Chemistry—Ch. 19: Acids, Bases, and Salts - PowerPoint PPT Presentation

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Chemistry—Ch. 19: Acids, Bases, and Salts
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Chemistry—Ch. 19: Acids, Bases, and Salts

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  1. 19.1—Acid-Base Theoriesmain ideas: properties of acids/bases, Arrhenius, Bronsted-Lowry 19.2—Hydrogen Ions and Aciditymain ideas: hydrogen ions from water, pH, indicators 19.3—Strengths of Acids and Basesmain ideas: what makes acid “strong” or “weak” 19.4—Neutralization Reactionsmain ideas: acid/base yields salt/water,titration, endpoint 19.5—Salts in Solutionmain idea: what is a buffer Chemistry—Ch. 19: Acids, Bases, and Salts

  2. 19.1 19.1 Acid-Base Theories Bracken Cave, near San Antonio, Texas, is home to twenty to forty million bats. Visitors to the cave must protect themselves from the dangerous levels of ammonia in the cave. Ammonia is a byproduct of the bats’ urine. You will learn why ammonia is considered a base.

  3. 19.1 Properties of Acids and Bases • Acids taste sour • pH less than 7 • Change blue litmus paper red • Red to orange color with universal indicator • Contain excess H+ ions • Corrosive to metal and skin • React with bases to form water and a salt • Citrus fruits contain citric acid. Tea contains tannic acid.

  4. Properties of Acids and Bases • Bases taste bitter • feel slippery • pH greater than 7 • change red litmus paper blue • Blue with universal indicator; pink with phenolphthalein • Caustic • React with acids to form water and a salt • Antacids use bases to neutralize excess stomach acid. The base calcium hydroxide is a component of mortar.

  5. 19.1 Properties of Acids and Bases • Both acids and bases can be strong or weak electrolytes in aqueous solution. • Conductivity tests the extent of ionization in water; weak acids and bases do not ionize completely whereas strong acids and bases do • Examples: • Strong acid is HCl; weak acid is CH3COOH (also known as HCH3COO or HC2H3O2) • Strong base is NaOH; weak base is NH3 (also known as NH4OH)

  6. 19.1 Acids and Bases There are several definitions of an acid and a base. Arrhenius is a simple, but limited definition. The Lewis definition is much more comprehensive. Some substances would be considered an acid using one definition, but not using another.

  7. 19.1 Arrhenius Acids and Bases Arrhenius said that acids are hydrogen-containing compounds that ionize to yield hydrogen ions (H+) in aqueous solution. He also said that bases are compounds that ionize to yield hydroxide ions (OH–) in aqueous solution.

  8. 19.1 Arrhenius Acids and Bases Arrhenius Acids • Acids that contain one ionizable hydrogen, such as nitric acid (HNO3), are called monoprotic acids. • Acids that contain two ionizable hydrogens, such as sulfuric acid (H2SO4), are called diprotic acids. • Acids that contain three ionizable hydrogens, such as phosphoric acid (H3PO4) are called triprotic acids.

  9. 19.1 Arrhenius Acids and Bases Arrhenius Bases Milk of magnesia is a base used as an antacid.

  10. 19.1 Brønsted-Lowry Acids and Bases • The Brønsted-Lowry theory defines an acid as a hydrogen-ion donor, and a base as a hydrogen-ion acceptor.

  11. 19.1 Brønsted-Lowry Acids and Bases • Conjugate Acids and Bases • 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. The acid and base are always on the left with the corresponding conjugate acid-base pair on the right.

  12. 19.1 Brønsted-Lowry Acids and Bases • A conjugate acid-base pair consists of two substances related by the loss or gain of a single hydrogen ion. • A substance that can act as both an acid and a base is said to be amphoteric.

  13. 19.1 Brønsted-Lowry Acids and Bases • A water molecule that gains a hydrogen ion becomes a positively charged hydronium ion (H3O+).

  14. 19.1 Lewis Acids and Bases • Lewis proposed that an acid accepts a pair of electrons during a reaction, while a base donates a pair of electrons. • A Lewis acid is a substance that can accept a pair of electrons to form a covalent bond. • A Lewis base is a substance that can donate a pair of electrons to forma covalent bond.

  15. 19.1 Section Quiz. 1. Which of the following is NOT a characteristic of acids? a) taste sour b) are electrolytes c) feel slippery d) affect the color of indicators

  16. 19.1 Section Quiz. 2.Which compound is most likely to act as an Arrhenius acid? a) H2O b) NH3. c) NaOH. d) H2SO4.

  17. 19.1 Section Quiz. 3. A Lewis acid is any substance that can accept a) a hydronium ion. b) a proton. c) hydrogen. d) a pair of electrons.

  18. 19.2 19.2 Hydrogen Ions and Acidity To test a diagnosis of diabetic coma, a doctor orders several tests, including the acidity of the patient’s blood. Results from this test will be expressed in units of pH. You will learn how the pH scale is used to indicate the acidity of a solution and why the pH scale is used.

  19. 19.2 Hydrogen Ions from Water The reaction in which water molecules produce ions is called the self-ionization of water. • In the self-ionization of water, a proton (hydrogen ion) transfers from one water molecule to another water molecule.

