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Acids and Bases

This guide explores the fundamental concepts of acids and bases, focusing on their definitions as proton donors and acceptors. It covers key properties, including pH levels, the role of hydrogen ions (H+) and hydroxide ions (OH-), and the self-ionization of water. You will learn to classify solutions as acidic, basic, or neutral based on ion concentrations and calculate pH and pOH. The guide also introduces acid-base indicators and the distinctions between Arrhenius, Bronsted-Lowry, and Lewis definitions of acids and bases.

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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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