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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. 20.1 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


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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


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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


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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


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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


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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


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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


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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


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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


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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


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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


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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


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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]


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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


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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


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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


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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


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