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Acids and Bases

Acids and Bases. Part 2. Classifying Acids and Bases. Arrhenius Acid Increases hydrogen ions (H + ) in water Creates H 3 O + (hydronium) Base Increases OH - in water. Classifying Acids and Bases. Brønsted - Lowry Acid Proton donor (H + ) Must have hydrogen in formula Base

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Acids and Bases

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  1. Acids and Bases Part 2

  2. Classifying Acids and Bases • Arrhenius • Acid • Increases hydrogen ions (H+) in water • Creates H3O+ (hydronium) • Base • Increases OH- in water

  3. Classifying Acids and Bases • Brønsted - Lowry • Acid • Proton donor (H+) • Must have hydrogen in formula • Base • Proton acceptor (H+) • Water can be an acid or a base

  4. Classifying Acids and Bases • Lewis • Acid • Electron pair acceptor • Usually positive ions • Base • Electron pair donor • Usually negative ions

  5. Pairs • Acid + Base Conjugate acid + Conjugate base • This is an equilibrium. • The acid becomes the conjugate base after it has donated the H+ • The base becomes the conjugate acid once it accepts the H+

  6. Conjugate Acids and Bases • Identify the acid, base, conjugate acid, and conjugate base in the below reaction: HF + H2O → F- + H3O+ Acid Base conjugate conjugate base acid

  7. Water • Water conducts electricity • Electrolyte • Thus, water self-ionizes • H2O → H+ + OH- • Water is also Amphoteric • Can act as either an acid or a base • Therefore: • 2H2O → H3O+ + OH-

  8. Types of Acids • Monoprotic • Only 1 acidic hydrogen • PolyproticAcids • More than 1 acidic hydrogen • Diprotic – H2SO4 • Triprotic– H3PO4 • Oxyacids • Proton is attached to the oxygen of an ion. • Organic acids • Contain the Carboxyl group –COOH • H attached to O – ex: CH3COOH – acetic acid • Generally very weak

  9. Strong Acids • They completely dissociate in water • HCl • HNO3 • HI • H2SO4 • HClO • HBr

  10. pH of Strong Acid • Calculate the pH and [OH]: • For a 1 x 10-3 M solution of HClO4

  11. Bases • The OH-is a strong base. • Hydroxides of the alkali metals are strong bases because they dissociate completely when dissolved. • Others are weak.

  12. Bases without OH- • Bases are proton acceptors. • NH3 + H2O NH4+ + OH- • It is the lone pair on nitrogen that accepts the proton.

  13. Salts as acids and bases • Salts are ionic compounds. • Salts of the cation of strong bases and the anion of strong acids are neutral. • Salts of a strong acid and weak base are acidic salts • Salts of strong base and weak acids are basic salts

  14. Lewis Acids and Bases • Most general definition. • Acids are electron pair acceptors. • Bases are electron pair donors. F H B F :N H F H

  15. Ion concentrations • The concentrations of H3O+ and OH- are based on water • Kw – ion-product constant of water • At 25°C, Kw= 1.0x10-14 • Kw = (Ka)(Kb) [H3O+][OH-] • [H3O+] – concentration of hydronium ion • [OH-] – concentration of hydroxide ion • Neutral: [H3O+] = [OH-]= 1.0 x10-7 • Acidic:[H3O+] > [OH-] • Basic: [H3O+] < [OH-]

  16. Acid dissociation constant • Acid dissociation constant Ka • Base dissociation constant Kb • The larger the number the stronger the substance • The smaller the number the weaker the substance

  17. Calculating pH • pH is a measure of hydronium [H3O+] concentration in solution • Lower pH = more acidic • pH = -log [H3O+]

  18. pH Scale

  19. pH Calculations • What is the pH of a neutral solution ([H3O+] = 1.0x10-7M)? • pH = 7 • What is the pH of a solution if: • [H3O+] = 1.0x10-4M? • pH = 4 • [H3O+] = 0.0015M? • pH = 2.8

  20. Calculating pOH • pOHis a measure of hydroxide [OH-] concentration in solution • Higher pOH= more acidic • pOH = - log [OH-] • pH + pOH = 14

  21. pOH Calculations • What is the pOH of a neutral solution ([OH-] = 1.0x10-7M)? • What is the pOH of a solution if: • [OH-] = 8.22x10-6M? • pOH = 5.09 • [OH-] = 0.0541M? • pOH = 1.27

  22. [H+] 100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14 pH 0 1 3 5 7 9 11 13 14 14 13 11 9 7 5 3 1 0 pOH 10-14 10-13 10-11 10-9 Basic 10-7 10-5 10-3 10-1 100 [OH-] Acidic Neutral Basic

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