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Ch. 20: Acids and Bases. Ch. 20.1 Describing Acids and Bases Ch. 20.2 Hydrogen Ions and Acidity Ch. 20.3 Acid-Base Theories Ch. 20.4 Strengths of Acids and Bases. Ch. 20.1 Describing Acids and Bases. Properties of Acids and Bases Acids Produce H + ions when dissolved in water

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ch 20 acids and bases
Ch. 20: Acids and Bases
  • Ch. 20.1 Describing Acids and Bases
  • Ch. 20.2 Hydrogen Ions and Acidity
  • Ch. 20.3 Acid-Base Theories
  • Ch. 20.4 Strengths of Acids and Bases
ch 20 1 describing acids and bases
Ch. 20.1 Describing Acids and Bases
  • Properties of Acids and Bases
    • Acids
      • Produce H+ ions when dissolved in water
      • Sour taste
      • Solutions are electrolytes (some strong, some weak)
      • React with metals to produce H2
      • React with a base to form water and salt
      • Turn litmus paper red
    • Bases
      • Produce OH- ions when dissolved in water
      • Bitter taste
      • Feel slippery
      • Solutions are electrolytes (strong and weak)
      • React with acids to form water and a salt
      • Turn litmus paper blue
ch 20 1 describing acids and bases3
Ch. 20.1 Describing Acids and Bases
  • Names and formulas of acids and bases
    • Acids
      • Acids have a hydrogen ion
      • The general formula for an acid is HX, where the X is a monatomic or polyatomic ion
    • Bases
      • Bases have an OH- ion
      • Ionic compounds that are bases are named like any other ionic compound
    • See Table 20.1, pg. 578
ch 20 2 hydrogen ions and acidity
Ch. 20.2 Hydrogen Ions and Acidity
  • Hydrogen ions from water
    • Water molecules that gain a hydrogen ion become a hydronium ion (H3O+)
    • Water molecules that lose a hydrogen ion become a hydroxide ion (OH-)
    • In pure water, the concentration of H+ and OH- ions are each 1.0 x 10-7 M
      • This means that the concentration of each are equal in pure water
      • Described as a neutral solution
ch 20 2 hydrogen ions and acidity5
Ch. 20.2 Hydrogen Ions and Acidity
  • Hydrogen ions from water
    • In any aqueous solution, the [H+] and [OH-] are interdependent
      • When [H+] decreases, the [OH-] increases
      • For aqueous solutions, [H+] x [OH-] = 1.0 x 10-14
      • This is also known as the ion-product constant for water (Kw)
      • An acidic solution is one in which the [H+] concentration is greater than the [OH-]
        • Therefore, the [H+] is greater than 1 x 10-7
      • A basic (alkaline) solution is one in which the [OH-] is greater than than the [H+] concentration
        • Therefore, the [H+]is less than 1 x 10-7
ch 20 2 hydrogen ions and acidity6
Ch. 20.2 Hydrogen Ions and Acidity
  • The pH concept
    • Expressing concentration in molarity is inefficient, so we use a pH scale
      • The scale ranges from 1 to 14
        • 1 is very acidic, 7 is neutral, and 14 is very basic
      • The pH of a solution is the negative logarithm of the hydrogen-ion concentration
        • pH = -log [H+]
      • The pOH of a solution equals the negative logarithm of the hydroxide-ion concentration
        • pOH = -log[OH-]
      • pH + pOH = 14
ch 20 2 hydrogen ions and acidity7
Ch. 20.2 Hydrogen Ions and Acidity
  • Calculating pH values
    • Most pH values are not whole numbers
    • You can calculate the hydrogen-ion concentration of a solution if you know the pH
      • If the pH is 3, then [H+] = 1.0 x 10 –3
      • If the pH is not a whole number, you will need a calculator to find antilog
        • [H+] = -pH antilog
ch 20 2 hydrogen ions and acidity8
Ch. 20.2 Hydrogen Ions and Acidity
  • Measuring pH
    • A pH meter is preferred for precise measurements
      • Must be calibrated first by dipping the electrodes in a solution of known pH
      • It is then rinsed and used to measure the pH of an unknown solution
    • Acid-base indicators
      • An indicator is an acid or base that dissociates in a known pH range
      • See Fig. 20.8, pg. 590
      • These have limitations
        • Some have a specific temperature range
        • Do not work well in colored/cloudy solutions
        • Can be affected by dissolved salts
ch 20 3 acid base theories
Ch. 20.3 Acid-Base Theories
  • Arrhenius acids and bases
