Acids & Bases

UNIT 13: ACIDS AND BASES (chapter 18)

I. PROPERTIES OF ACIDS AND BASES

A. Introduction
  • Common Examples:
    • Inorganic acids
    • Organic acids: contain the carboxyl group (-COOH)
      • Formic acid - HCOOH
      • Acetic acid - CH3COOH
      • Citric acid
      • Ascorbic acid (vitamin c)
      • Lactic acid (found in milk)
    • Inorganic bases:
      • Sodium hydroxide (NaOH)
      • Magnesium hydroxide (Mg(OH)2)
      • NH4OH
    • Organic bases: aka "amines" - has R-NH2
      • CH3NH2
      • C2H3NH2
  • Acid solutions: sour taste
  • Basic solutions: bitter taste and feel slippery
  • Can both be corrosive and/or caustic (causes burns)
  • Solutions conduct electricity (known as electrolytes)
  • Have an effect on indicators (things that turn diff colors) - dyes whose colors vary with type of solution.
    • Ex:
      • litmus: pink/red in acid, blue in base
      • phenolphthalein: colorless in acid, pink in base
  • Rlly like to react
  • Reactions:
    • Acid + Base à
    • Acids react with some metals
    • Acids react with carbonates
  • All aqueous solutions contain H+ and OH- ions.
    • acidic solution: H+ > OH-
    • basic solution: H+ < OH-
    • neutral solution: H+ = OH-
B. Theories
  • Arrhenius model:
    • acid - contains hydrogen and the hydrogen ionizes (H^+).
    • base
  • Bronsted & Lowry model:
    • acid: Donates H+ (proton donor)
    • base: Accepts a hydrogen ion (proton)
    • Ex:
      • HF(g)+H2O(l)⇌H3O+(aq)+F−(aq)HF (g) + H2O (l) \rightleftharpoons H3O+ (aq) + F- (aq)
      • NH3(g)+H2O(l)⇌NH4+(aq)+OH−(aq)NH3 (g) + H2O (l) \rightleftharpoons NH4 + (aq) + OH - (aq)
    • conjugate acid
      • Acts like a bronsted-lowry acid in the in the reverse reaction.
      • Base w/ extra hydrogen
    • conjugate base
      • Acts like a bronsted-lowry acid in the in the reverse reaction.
      • What's left over after the acid loses one hydrogen ion.
    • Amphoteric
      • A substance that behaves like B-L acid and base (ex: water)
  • Lewis model:
    • acid: electron pair acceptors; Ex: BF3
    • base: electron pair donors; Ex: NH3
C. Protolysis:
  • Ability to transfer a proton (P+) (H+)
  • monoprotic acids:
    • Can only donate one proton p+
    • HCl, HBr, HF, HNO3
    • Note: Organic acids: Contain CH3COOH
  • polyprotic acids:
    • Has more than one ionizable H
      • H2CO3, H2SO4, H3PO4
    • Note: Polyprotic acids ionize in steps.
      • H3PO4+H2O(L)⇌H3O+(aq)+H2PO4−(aq)H3PO4 + H2O(L) \rightleftharpoons H3O+ (aq) + H2PO4- (aq)
      • H2PO4−+H2O(L)⇌H3O+(aq)+HPO42−(aq)H2PO4- + H2O (L) \rightleftharpoons H3O+ (aq) + HPO42 - (aq)
      • HPO42−+H2O(L)⇌H3O+(aq)+PO43−(aq)HPO42 - + H2O (L) \rightleftharpoons H3O + (aq) + PO43 - (aq)

II. STRENGTHS OF ACIDS AND BASES

A. Strong Acids and Bases:
  • Strong acids and bases completely dissociate (separate) into ions.
    • All solute particles
    • binary acids
      • Most binary acids are strong (HCL, HBr, HI)
      • Exception: HF is weak
    • ternary acids (oxyacids) -called oxyacids bc contain oxygen
      • Strong if they have two or more oxygens than hydrogens
    • bases (metallic hydroxides)
      • Metallic hydroxides from group one or two are strong
      • Except Be(OH)2 is weak
B. Weak Acids and Bases:
  • Partially ionize in water.
    • Produce fewer ions
    • Less acidic/basic per mol
    • An equilibrium can be established (D)
    • All products and reactants are present in solution
    • Rate of forward and reverse reactions are equal
    • Organic acids and bases
      • They are weak
  • “weak” vs. “strong”
    • How much the acid/base ionizes
    • “strong" acid or base produces more hydrogen/hydroxide ions per molecule
    • So, we need less volume/mass to do acidic things
  • “dilute” vs. “concentrated”
    • How much solute per solvent that exists
    • Refers to molarity
    • Will affect number of hydrogen/hydroxide ions present
    • Refers to molarity

III. DETERMINING ACIDITY OR BASICITY

A. The Ionization of Water
  • Self-ionization - water rips itself into ions
  • Equal amt of both (water is equal)
  • H2O(l)⇌H+(aq)+OH−(aq)H2O (l) \rightleftharpoons H+ (aq) + OH- (aq)
  • H+(aq)+H2O(l)⇌H3O+(aq)H+ (aq) + H2O (l) \rightleftharpoons H3O+ (aq)(forms hydronium ion, H3O+)
  • the [H+] = [OH-] = 1.0x10−71.0 x 10^{-7}M
  • the concentration of hydrogen ions = concentration of hydroxide ions
  • ion-product constant for water, Kw: Kw=[H+][OH−]=1.0x10−14Kw = [H+][OH-] = 1.0 x 10^{-14}
  • Solution types:
    • acidic: [H+] > [OH-]
    • basic: [H+] < [OH-]
    • neutral: [H+] = [OH-]
  • See problems in text.
B. The pH and pOH Scales
  • pH=−log[H+]pH = -log[H+]
  • pOH=−log[OH−]pOH = -log[OH-]
  • pH+pOH=14.00pH + pOH = 14.00
  • Since pH is a logarithmic scale, a change of 1 pH unit represents a 10-fold change in ion concentration.
    • Ex: A pH of 3 has 10 times more H+ ions than a pH of 4.
  • pH of common substances: see figure 14, p.652
  • pH and sig. figs.: The #sig figs in the [H+] or [OH-] determine the #decimal places in the pH or pOH.
    • Ex:
      • if [H+] = 1.00x10−71.00 x 10^{-7}M, then the pH = 7.000
      • if [H+] = 1.0x10−71.0 x 10^{-7}M, then the pH = 7.00
      • if [H+] = 1x10−71 x 10^{-7}M, then the pH = 7.0
  • See problems in text.
  • Given the pH or pOH, find the ion concentration:
    • Since pH = -log[H+], then the antilog (-pH) = [H+]
    • On calculator:
      • enter 10, yx, with x = -pH (taking inverse log (2nd function) of the –pH)
    • The values for the ion concentration are expressed with as many significant figures as the number of decimal places in the pH.

IV. NEUTRALIZATION

A. The Reaction Between Acids and Bases
  • neutralization reaction: Acid (aq) + Base (aq) à Salt (aq) + Water (l)
  • For any strong acid & strong base neutralization reaction, the net ionic equation is: H+ (aq) + OH- (aq) à H2O (l)
B. Salt Hydrolysis
  • Definition
  • Types of solutions:
    • The salt of a SA & SB -neutral
    • The salt of a SA & WB - acidic
    • The salt of a WA & SB - basic
    • The salt of a WA & WB - determine which is stronger