Chapter 11: Acids and Bases

Defining and Properties of Acids and Bases

• Arrhenius Definition of Acids: An acid is any substance that produces hydrogen ions ($H^+$) when dissolved in water.
• Properties of Acids:   - They possess a sour taste.   - They are capable of corroding some metals.   - They cause blue litmus paper to turn red.   - They are electrolytes due to the production of $H^+$ ions in water.
• Arrhenius Definition of Bases: A base is any substance that produces hydroxide ions ($OH^-$) when dissolved in water.
• Properties of Bases:   - They possess a bitter or chalky taste.   - They feel soapy and slippery to the touch.   - They cause red litmus paper to turn blue.   - They cause phenolphthalein indicator to turn pink.   - They are electrolytes due to the production of hydroxide ions ($OH^-$) in water.
• Limitations of the Arrhenius Definition: This definition is limited by the requirement that acids and bases only be present in aqueous (water-based) conditions.

Naming Consistency for Acids and Bases

• Naming Binary Acids: Acids composed of a hydrogen ion ($H^+$) and a nonmetal (or $CN^-$) ion are named using the prefix "hydro-" and the suffix "-ic acid."   - Example: $HCl(aq)$ is named hydrochloric acid.
• Naming Oxyacids: Acids comprising a hydrogen ion ($H^+$) and a polyatomic ion are named based on the suffix of the polyatomic ion:   - If the polyatomic ion ends in "-ate," the acid name ends in "-ic acid."     - Example: $ClO_3^-$ (chlorate anion) becomes $HClO_3$ (chloric acid).   - If the polyatomic ion ends in "-ite," the acid name ends in "-ous acid."     - Example: $ClO_2^-$ (chlorite anion) becomes $HClO_2$ (chlorous acid).
• Naming Bases: Typical Arrhenius bases are named as the metal followed by "hydroxide."   - $NaOH$: sodium hydroxide   - $KOH$: potassium hydroxide   - $Ba(OH)_2$: barium hydroxide   - $Al(OH)_3$: aluminum hydroxide   - Industrial Note: Calcium hydroxide, $Ca(OH)_2$, is employed in the food industry for beverage production and in dentistry as a filler for root canals.

Br nsted–Lowry Theory and Conjugate Pairs

• Br nsted-Lowry Definition:   - Acid: A substance that serves as a proton ($H^+$) donor.   - Base: A substance that serves as a proton ($H^+$) acceptor.
• Interaction Example (Ammonia and Water):   - In the reaction $NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$, $NH_3$ acts as the base because it accepts a proton, and $H_2O$ acts as the acid because it donates a proton.
• Conjugate Acid-Base Pairs:   - Every acid-base reaction involves two conjugate acid-base pairs.   - Each pair is related by the specific loss and gain of one $H^+$ ion.   - One pair operates in the forward direction, while the other operates in the reverse direction.   - General Equation: $HA + B \rightleftharpoons A^- + BH^+$     - Pair 1: $HA$ (Acid) and $A^-$ (Conjugate Base).     - Pair 2: $B$ (Base) and $BH^+$ (Conjugate Acid).

Strong vs. Weak Acids and Bases

• Strong Acids:   - Dissociate 100% into ions in water (strong electrolytes).   - Produce large concentrations of hydronium ($H_3O^+$) and anions ($A^-$).   - The chemical reaction is designated with a single forward arrow ($\rightarrow$).   - Example: $HCl(g) + H_2O(l) \rightarrow H_3O^+(aq) + Cl^-(aq)$.
• The Six Common Strong Acids:   1. Hydroiodic acid ($HI$)   2. Hydrobromic acid ($HBr$)   3. Perchloric acid ($HClO_4$)   4. Hydrochloric acid ($HCl$)   5. Sulfuric acid ($H_2SO_4$)   6. Nitric acid ($HNO_3$)
• Weak Acids:   - Only a few molecules dissociate in water (weak electrolytes).   - The majority of the acid remains in its undissociated molecular form.   - Concentrations of $H_3O^+$ and the anion ($A^-$) are small.   - The reaction is designated with a double arrow ($\rightleftharpoons$).   - Example: Carbonic acid ($H_2CO_3(aq) + H_2O(l) \rightleftharpoons H_3O^+(aq) + HCO_3^-(aq)$).
• Common Weak Acids and Conjugate Bases:   - Hydronium ion ($H_3O^+$) → Water ($H_2O$)   - Hydrogen sulfate ion ($HSO_4^-$) → Sulfate ion ($SO_4^{2-}$)   - Phosphoric acid ($H_3PO_4$) → Dihydrogen phosphate ion ($H_2PO_4^-$)   - Hydrofluoric acid ($HF$) → Fluoride ion ($F^-$)   - Nitrous acid ($HNO_2$) → Nitrite ion ($NO_2^-$)   - Acetic acid ($HC_2H_3O_2$) → Acetate ion ($C_2H_3O_2^-$)   - Carbonic acid ($H_2CO_3$) → Bicarbonate ion ($HCO_3^-$)   - Hydrosulfuric acid ($H_2S$) → Hydrogen sulfide ion ($HS^-$)   - Dihydrogen phosphate ion ($H_2PO_4^-$) → Hydrogen phosphate ion ($HPO_4^{2-}$)   - Ammonium ion ($NH_4^+$) → Ammonia ($NH_3$)   - Bicarbonate ion ($HCO_3^-$) → Carbonate ion ($CO_3^{2-}$)   - Hydrogen sulfide ion ($HS^-$) → Sulfide ion ($S^{2-}$)   - Water ($H_2O$) → Hydroxide ion ($OH^-$)
• Strong Bases:   - Specifically include: Lithium hydroxide ($LiOH$), Sodium hydroxide ($NaOH$), Potassium hydroxide ($KOH$), Strontium hydroxide ($Sr(OH)_2$), Calcium hydroxide ($Ca(OH)_2$), and Barium hydroxide ($Ba(OH)_2$).   - Note: Alkaline earth hydroxides like $Ca(OH)_2$ have low solubility but dissociate completely.
• Weak Bases:   - Everything besides the strong bases listed is considered weak.   - Example: $NH_3(g) + H_2O(l) \rightleftharpoons NH_4^+(aq) + OH^-(aq)$.
• Equilibrium Constant ($K_a$ and $Kb$):   - $K_a$ measures weak acid strength; $K_b$ measures weak base strength.   - The larger the value of $K_a$ or $K_b$, the stronger the respective acid or base.

