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Chapter 10 Acids and Bases

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Chapter 10 Acids and Bases

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    1. Chapter 10 Acids and Bases John D. Bookstaver St. Charles Community College St. Peters, MO ? 2006, Prentice Hall, Inc.

    3. Polyprotic Acids

    4. Arrhenius Acids and Bases In 1884, Svante Arrhenius proposed these definitions acid: a substance that produces H3O+ ions aqueous solution base: a substance that produces OH- ions in aqueous solution this definition of an acid is a slight modification of the original Arrhenius definition, which was that an acid produces H+ in aqueous solution today we know that H+ reacts immediately with a water molecule to give a hydronium ion

    5. Arrhenius Acids and Bases when HCl, for example, dissolves in water, its reacts with water to give hydronium ion and chloride ion

    6. Arrhenius Acids and Bases With bases, the situation is slightly different many bases are metal hydroxides such as KOH, NaOH, Mg(OH)2, and Ca(OH)2 these compounds are ionic solids and when they dissolve in water, their ions merely separate other bases are not hydroxides; these bases produce OH- by reacting with water molecules

    7. Acid and Base Strength Strong acid: one that reacts completely or almost completely with water to form H3O+ ions Strong base: one that reacts completely or almost completely with water to form OH- ions here are the six most common strong acids and the four most common strong bases

    15. Acid and Base Strength Strong acids are completely dissociated in water. Weak acids only dissociate partially in water.

    16. Acid and Base Strength Substances with negligible acidity do not dissociate in water.

    17. Acid and Base Strength Weak acid: a substance that dissociates only partially in water to produce H3O+ ions acetic acid, for example, is a weak acid; in water, only 4 out every 1000 molecules are converted to acetate ions Weak base: a substance that dissociates only partially in water to produce OH- ions ammonia, for example, is a weak base

    18. Some Definitions Brønsted–Lowry Acid: Proton donor Base: Proton acceptor

    19. Strong Acids You will recall that the seven strong acids are HCl, HBr, HI, HNO3, H2SO4, HClO3, and HClO4. These are, by definition, strong electrolytes and exist totally as ions in aqueous solution. For the monoprotic strong acids, [H3O+] = [acid].

    20. Strong Bases Strong bases are the soluble hydroxides, which are the alkali metal and heavier alkaline earth metal hydroxides (Ca2+, Sr2+, and Ba2+). Again, these substances dissociate completely in aqueous solution.

    23. What Happens When an Acid Dissolves in Water? Water acts as a Brønsted–Lowry base and abstracts a proton (H+) from the acid. As a result, the conjugate base of the acid and a hydronium ion are formed.

    24. Conjugate Acids and Bases: From the Latin word conjugare, meaning “to join together.” Reactions between acids and bases always yield their conjugate bases and acids.

    27. Brønsted-Lowry Acids & Bases Acid: a proton donor Base: a proton acceptor Acid-base reaction: a proton transfer reaction Conjugate acid-base pair: any pair of molecules or ions that can be interconverted by transfer of a proton

    28. Brønsted-Lowry Acids & Bases Brønsted-Lowry definitions do not require water as a reactant

    29. Brønsted-Lowry Acids & Bases Note the following about the conjugate acid-base pairs in the table 1. an acid can be positively charged, neutral, or negatively charged; examples of each type are H3O+, H2CO3, and H2PO4- 2. a base can be negatively charged or neutral; examples are OH-, Cl-, and NH3 3. acids are classified a monoprotic, diprotic, or triprotic depending on the number of protons each may give up; examples are HCl, H2CO3, and H3PO4

    30. Brønsted-Lowry Acids & Bases carbonic acid, for example can give up one proton to become bicarbonate ion, and then the second proton to become carbonate ion 4. several molecules and ions appear in both the acid and conjugate base columns; that is, each can function as either an acid or a base

    31. If it can be either… ...it is amphiprotic. HCO3- HSO4- H2O

    32. pH pH is defined as the negative base-10 logarithm of the hydronium ion concentration. pH = -log [H3O+]

    33. pH Therefore, in pure water, pH = -log (1.0 ? 10-7) = 7.00 An acid has a higher [H3O+] than pure water, so its pH is <7 A base has a lower [H3O+] than pure water, so its pH is >7.

