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Acids and Bases

Acids and Bases. Acids and Bases. In your groups, separate your cards into three categories. Acids = Proton (H + ) Donors Bases = Proton (H + ) Acceptors Both (Acids and Bases) Why is H + a proton? H + has 1 proton, 0 neutrons, 0 electrons. Acids. pH below 7. turns litmus paper red.

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Acids and Bases

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  1. Acids and Bases

  2. Acids and Bases • In your groups, separate your cards into three categories. • Acids = Proton (H+) Donors • Bases = Proton (H+) Acceptors • Both (Acids and Bases) Why is H+ a proton? H+ has 1 proton, 0 neutrons, 0 electrons.

  3. Acids pH below 7 turns litmus paper red taste sour reacts with metals to produce H2(g) generally starts with a hydrogen ion [H+] > [OH-] HCl

  4. Bases pH greater than 7 turns litmus paper blue taste bitter feel slippery generally contains a hydroxide ion [H+] < [OH-] NaOH

  5. BothAcids and Bases amphoteric an electrolyte contains H+ and OH- ions

  6. Water molecules are polar and are in constant motion. Sometimes the collisions between two water molecules have enough energy to cause to transfer of a hydrogen. • This reaction in which two water molecules react to form ions is called the self-ionization of water

  7. H2O H2O H3O+ OH- H3O+ (hydronium ion) is often simplified to H+

  8. In pure water the hydrogen ion concentration [H+] and the hydroxide ion concentration [OH-] are both 1.0 x 10-7 M. • A solution where [H+] and [OH-] are equal is called a neutralsolution.

  9. For all aqueous solution the following is true: [H+] [OH-] = 1.0 x 10-14 this is Kw – ion product constant for H20

  10. [H+] & [OH-] are interdependent. If [H+] goes up then [OH-] goes down & vice versa • When [H+] is greater than [OH-] the solution is considered ACIDIC. • When [H+] is less than the [OH-] the solution is considered BASIC. • Basic solutions are also known as alkaline solutions.

  11. The[H+] of a solution is 1.0 x 10-3. What is the[OH-] in this solution? Is the solution acidic, basic or neutral? [H+] [OH-] = 1.0 x 10-14 1.0 x 10-3[OH-] = 1.0 x 10-14 [OH-] = 1.0 x 10-14 1.0 x 10-3 [OH-] = 1.0 x 10-11 Acidic!

  12. The[H+] of a solution is 3.5 x 10-9. What is the[OH-] in this solution? Is the solution acidic, basic or neutral? [H+] [OH-] = 1.0 x 10-14 3.5 x 10-9[OH-] = 1.0 x 10-14 [OH-] = 1.0 x 10-14 3.5 x 10-9 [OH-] = .29 x 10-5 or 2.9 x 10-6 Basic!

  13. Are the following solutions acidic, basic, or neutral? • [H+] = 6.0 x 10-10 • [H+] = 3.0 x 10-2 • [OH-] = 1.0 x 10-7 • [OH-] = 2.0 x 10-7 Basic Acidic Neutral Basic

  14. Acidic, Basic, and Neutral Solutions

  15. pH Scale More Basic More Acidic

  16. Acid-Base Indicators Acid-base indicators are substances that change their color based on the pHof a solution. Examples of acid-base indicators include the following: Litmus Paper pH Paper Phenolphthalein Universal Indicator Bromothymol Blue Some Fruits and Vegetables

  17. Litmus and pH Paper

  18. The pH of a solution is the –log of the hydrogen ion concentration • pH = - log[H+] • The pOH of a solution is the –log of the hydroxide ion concentration • pOH = - log[OH-] • pH + pOH = 14

  19. What is the pH of a solution with a [H+] of 1.0x10-7? pH = -log[1.0x10-7] pH = 7 Neutral!

  20. What is the pH of a solution with a [OH-] of 1.0x10-11? pOH = -log[1.0x10-11] pOH = 11 pH = 3 Acidic!

  21. What is the pH of a solution with a [H+] of 1.0x10-9? pH = -log[1.0x10-9] pH = 9 Basic!

  22. What is the pH of a solution with a [H+] of 2.5x10-4? pH = -log[2.5x10-4] pH = 3.6 Acidic!

  23. What is the pH of a solution with a [OH-] of 3.0x10-11? pOH = -log[3.0x10-11] pOH = 10.52 pH = 3.48 Acidic!

  24. Arrhenius Acids and Bases • Acids- hydrogen containing compounds that ionize to yield hydrogen ions in aqueous solutions. • Bases- hydroxide containing compounds that ionize to yield hydroxide ions in aqueous solutions. • Not all compounds containing hydrogen are arrhenius acids, because only hydrogen ions in very polar bonds are ionizable.

  25. Arrhenius Acids and Bases • Example: HCl, H: 2.1 and Cl: 3.0 very polar bond, so it is ionizable  +  - H2O H – Cl H+(aq)+ Cl-(aq)

  26. Arrhenius Acids and Bases • Example: CH4 • C: 2.5 and H: 2.1 • C-H bonds are weakly polar, so the hydrogen’s are not ionizable. • CH4 is not an acid by Arrhenius definition.

  27. Common Arrhenius Bases • Group 1A bases are very soluble in water and produce very concentrated solutions with ease. These concentrated solutions cause very deep, painful, slow-healing wounds. Example: KOH and NaOH • Group 2A bases are not very soluble in water and therefore are very dilute. These bases are used in antacids and laxatives. Ex: Ca(OH)2 and Mg(OH)2

  28. Bronsted-Lowry Acids and Bases • Acid – hydrogen-ion donor • Base – hydrogen ion acceptor • All acids and bases under Arrhenius definition are also Bronsted-Lowry bases. • Arrhenius definitions are very narrow and do not include all acids and bases. Therefore, not all Bronsted-Lowry acids and bases are Arrhenius acids and bases.

  29. Bronsted-Lowry Acids and Bases • Example: NH3 • NH3 + H2O  NH4++ OH- • NH3 accepts a H+ • H2O donates a H+ BASE ACID BASE ACID

  30. Bronsted-Lowry Acids and Bases • Water acted as an acid in the previous example. HCl + H2O  Cl- + H3O+ Here water is acting as a bronsted-lowry base. Water is ACID BASE amphoteric!!!

  31. Lewis Acids and Bases • Definition is more general then both Arrhenius and Bronsted-Lowry definitions. • Lewis Acid- accepts a pair of electrons to form a covalent bond • Lewis Base- donates a pair of electrons to form a covalent bond.

  32. Lewis Acids and Bases

  33. Strong Acids • HClO4 • HClO3 • HCl • HBr • HI • HNO3 • H2SO4 Memorize These!!!! Strong acids completely ionize in aqueous solutions, while weak acids only slightly ionize in aqueous solutions.

  34. Strong Bases Group I and II elements with hydroxide ion. Examples: KOH and Sr(OH)2

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