g., $NaCl, MgBr_2, HCl$ ). Strong electrical conductors.
  - Strong electrolyte (aq or s) → Cation+(aq) + Anion-(aq)
  - Example: $NaCl(s) \longrightarrow Na^+(aq) + Cl^-(aq)$
- Weak electrolytes: Solutes that, when dissolved in water, only partially dissociate into ions (e.g., $HF, NH_3$ , acetic acid). Weak electrical conductors.
  - Weak electrolyte (aq) → Cation+(aq) + Anion-(aq)
  - Example: $HF(aq) \rightleftharpoons H^+(aq) + F^-(aq)$
Nonelectrolytes: Solutes that dissolve in water without separating into ions (e.g., Sucrose, ethanol). Do not conduct electricity.
- Nonelectrolyte (s or l) → Nonelectrolyte (aq)
- Example: $C{12}H{22}O{11}(s) \longrightarrow C{12}H{22}O{11}(aq)$

Neutralization Reactions

When strong acids and bases in aqueous solution react with each other, they form water and a salt. $HX(aq) + MOH(aq) \longrightarrow H2O(l) + MX(aq)$ $HCl(aq) + NaOH(aq) \longrightarrow H2O(l) + NaCl(aq)$
- HX is the acid.
- MOH is the base.
- MX is the salt.

Titration

A procedure for the quantitative analysis of a substance of unknown concentration whereby a measured quantity of another substance, of known concentration, is completely reacted with the original substance.
Often used to determine the concentration of acids and bases.
Equivalence point: The point in a titration at which one reactant has been exactly consumed by the addition of another reactant.
- Midpoint of vertical rise.
- Occurs at pH = 7 in a strong acid-strong base titration.
- $[H_3O^+] = [OH^-]$

Indicators

Acid-Base Indicator: a chemical that changes color with a change in pH.
Added to solutions in small amounts in order to determine the solution's pH visually.
Usually organic compounds.
Weak acid or base that establishes an equilibrium with the $H2O$ and $H3O^+$ in the solution.

Buffers

If an acid is added to a buffer equilibrium, the excess acid reacts with the conjugate base, so the overall pH does not change much. If a base is added to a buffer equilibrium, the excess base reacts with the conjugate acid, so the overall pH does not change much.