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Chapter 19 Acids & Bases

Chapter 19 Acids & Bases. Section 19.1 Acid – Base Theories. Acids vinegar  citrus fruits carbonated drinks  car battery lemon juice  tea. Bases calcium hydroxide in mortar  antacids household cleaning agents. Properties of Acids. Give foods a tart or sour taste

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Chapter 19 Acids & Bases

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  1. Chapter 19Acids & Bases Section 19.1Acid – Base Theories

  2. Acids • vinegar  citrus fruits • carbonated drinks  car battery • lemon juice  tea • Bases • calcium hydroxide in mortar  antacids • household cleaning agents

  3. Properties of Acids • Give foods a tart or sour taste • e.g. lemon & vinegar for example • Aqueous solutions of acids are electrolytes (conduct electricity) • Acids cause certain chemical indicators to change color. • Acid + Base Salt + water

  4. Properties of Bases • Bases have a bitter taste • e.g. soap • Bases have a slippery feel • Aqueous solutions of bases are electrolytes (conduct electricity) • Bases cause certain chemical indicators to change color. • Acid + Base Salt + water

  5. Arrhenius Acids & Bases In 1887, Swedish chemist Svante Arrhenius proposed a revolutionary way of defining and thinking about acids and bases • Acids are hydrogen-containing compounds that ionize to yield hydrogen ions (H+) in aqueous solution. • Bases are compounds that ionize to yield hydroxide ions (OH-) in aqueous solution

  6. Monoprotic acids: acids that contain one ionizable hydrogen HNO3 – nitric acid Diprotic acids: acids that contain two ionizable hydrogens H2SO4 – sulfuric acid Triprotic acids:acids that contain three ionizablehydrogens H3PO4 – phosphoric acid

  7. Not all compounds that contain hydrogen are acids • e.g. CH4 – methane has weak polar C – H bonds and no ionizable hydrogens. Not an acid. • Not all hydrogens in an acid may be released as hydrogen ions. • Only hydrogens in very polar bonds are ionizable. In the case where hydrogen is joined to a very electronegative element. • e.g. HCl hydrogen chloride very polar covalent molecule

  8. When HCl dissolves in water, it releases hydrogen ions because the hydrogen ions are stabilized by solvation. H2O H – Cl(g) H+(aq) + Cl-(aq) Hydrogen Hydrogen Chloride chloride ion ion Ionizes to form an aqueous solution of hydronium ions and chloride ions HCl + H2O H3O+ + Cl-

  9. Ethanoic acid CH3COOH is a monoprotic acid due to its structure • H O • H C C O H • H • The three H attached to the carbon are in weak polar bonds. They do not ionize. • Only the H bonded to the highly electronegative O can be ionized

  10. Sodium hydroxide dissociates into sodium ions and hydroxide ions in aqueous solution. H2O NaOH(s) Na+(aq) + OH- (aq) Sodium Sodium Hydroxide Hydroxide Ion ion Potassium hydroxide dissociates into potassium ions and hydroxide ions in aqueous solution. H2O KOH (s) K+(aq) + OH- (aq) Potassium Potassium Hydroxide Hydroxide Ion ion

  11. Arrhenius Bases Group one, the alkali metals, react with water to produce solutions that are basic. Group one metals are very soluble in water and can produce concentrated solutions. Group two metals are not very soluble in water. Their solutions are always very dilute.

  12. Bronsted – Lowry Acids and Bases The Bronsted – Lowry theory defines Acid: a hydrogen-ion donor Base: a hydrogen-ion acceptor All acids and bases included in the Arrhenius theory are also acids and bases according to the Bronsted-Lowry theory.

  13. NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq) • base acid conjugate conjugate • acid base • Ammonia is the hydrogen-ion acceptor • therefore it is a base. • Water is the hydrogen-ion donor and • therefore it is an acid. • Hydrogen ions are transferred from water to ammonia, which causes the hydroxide-ion concentration to be greater than it is in pure water. • When ammonia dissolves and reacts with water, NH4+ is the conjugate acid of the base NH3. • OH- is the conjugate base of acid H2O

  14. Conjugate Acids and Bases • HCl(g) + H2O (l) H3O+(aq) + Cl-(aq) • acid base conjugateconjugate • acid base • HCl is the hydrogen-ion donor – thus it is an acid. • Water is the hydrogen-ion acceptor – thus it is a base

  15. Conjugate Acid – Base Pair • Conjugate acid:the particle formed when a base gains a hydrogen ion • Conjugate base: the particle that remains when an acid has donated a hydrogen ion. • Conjugate acids and bases are always paired with a base or an acid, respectively. • Conjugate acid-base pairs consists of two substances related by the loss or gain of a single hydrogen ion.