The Autoionization of Water and $K_w$

• Amphoteric Nature: Water is amphoteric, meaning it can act as both an acid and a base.
• Autoionization Reaction: Water reacts with itself to form acid and base ions: $H_2O(l) + H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)$.
• Ion Product Constant for Water ($K_w$): At $25^\circ C$, $K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$.
• Particle Statistics: If you randomly took 10 particles from a liter of water every 10 seconds, it would take 2 years of non-stop work to find a single $H_3O^+$ ion.
• Solution Classification based on Concentration:   - Acidic solution: $[H_3O^+] > [OH^-]$ ($[H_3O^+] > 1.0 \times 10^{-7}\,M$).   - Neutral solution: $[H_3O^+] = [OH^-]$ ($[H_3O^+] = 1.0 \times 10^{-7}\,M$).   - Basic solution: $[H_3O^+] < [OH^-]$ ($[H_3O^+] < 1.0 \times 10^{-7}\,M$).

The pH Scale and Logarithmic Calculations

• pH Significance: The pH scale indicates a solution’s acidity, ranging from 0 to 14, representing the hydronium concentration.   - Acidic: $pH < 7.0$   - Neutral: $pH = 7.0$   - Basic: $pH > 7.0$
• Calculations and Formulas:   - $pH = -\log[H^+]$   - $pOH = -\log[OH^-]$   - $pH + pOH = 14$   - $[H^+] = 10^{-pH}$   - $[OH^-] = 10^{-pOH}$
• Guide to Calculating pH of an Aqueous Solution:   1. Enter the $[H_3O^+]$ concentration.   2. Press the log key and reverse the sign.   3. Significant Figures (SFs): Adjust the number of decimal places on the right side of the pH value decimal point to equal the number of SFs in the original coefficient of the concentration.
• Specific Example: For $[H_3O^+] = 1.7 \times 10^{-3}\,M$.   - Log of $1.7 \times 10^{-3}$ is $-2.7695$.   - Reversing sign gives $2.7695$.   - Since $1.7$ has 2 SFs, the pH is rounded to 2 decimal places: $2.77$.

Buffers: Resistance to pH Change

• Definition: A buffer is a solution capable of resisting large changes in pH when small amounts of acid or base are added.
• Components: A buffer consists of acid-base conjugate pairs in nearly equal concentrations.   - Combination 1: A weak acid and a salt containing its conjugate base.   - Combination 2: A weak base and a salt containing its conjugate acid.
• Buffer Action Mechanism:   - The weak acid in the buffer neutralizes added base.   - The conjugate base in the buffer neutralizes added acid.   - This dual role maintains the solution's pH.
• Buffer Identification Examples:   - $HCl$ and $KCl$: Not a buffer (contains a strong acid).   - $H_2CO_3$ and $NaHCO_3$: A buffer (weak acid + salt of conjugate base).   - $H_3PO_4$ and $NaCl$: Not a buffer (acid and salt are not a conjugate pair).   - $HC_2H_3O_2$ and $KC_2H_3O_2$: A buffer (weak acid + salt of conjugate base).

Skill Checks and Quantitative Examples

• Naming Skill Check:   - $HBr$: Hydrobromic acid (Hydrogen + nonmetal).   - $H_2CO_3$: Carbonic acid ($CO_3^{2-}$ is carbonate).   - $HBrO_2$: Bromous acid ($BrO_2^-$ is bromite).
• Acid/Base Strength Identification:   - $HBr$: Strong acid.   - $HNO_2$: Weak acid.   - $NaOH$: Strong base.   - $H_2SO_4$: Strong acid.   - $Cu(OH)_2$: Strong base (it is a metal hydroxide despite solubility).   - $NH_3$: Weak base.
• Concentration Problem: If lemon juice has $[H_3O^+] = 2 \times 10^{-3}\,M$, solve for $[OH^-]$.   - $[OH^-] = \frac{1.0 \times 10^{-14}}{2.0 \times 10^{-3}}$   - Result: $[OH^-] = 5.0 \times 10^{-12}\,M$.
• pH Practice Problems:   - Coffee: If $[H_3O^+] = 1 \times 10^{-5}\,M$, then $pH = -\log(1 \times 10^{-5}) = 5$.   - pH from $OH^-$: If $[OH^-] = 1.0 \times 10^{-3}\,M$, then $pOH = -\log(1.0 \times 10^{-3}) = 3.0$. Therefore, $pH = 14 - 3.0 = 11.0$.   - $[H_3O^+]$ from pH: If $pH = 11.2$, then $[H_3O^+] = 10^{-11.2} = 6.31 \times 10^{-12}\,M$.