  20. Ion Product Constant for Water The product of the concentrations of the hydrogen ions and hydroxide ions in water is called the ion-product constant for water (Kw). • For aqueous solutions, the product of the hydrogen-ion concentration and the hydroxide-ion concentration equals 1.0  10-14. • Any aqueous solution in which [H+] and [OH-] are equal is described as a neutral solution. • The [ ] represents concentration of the ions. The appropriate unit is M (molarity).

  21. 19.2 Ion Product Constant for Water • A basic solution is one in which [H+] is less than [OH]. Basic solutions are also known as alkaline solutions. An acidic solution is one in which [H+] is greater than [OH-]. • Unrefined hydrochloric acid, commonly called muriatic acid, is used to clean stone buildings and swimming pools. • Sodium hydroxide, or lye, is commonly used as a drain cleaner.

  22. 19.2 The pH Concept • pH and Significant Figures 10-14 [OH-] 100 • A solution in which [H+] is greater than 1  10–7M has a pH less than 7.0 and is acidic. The pH of pure water or a neutral aqueous solution is 7.0. A solution with a pH greater than 7 is basic and has a [H+] of less than 1  10–7M.

  23. The pH Concept

  24. The pH Concept

  25. 19.2 The pH Concept Kw= [H+][OH-]= 1.0 x 10-14 Use these equations for pH problems. Not written exactly in this manner in the textbook.

  26. 19.2 Measuring pH • Universal Indicator An indicator is a valuable tool for measuring pH because its acid form and base form have different colors in solution. • Phenolphthalein changes from colorless to pink at pH 7–9.

  27. 19.2 Measuring pH pH Meters

  28. 19.2 Section Quiz. 1. If the [OH-] in a solution is 7.65  10-3M, what is the [H+] of this solution? a) 7.65  10-17M b) 1.31  10-12M c) 2.12M d) 11.88M

  29. 19.2 Section Quiz. 2. The [OH-] for four solutions is given below. Which one of the solution is basic? a) 1.0 x 10-6M b) 1.0 x 10-8M c) 1.0 x 10-7M d) 1.0 x 10-14M

  30. 19.2 Section Quiz. 3. What is the pH of a solution with a hydrogen-ion concentration of 8.5 x 10-2M? a) 12.93 b) 8.50 c) 5.50 d) 1.07

  31. 19.3 19.3 Strengths of Acids and Bases Lemons and grapefruits have a sour taste because they contain citric acid. Sulfuric acid is a widely used industrial chemical that can quickly cause severe burns if it comes into contact with skin. You will learn why some acids are weak and some acids are strong.

  32. 19.3 Strong and Weak Acids and Bases • An acid dissociation constant (Ka) is the ratio of the concentration of the dissociated (or ionized) form of an acid to the concentration of the undissociated (nonionized) form. • Weak acids have small Ka values. The stronger an acid is, the larger is its Ka value.

  33. 19.3 Strong and Weak Acids and Bases • Strong acids are completely ionized in aqueous solution. • Weak acids ionize only slightly in aqueous solution.

  34. 19.3 Strong and Weak Acids and Bases • In general, the base dissociation constant (Kb) is the ratio of the concentration of the conjugate acid times the concentration of the hydroxide ion to the concentration of the base. • Strong bases dissociate completely into metal ions and hydroxide ions in aqueous solution. • Weak bases react with water to form the hydroxide ion and the conjugate acid of the base.

  35. 19.3 Strong and Weak Acids and Bases

  36. 19.3 Calculating Dissociation Constants • How can you calculate an acid dissociation constant (Ka) of a weak acid? • Ex. The dissociation constant, Ka, of ethanoic acid is calculated from the equilibrium concentrations of all of the molecules and ions in the solution.

  37. 19.3 Section Quiz. 1. H2S is considered to be a weak acid because it a) is insoluble in water. b) ionizes only slightly. c) is completely ionized. d) is dilute.

  38. 19.3 Section Quiz. 2. Calcium hydroxide, Ca(OH)2, is a strong base because it a) has a large Kb. b) has a small Kb. c) forms concentrated solutions. d) is highly soluble in water.

  39. 19.3 Section Quiz. 4. The Ka of three acids is given below. (1) 5.1  10–3 (2) 4.8  10–11 (3) 6.3  10–5 Put the acids in order from the strongest acid to the weakest acid. a) 1, 3, 2 b) 2, 3, 1 c) 3, 1, 2 d) 2, 1, 3

  40. 19.4 19.4 Neutralization Reactions Excess hydrochloric acid in the stomach can cause heartburn and a feeling of nausea. Antacids neutralize the stomach acid and relieve the pain of acid indigestion. You will learn what a neutralization reaction is.

  41. 19.4 Acid-Base Reactions • In general, the reaction of an acid with a base produces water and one of a class of compounds called salts. • Reactions in which an acid and a base react in an aqueous solution to produce a salt and water are generally called neutralization reactions.

  42. 19.4 Acid-Base Reactions

  43. 19.4 Titration The process of adding a known amount of solution of known concentration to determine the concentration of another solution is called titration. • When an acid and base are mixed, the equivalence point is when the number of moles of hydrogen ions equals the number of moles of hydroxide ions.

  44. 19.4 Titration • The solution of known concentration is called the standard solution. • Indicators are often used to determine when enough of the standard solution has been added to neutralize the acid or base. • The point at which the indicator changes color is the end point of the titration. Added base is measured with a buret. Acid solution with indicator Color change shows neutralization.