  16. Common Conjugate Acid – Base Pairs

  17. Bronsted – Lowry Acids and Bases A water molecule that gains a hydrogen ion becomes a positively charged hydronium ion(H3O+) Amphoteric – a substance that can act as both an acid and a base e.g. water H2SO4 + H2O H3O+ + HSO4- NH3 + H2O NH4+ + OH-

  18. Lewis Acids and Bases Gilbert Lewis proposed a third Acid Base theory Acid– accepts a pair of electrons during a reaction Base –donates a pair of electrons during a reaction Concept is more general than either the Arrhenius theory or the Bronsted-Lowry theory.

  19. Lewis Acids and Bases Lewis Acid – a substance that can accept a pair of electrons to form a covalent bond. Lewis Base –a substance that can donate a pair of electrons to form a covalent bond. .. H+: O – H :O:.. H H Lewis Lewis Acid Base

  20. Acid Base Definitions

  21. End of Section 19.1

  22. Section 19.2Hydrogen Ions & Acidity

  23. Hydrogen Ions From Water A water molecule that loses a hydrogen ion becomes a negatively charged hydroxide ion A water molecule that gains a hydrogen ion becomes a positively charged hydronium ion H2O (l) OH- (aq) + H+ (aq) Hydroxide ion Hydrogen ion Self ionization of water – the reaction in which water molecules produce ions

  24. Self Ionization of Water • Hydrogen ions in aqueous solution have several names. • Some chemists call them protons • Some chemists call them hydrogen ions or hydronium ions. • For our purposes, either H+ or H3O+ will represent hydrogen ions in aqueous solution. • H2O + H2O H3O+ + OH-

  25. In pure water at 25˚C, the equilibrium concentration of hydrogen ions and hydroxide ions are each only 1 x 10-7. • In other words the concentration of OH- and H+ are equal in pure water Any aqueous solution in which H+ and OH- are equal is a neutral solution. When [H+] increases [OH-] decreases When [H+] decreases [OH-] increases

  26. For aqueous solutions, the product of the hydrogen ion concentration and the hydroxide ion concentration equals 1.0 x 10-14 [H+] x [OH-] = 1 x 10-14 1 x 10-7 x 1 x 10-7 = 1 x 10-14 This equation is true for all dilute aqueous solutions at 25˚C. Ion-Product Constant for Water (Kw) theproduct of the concentrations of the hydrogen ions and hydroxide ions in water Kw = [H+] x [OH-] = 1.0 x 10-14

  27. But not all solutions are neutral When some substances dissolve in water, they release hydrogen ions. When hydrogen chloride dissolves in water, it forms hydrochloric acid. H2O HCl (g) H+ (aq) + Cl- (aq)

  28. Ion Product Constant for Water In the previous HCl solution, the hydrogen-ion concentration is greater than the hydroxide-ion concentration. Acidic Solution: A solution in which [H+] is greater than [OH-]. The [H+] of an acidic solution is greater than 1 x 10-7 M

  29. When sodium hydroxide dissolves in water, it forms hydroxide ions in solution. H20 NaOH(s) Na+(aq) + OH-(aq) In the above solution, the hydrogen-ion concentration is less than the hydroxide-ion concentration. Basic Solution: A solution in which [H+] is less than [OH-] The [H+] of a basic solution is less than 1 x 10-7 Basic solutions are also known as alkaline solutions.