    • Acids dissociate in water to produce H+ ions
    • Bases dissociate in water to produce OH- ions
    • The Arrhenius definition is very broad
      • Does not include certain chemicals that should be classified as an acid or base
      • NH3 and Na2CO3 are both bases but would not be classified as such under the Arrhenius definition
ch 20 3 acid base theories10
Ch. 20.3 Acid-Base Theories
  • Types of acids
    • Monoprotic acids – acids that contain one ionizable hydrogen
    • Diprotic acids – acids that produce two ionizable hydrogens
    • Triprotic acids – acids that contain three ionizable hydrogens
      • Not all compounds that contain H are acids
      • Not all hydrogens in an acid may be released
  • Acid and base strength is based on solubility
    • Greater dissociation means greater strength
    • Group 1 metals are more soluble than Group 2 metals
ch 20 3 acid base theories11
Ch. 20.3 Acid-Base Theories
  • Bronsted-Lowry acids and bases
    • Defines an acid as a hydrogen-ion donor (proton donor)
    • Defines a base as a hydrogen-ion acceptor (proton acceptor)
    • Conjugate acid-base pairs
      • A conjugate acid is the particle formed when a base gains a hydrogen ion
      • A conjugate base is the particle that remains when an acid has donated a hydrogen ion
      • A conjugate acid-base pair is made up of two substances related by the loss or gain of a single hydrogen ion
        • Water is amphoteric (amphoprotic) – it can accept or donate a hydrogen ion
ch 20 3 acid base theories12
Ch. 20.3 Acid-Base Theories
  • Lewis acids and bases
    • Focuses on the donation or acceptance of a pair of electrons during a reaction
      • More general than the Arrhenius or Bronsted-Lowry definitions
      • A Lewis acid is one that can accept a pair of electrons to form a covalent bond
      • A Lewis base is one that can donate a pair of electrons to form a covalent bond
      • See Table 20.6, pg. 598 for a summary of the three definitions
ch 20 4 strengths of acids and bases
Ch. 20.4 Strengths of Acids and Bases
  • Strong and weak acids and bases
    • Acids
      • Strong acids are completely ionized in aqueous solution
      • Weak acids are only partially ionized in aqueous solutions
        • See Table 20.7, pg. 600 for a list of acids/bases
      • Ka is the acid dissociation constant
        • The ratio of the concentration of the dissociated acid to the concentration of the undissociated acid in a solution
        • Ka = [H+][A-] / [HA]
ch 20 4 strengths of acids and bases14
Ch. 20.4 Strengths of Acids and Bases
  • Ka
    • Reflects the fraction of an acid formed
      • If the Ka is small, then the then the degree of dissociation is small (weak acid)
      • If the Ka is large, then the degree of dissociation is large (strong acid)
    • Diprotic and triprotic acids lose their H+ ions one at a time
      • Each reaction has its own Ka
ch 20 4 strengths of acids and bases15
Ch. 20.4 Strengths of Acids and Bases
  • Strong and weak acids and bases
    • Bases
      • Strong bases dissociate completely into metal ions and hydroxide ions in aqueous solutions
      • Weak bases react with water to form the hydroxide ion and the conjugate acid of the base
        • The base dissociation constant (Kb) is the ratio of the concentration of the hydroxide ion to the concentration of the conjugate base
        • Kb = [HB+][OH-] / [B]
        • The smaller the value of Kb, the weaker the base
ch 20 4 strengths of acids and bases16
Ch. 20.4 Strengths of Acids and Bases
  • Strong and weak acids and bases
    • Concentrated and dilute refer to how much of an acid or base is dissolved in solution
      • Moles of acid/base per liter
    • Strong and weak refer to the extent of ionization or dissociation of an acid or base
      • a solution of ammonia, whether concentrated or dilute, will be a weak base because the ionization NH3 will be small
ch 20 4 strengths of acids and bases17
Ch. 20.4 Strengths of Acids and Bases
  • Calculating dissociation constants (Ka and Kb)
    • It is possible to calculate Ka and Kb from experimental data
    • First, measure the equilibrium concentrations of all the substances present at equilibrium
    • Then put into the Ka or Kb formula
      • See Sample Problem 20-8, pg. 604