Chem 2 Chapter 16 Questions (medium)

Introduction to Acids, Bases, and pH

• Arrhenius acids produce H+ ions in water
• Arrhenius bases produce OH- ions in water
• Bronsted-Lowry acids are proton donors
• Bronsted-Lowry bases are proton acceptors
• A conjugate acid-base pair includes two substances that differ by one proton (H+)
• HF and F- are a conjugate acid-base pair
• HF donates a proton to become F-
• F- can accept a proton to become HF again

Conjugate Acid/Base Examples

• Conjugate base for NH3 (ammonia) is NH2-
• Conjugate acid for NH3 is NH4+ (ammonium ion)
• Conjugate base involves the initial species acting like an acid and donating a proton
• Conjugate acid involves the initial species acting like a base and accepting a proton

Autoionization of Water

• Autoionization of water equation: H2O(l) + H2O(l) = H3O+(aq) + OH-(aq)
• Kw = [H3O+][OH-] = 1.0 × 10-14 at 25°C
• Kw comes from the equilibrium constant expression (products over reactants) for the autoionization of water

Amphoteric Substances

• Amphoteric substances can act as both an acid and a base
• Water (H2O) can donate a proton (acting as an acid) or accept a proton (acting as a base)

Hydrogen Ions

• H+ is a proton, a hydrogen atom that has lost its electron
• In water, H+ combines with a water molecule to form H3O+ (hydronium ion)
• H+ and H3O+is often used interchangeably

Kw and the pH Scale

• Kw is the product of the concentrations of H3O+ (hydronium) and OH- (hydroxide) ions in water
• pH scale is based on the concentration of H3O+ in a solution
• Increased [H3O+] decreases pH
• Decreased [H3O+] increases pH
• A neutral solution has a pH of 7, where [H3O+] = [OH-]
• The pH scale is only accurate at 25 °C
• Acidic solutions have pH less than 7
• Basic solutions have pH greater than 7

pH Calculation

• Strong acids 100% dissociate
• pH of a strong acid can be found by taking the negative logarithm of the concentration of the strong acid ([HCl] = [H+] = [H3O+])
• For [NaOH] = [OH-]
• For Ba(OH)2 the concentration of OH is twice the concentration of Ba(OH)2: Ba(OH)2 → Ba2+ + 2 OH
• pOH = -log[OH]
• pH = 14 – pOH

Strong Acids/Bases

• The 7 strong acids: HCl (hydrochloric acid), HBr (hydrobromic acid), HI (hydroiodic acid), HNO3 (nitric acid), HClO4 (perchloric acid), H2SO4 (sulfuric acid), HClO3 (chloric acid)
• Strong bases completely dissociates in water, producing OH- ions
• Common strong bases include hydroxides of Group 1 metals like NaOH (sodium hydroxide) and KOH (potassium hydroxide), and some Group 2 metals (Ca2+, Sr2+, and Ba2+) like Ba(OH)2 (barium hydroxide)

pH and pOH

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

Weak Acids and Bases

• A weak acid or base only partially dissociates in water
• Makes it less effective at donating or accepting protons compared to a strong acid or base
• Dissociation of HCl would get a ‘straight' arrow, which means 100% dissociation, but weak acids/bases will get

Equilibrium Expression

• For a weak acid: HA + H2O = H3O+ + A-
• The equilibrium expression is: K₁ = [H3O+][A-] / [HA]
• For a weak base: B + H2O = BH+ + OH-
• The equilibrium expression is: K₁ = [BH+][OH-] / [B]
• The pH of a weak acid or base is calculated using the equilibrium expression (Kaor K♭) and an ICE table

ICE Charts

• The pH for strong acids/bases can be calculated directly from their intial concerations
• The chemical equation at the top of the ICE chart is: acid + water = hydronium ion + conjugate base or base + water hydroxide ion + conjugate acid

Acid and Base Strength

• The stronger an acid, the weaker its conjugate base, and the stronger a base, the weaker its conjugate acid
• Conjugates of strong acids and bases are actually considered so weak that they don't have any acid/base properties at all
• For a given conjugate acid-base pair, the product of Ka and K♭ is constant (Kw = Ka × K♭)
• If an acid has a high Ka, its conjugate base will have a low K♭, and vice versa

Acids

• A diprotic acid can donate two protons (H+), like H2SO4
• A polyprotic acid can donate more than one proton, such as H3PO4, which can donate three protons
• HF is monoprotic and only donates 1 proton

pH of 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 has an acidic solution
• A salt from a weak acid and strong base has a basic solution
• A salt from a strong acid and strong base has a neutral solution
• Conjugates of a strong acid or base (like Cl is the conjugate of a weak acid) are considered so weak that they are neutral

Calculating pH

• To find the pH, you need the concentration of hydronium ions
• The amount of hydronium ions the weak acid produces is so small compared to the amount produced by the strong acid that it can be ignored
• To calculate the pH of a solution with a strong acid and a weak acid, you can just use the concentration of the strong acidbecause it completely dissociates and dominates the pH

Salt Solutions

• The hydrolysis equation for NH4Cl: NH4Cl (s) → NH4+ (aq) + Cl- (aq)
• The NH4+ acts as a weak acid, donating a proton to water to form H3O+, which makes the solution acidic: NH4+ (aq) + H2O (l) = NH3 (aq) + H3O+ (aq)
• The Cl- ion does not affect pH because it comes from a strong acid (HCl) and doesn't react with water

Acids and Base Definitions

• A Lewis acid is a substance that accepts a pair of electrons, while a Lewis base donates a pair of electrons
• The Arrhenius definition is the most specific definition, while the Lewis definition is the broadest
• All squares (specific) are rectangles (broad), but not all rectangles (broad) are squares (specific)
• Similarly, all Arrhenius acids (specific) are Lewis acids (broad), but not all Lewis acids (broad) are Arrhenius acids (specific)