  30. The pH Concept The pH scale was proposed by Danish Scientist Soren Sorensen in 1909. The pH scale is used to express [H+] 1 2 3 4 5 6 7 8 9 10 11 12 13 14 Strongly Neutral Strongly Acidic Basic

  31. Calculating pH The pH of a solution is the negative logarithm of the hydrogen-ion concentration. pH = - log [H+]

  32. In neutral solution, the [H+] = 1 x 10-7M pH = - log [H+] pH = - log (1 x 10-7) pH = 7

  33. Classifying Solutions A solution in which [H+] is greater than 1 x 10-7 has a pH less than 7 is called acidic. A solution in which [H+] is less than 1 x 10-7 has a pH greater than 7 is called basic. The pH of pure water or a neutral aqueous solution is 7 Acidic solution: pH < 7.0 [H+] > 1 x 10-7 M Neutral solution: pH = 7.0 [H+] = 1 x 10-7 M Basic solution: pH > 7.0 [H+] < 1 x 10-7 M

  34. pH can be read from the value of [H+] if it is written in scientific notation and has a coefficient of 1. Then the pH of the solution equals the exponent, with the sign changed from minus to plus e.g. [H+] = 1 x 10-2 has a pH of 2 e.g. [H+] = 1 x 10-13 has a pH of 13

  35. If the pH is an integer, it is also possible to directly write the value of [H+]. pH = 9 ; then [H+] = 1 x 10-9 M pH = 4 ; then [H+] = 1 x 10-4M

  36. The pOH of a solution equals the negative logarithm of the hydroxide-ion concentration pOH = - log [OH-] A neutral solution has a pOH of 7 Acidic solution: pOH > 7 [OH-] < 1 x 10-7 M Neutral solution: pOH = 7 [OH-] = 1 x 10-7 M Basic solution: pOH < 7 [OH-] > 1 x 10-7 M

  37. pH and pOH Relationship pOH + pH = 14 pH = 14 – pOH pOH = 14 – pH

  38. Problem Example Colas are slightly acidic. If the [H+] in a solution is 1.0 X 10 - 5 M , is the solution acidic, basic or neutral. What is the [OH-] of this solution? [H+] = 1.0 X 10 - 5 M which is greater than 1.0 X 10 -7M so the solution is acidic Kw = [OH-] x [H+] = 1.0 X 10 - 14 [OH-] = 1.0 X 10 - 14 ÷[H+] [OH-] = 1.0 X 10 - 14 ÷ 1.0 X 10 - 5 [OH-] = 1.0 X 10 - 9 M

  39. Problem Example What is the pH of a solution with a hydrogen-ion concentration of 4.2 x 10 - 10 M? pH = - log [H+] pH = - log (4.2 x 10 - 10) pH = 9.38

  40. Problem Example pH of an unknown solution is 6.35. What is its hydrogen-ion concentration? pH = -log [H+] 6.35 = -log [H+] - 6.35 = log [H+] Using calculator find the anti log of - 6.35 [H+] = 4.5 x 10 - 7 M

  41. Problem Example What is the pH of a solution if the [OH-] = 4.0 X10 – 11 M? Kw = [H+] x [OH-] = 1 x 10 -14 [H+] = 1 x 10 -14 ÷ [OH-] [H+] = 1 x 10 -14 ÷ 4.0 x 10 -11 [H+] = 2.5 x 10 - 4 M

  42. Problem Example (con’t) What is the pH of a solution if [OH-] = 4.0 X 10 - 11 M? pH = - log [H+] pH = - log (2.5 x 10 - 4) pH = 3.60

  43. Acid – Base Indicators Indicator - is an acid or a base that undergoes dissociation in a known pH range An indicator is a valuable tool for measuring pH because its acid form and base form have different color in solution. For each indicator, the change from dominating acid form to dominating base form occurs in a narrow range of approximately two pH units. Within this range, the color of the solution is a mixture of the colors of the acid and the base forms. Knowing the pH range over which this color change occurs, can give you a rough estimate of the pH of the solution.

  44. End of Section 19.2

  45. Section 19.3Hydrogen Ions & Acidity

  46. Strong Acids • Acids are classified as strong or weak depending on the degree to which they ionize in water. • In general, strong acids are completely ionized in aqueous solution. • HNO3 - nitric acid HCl - hydrochloric acid H2SO4 - sulfuric acid HClO4 - perchloric acid HBr - hydrobromic acid HI - hydroiodic acid • HCl(g) + H2O(l) H3O +(aq) + Cl¯(aq)

  47. Weak Acids • Weak acids ionize only slightly in aqueous solution. • Some Weak Acids • Acetic Acid CH3COOH • Boric Acid H3BO3 (all three are weak) • Phosphoric Acid H3PO4 (all three are weak) • Sulfuric Acid HSO4- (first ionization is strong) • CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO ¯(aq) • ethanoic acid water hydronium ethanoate • ion ion

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