    35. Other “p” Scales The “p” in pH tells us to take the negative log of the quantity (in this case, hydrogen ions). Some similar examples are pOH -log [OH-] pKw -log Kw

    38. pH These are the pH values for several common substances.

    39. How Do We Measure pH? For less accurate measurements, one can use Litmus paper “Red” paper turns blue above ~pH = 8 “Blue” paper turns red below ~pH = 5 An indicator

    40. How Do We Measure pH? For more accurate measurements, one uses a pH meter, which measures the voltage in the solution.

    42. Brønsted-Lowry Acids & Bases the HCO3- ion, for example, can give up a proton to become CO32-, or it can accept a proton to become H2CO3 a substance that can act as either an acid or a base is said to be amphiprotic the most important amphiprotic substance in Table 8.2 is H2O; it can accept a proton to become H3O+, or lose a proton to become OH- 5. a substance cannot be a Brønsted-Lowry acid unless it contains a hydrogen atom, but not all hydrogen atoms in most compounds can be given up acetic acid, for example, gives up only one proton

    43. Brønsted-Lowry Acids & Bases 6. there is an inverse relationship between the strength of an acid and the strength of its conjugate base the stronger the acid, the weaker its conjugate base HI, for example, is the strongest acid in Table 8.2, and its conjugate base, I-, is the weakest base in the table CH3COOH (acetic acid) is a stronger acid that H2CO3 (carbonic acid); conversely, CH3COO- (acetate ion) is a weaker base that HCO3- (bicarbonate ion)

    44. Properties of Acids & Bases Neutralization acids and bases react with each other in a process called neutralization; these reactions are discussed in Section 8.10 Reaction with metals strong acids react with certain metals (called active metals) to produce a salt and hydrogen gas, H2 reaction of a strong acid with a metal is a redox reaction; the metal is oxidized to a metal ion and H+ is reduced to H2

    45. Properties of Acids & Bases Reaction with metal hydroxides reaction of an acid with a metal hydroxide gives a salt plus water the reaction is more accurately written as omitting spectator ions gives this net ionic equation

    46. Properties of Acids & Bases Reaction with metal oxides strong acids react with metal oxides to give water plus a salt

    47. Properties of Acids & Bases Reaction with carbonates and bicarbonates strong acids react with carbonates to give carbonic acid, which rapidly decomposes to CO2 and H2O strong acids react similarly with bicarbonates

    50. Acid-Base Titrations Titration: an analytical procedure in which a solute in a solution of known concentration reacts with a known stoichiometry with a substance whose concentration is to be determined in this chapter, we are concerned with titrations in which we use an acid (or base) of known concentration to determine the concentration of a base (or acid) in another solution

    51. Titration A known concentration of base (or acid) is slowly added to a solution of acid (or base).

    52. Titration A pH meter or indicators are used to determine when the solution has reached the equivalence point, at which the stoichiometric amount of acid equals that of base.

    53. Acid-Base Titrations As an example, let us use 0.108 M H2SO4 to determine the concentration of a NaOH solution requirement 1: we know the balanced equation requirement 2: the reaction between H3O+ and OH- is rapid and complete requirement 3: we can use either an acid-base indicator or a pH meter to observe the sudden change in pH that occurs at the end point of the titration requirement 4: we use volumetric glassware

    54. Acid-Base Titrations experimental measurements doing the calculations

    55. Buffers: Solutions of a weak conjugate acid-base pair. They are particularly resistant to pH changes, even when strong acid or base is added.

    56. Blood Buffers The average pH of human blood is 7.4 any change larger than 0.10 pH unit in either direction can cause illness To maintain this pH, the body uses three buffer systems carbonate buffer: H2CO3 and its conjugate base, HCO3- phosphate buffer: H2PO4- and its conjugate base, HPO42- proteins: discussed in Chapter 21

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