Chem 2 Chapter 16 Questions (Hard)

Introduction to Acids, Bases, and pH

• Arrhenius acids yield H⁺ ions in water; Arrhenius bases yield OH⁻ ions in water
• Brønsted-Lowry acids are proton donors; Brønsted-Lowry bases are proton acceptors
• A conjugate acid-base pair includes two substances differing by one proton (H⁺)
• HF and F⁻ are a conjugate acid-base pair because HF donates a proton to become F⁻ and F⁻ accepts a proton to become HF

Conjugate Acids and Bases

• The conjugate base for NH₃ (ammonia) is NH₂⁻
• The conjugate acid for NH₃ is NH₄⁺ (ammonium ion)
• When writing the conjugate base, the initial species acts like an acid by donating a proton
• To write the conjugate acid, the starting species acts like a base that accepts a proton

Autoionization of Water

• The autoionization of water is represented by: H₂O(l) + H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq)
• The expression for Kw is: Kw = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25°C
• Kw is derived from the equilibrium constant expression (products over reactants) for water's autoionization

Amphoteric Substances

• An amphoteric substance acts as both an acid and a base
• Water (H₂O) is a good example, as it can donate (acting as an acid) or accept a proton (acting as a base)

Relationship Between H+, Hydronium Ion, and a Proton

• H⁺ is simply a proton, a hydrogen atom lacking its electron
• H⁺ doesn't exist alone in water, it combines with a water molecule to form H₃O⁺ (hydronium ion)
• H⁺ and H₃O⁺ are often used interchangeably to denote hydrogen ions in solution, especially in pH calculations
• Although technically more correct, some general chemistry resources still use H⁺

Kw and the pH Scale

• Kw is the product of H₃O⁺ (hydronium) and OH⁻ (hydroxide) concentrations in water
• The pH scale relies on the concentration of H₃O⁺ in a solution
• When [H₃O⁺] increases, pH drops; when [H₃O⁺] decreases, pH rises
• A neutral solution has a pH of 7, where [H₃O⁺] = [OH⁻]
• The pH scale is most accurate at 25°C
• An acidic solution has a pH under 7, and a basic solution has a pH above 7

pH Calculation for Strong Acids and Bases

• Strong acids fully dissociate, all strong acid produces H⁺ ions (HCl → H⁺ + Cl⁻)
• You can determine the pH of a strong acid via the negative logarithm of the strong acid's concentration: [HCl] = [H⁺] = [H₃O⁺]
• For strong bases, [NaOH] = [OH⁻] but for Ba(OH)₂, the [OH⁻] is twice the [Ba(OH)₂]: Ba(OH)₂ → Ba²⁺ + 2OH⁻
• Can use the pH equations, or pOH = log[OH⁻], then, pH = 14 – pOH

Strong Acids and Bases

• The 7 strong acids include: HCl (hydrochloric), HBr (hydrobromic), HI (hydroiodic), HNO₃ (nitric), HClO₄ (perchloric), H₂SO₄ (sulfuric), and HClO₃ (chloric)
• Strong bases completely dissociate in water to produce OH⁻ ions
• Common examples include Group 1 metal hydroxides like NaOH (sodium hydroxide) and KOH (potassium hydroxide), and Group 2 metal hydroxides like Ba(OH)₂ (barium hydroxide)

pH, pOH, and the 'p' Function

• The 'p' function in chemistry signifies "the negative logarithm of"
• pH = -log[H⁺]
• pOH = log[OH-]

Weak Acids and Bases

• Weak acids/bases only partially dissociate in water; not all molecules break apart into ions
• This makes them less effective at donating/accepting protons compared to strong acids/bases

Equilibrium Expressions

• For a weak acid: HA + H₂O ⇌ H₃O⁺ + A⁻, equilibrium expression: Ka = [H₃O⁺][A⁻] / [HA]
• For a weak base: B + H₂O ⇌ BH⁺ + OH⁻, the equilibrium expression: Kb = [BH⁺][OH⁻] / [B]

Calculating pH for Weak Acids or Bases

• Calculated using the equilibrium expression (Ka or Kb) and an ICE table
• Is different than strong acids/bases because they fully dissociate, making it easier to calculate pH directly from their initial concentrations—no ICE table needed for strong species.
• Requires finding the equilibrium concentrations of hydronium or hydroxide ions, depending on the acid or base

ICE Charts

• In ICE charts for weak acids/bases, the chemical equation at the top is:
  • acid + water ⇌ hydronium ion + conjugate base, or
• base + water ⇌ hydroxide ion + conjugate acid

ICE Chart Problems

• In ICE chart problems, use the equilibrium constant to find missing info or use the pH and concentration to completely fill out the ICE table to calculate the equilibrium constant
• Similar flavors to problems in chapter 15, but need the additional step of using pH equations to go between concentrations and pH/pOH
• Can almost always use approximation in chapter 16

Acid and Base Strength

• The stronger acid will have a weaker conjugate base, and the stronger base will have a weaker conjugate acid.
• Conjugates of strong acids and bases are considered so weak that they don’t have any acid/base properties
• The relationship Ka × Kb = Kw, where if an acid has a high Ka, its conjugate base will have a low Kb, and vice versa.

Polyprotic Acids

• Diprotic acids like H₂SO₄ donate two protons (H⁺)
• Polyprotic acids like H₃PO₄ can donate more than one proton
• HF is monoprotic and only donates 1 proton

Salt Solutions

• The pH of a salt solution depends on the strength of the acid and base from which the salt is made
• A salt from a strong acid and weak base will form an acidic solution
• A salt from a weak acid and strong base will form a basic solution
• A salt from a strong acid and base will make a neutral solution

pH Calculation

• Requires calculating by breaking the salt into component ions and determining whether each ion is a strong/weak acid/base
• Conjugates of strong acids/bases are considered so weak they are neutral

Mixed Acid/Base Solutions

• To find the pH of a solution with a strong and weak acid, you need the concentration of hydronium ions
• Determine the pH by ignoring the hydronium ions the weak acid/base produces because they are so small compared to the amount produced by the strong acid that it can be ignored

Salt Hydrolysis

• When ammonium chloride (NH₄Cl) dissolves in water, it hydrolyzes: NH₄Cl (s) → NH₄⁺ (aq) + Cl⁻ (aq)
• NH₄⁺ acts as a weak acid, donating a proton to water: NH₄⁺ (aq) + H₂O (l) ⇌ NH₃ (aq) + H₃O⁺ (aq), making the solution acidic
• Cl⁻ ions does not affect pH

Lewis Acids and Bases

• Lewis acids accept a pair of electrons; Lewis bases donate a pair of electrons
• Arrhenius is the most specific definition of acids and bases
• Lewis is the broadest
• All Arrhenius acids are Lewis acids
• Not all Lewis acids are Arrhenius acids

Learn about acids, bases, and pH, including Arrhenius and Brønsted-Lowry definitions. Understand conjugate acid-base pairs and the autoionization of water. Explore calculating Kw and its significance.