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ACIDS and BASES

ACIDS and BASES. Chapter 18. Acids and Bases: An Introduction. Acidic solution – contains more hydrogen ions than hydroxide ions. [H + ]>[OH - ] Basic solution – contains more hydroxide ions than hydrogen ions. [OH - ]>[H + ]

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ACIDS and BASES

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  1. ACIDS and BASES Chapter 18

  2. Acids and Bases: An Introduction • Acidic solution – contains more hydrogen ions than hydroxide ions. [H+]>[OH-] • Basic solution – contains more hydroxide ions than hydrogen ions. [OH-]>[H+] • Arrhenius model – states that an acid is a substance that contains hydrogen and ionizes to produce hydrogen ions in aqueous solution. • HCl→ H+(aq) + Cl-(aq) • NaOH→ Na+(aq) + OH-(aq) • Bronsted-Lowry model – an acid is a hydrogen-ion donor and a base is a hydrogen-ion acceptor. • HX(aq) + H2O(l) ←→ H3O+(aq) + X-(aq)

  3. Acids and Bases • Conjugate acid – the species produced when a base accepts a hydrogen ion from an acid. • Conjugate base – the species that results when an acid donates a hydrogen ion to a base. • Conjugate acid-base pair – consists of two substances related to each other by the donating and accepting of a single hydrogen ion. • Amphoteric – substances that can act as both acids and bases. (H2O) • Anhydrides – oxides that become acids or bases by adding the elements contained in water. • CO2(g) + H2O(l) → H2CO3(aq) • CaO(s) + H2O(l) → Ca2+(aq) + 2OH-(l)

  4. Strengths of Acids and Bases • Strong Acids – acids that ionize completely. • Produce the maximum number of ions, they are good conductors of electricity. HClO4, HNO3, HI • HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq) • Weak Acids – acids that ionizes only partially in dilute aqueous solutions. • Strong Acids produce a weak conjugate base. • Weak Acids produce a strong conjugate base.

  5. Strengths of Acids and Bases • Acid ionization constant (Ka) – the value of the equilibrium constant expression for the ionization of a weak acid. • Table 19-2 pg.605 • Example: Write ionization equations and acid ionization constant expressions for the following acids. • HClO2 & HNO2 • HClO2(aq) + H2O(l) ←→ H3O+(aq) + ClO-(aq) • Ka=[H3O+][ClO2-]/[HClO2] • HNO2(aq) + H2O(l) ←→ H3O+(aq) + NO2-(aq) • Ka=[H3O+][NO2-]/[HNO2]

  6. Strengths of Acids and Bases • Strong Bases – dissociate entirely into metal ions and hydroxide ions. • Examples: metallic hydroxides, sodium hydroxide • NaOH(s) → Na+(aq) + OH-(aq) • Weak Bases – ionizes only partially in dilute aqueous solution to form the conjugate acid of the base and hydroxide ion. • CH3NH2(aq) + H2O(l) ←→ CH3NH3+(aq) + OH-(aq) • Strong bases produce a weak conjugate acid • Weak bases produce a strong conjugate acid

  7. Strengths of Acids and Bases • Base ionization constant (Kb)- the value of the equilibrium constant expression for the ionization of a base. • Table 19-4 pg.607 • Example: Write the ionization equations and base ionization constant expressions for the following bases. • Hexylamine (C6H13NH2) & Propylamine (C3H7NH2) • C6H13NH2(aq) + H2O(l) ←→ C6H13NH3+(aq) + OH-(aq) • Kb=[C6H13NH3+][OH-]/[C6H13NH3] • C3H7NH2(aq) + H2O(l) ←→ C3H7NH3+(aq) + OH-(aq) • Kb=[C3H7NH3+][OH-]/[C3H7NH2]

  8. What is pH? • Concentration of pure water. • Keq[H2O] = Kw = [H+][OH-] = (1.0x10-7)(1.0x10-7) • Kw=1.0x10-14 • Ion product constant for water – the value of the equilibrium constant expression for the self-ionization of water. • Example: Calculate [H+] or [OH-]. Is the solution acidic, basic, or neutral? • [H+] = 1.0x10-13M • [OH-] = 1.0x10-7M

  9. What is pH? • pH – the negative logarithm of the hydrogen ion concentration. • pH = -log[H+] • pOH – negative logarithm of the hydroxide ion concentration. • pOH = -log[OH-] • pH + pOH = 14.00

  10. What is pH? • Examples: Calculate the pH or pOH of solutions having the following ion concentrations. • [H+] = 1.0x10-2M • [OH-] = 8.2x10-6M • [OH-] = 6.5x10-4M • [H+] = 0.025M

  11. What is pH? • Calculating ion concentrations from pH & pOH • Antilog(-pH) = [H+] • Antilog(-pOH) = [OH-] • Example: Calculate [H+] and [OH-]. • pH = 2.37 • pH = 11.05 • Example: Calculate the pH of the following. • 1.0M HI • 0.050M HNO3 • 1.0M KOH

  12. What is pH? • Example: Calculate the Ka for the following acid using the given information. • 0.0400M solution of HClO2, pH = 1.80

  13. Neutralization • Neutralization reaction – a reaction in which an acid and a base react in aqueous solution to produce a salt and water. • Double-replacement reaction • Salt – an ionic compound made up of a cation from a base and an anion from an acid. • NaOH + HCl → NaCl + H2O • base + acid → salt + water

  14. Neutralization • Example: Write a balanced formula equation for the following acid-base neutralization reaction. • Nitric acid and cesium hydroxide • Hydrobromic acid and calcium hydroxide

  15. Acid-base Titration • Titration – a method for determining the concentration of a solution by reacting a known volume of the solution with a solution of know concentration. • Acid-base indicators – chemical dyes whose colors are affected by acidic and basic solutions. • End point – point at which the indicator used in a titration changes color.

  16. Acid-base reactions • Salt hydrolysis – the anions of the dissociated salt accept hydrogen ions from water or the cations of the dissociated salt donate hydrogen ions to water. • Buffers – solutions that resist changes in pH when limited amounts of acid or base are added. • Buffer capacity – amount of acid or base a buffer solution can absorb without a significant change in pH.

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