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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. Properties of Acids and Bases Acids Produce H + ions when dissolved in water

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Ch. 20: Acids and Bases

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  1. 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

  2. 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

  3. 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

  4. 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

  5. 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

  6. 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

  7. 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

  8. 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

  9. 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

  10. 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

  11. 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

  12. 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

  13. 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]

  14. 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

  15. 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

  16. 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

  17. 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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