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Acids and Bases

Acids and Bases

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Acids and Bases

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  1. Acids and Bases Acids and Bases 2007-2008

  2. Acids • Svante Arrhenius, a Swedish chemist, defines an acid as a substance that yields hydrogen ions (H+) when dissolved in water. • Formulas for acids contain one or more ionizable hydrogen atoms as well as an anion. Acids and Bases 2007-2008

  3. Naming Acids In some cases two different names seem to be assigned to the same chemical formula. HCl(g) hydrogen chloride HCl(l) hydrogen chloride HCl(aq) hydrochloric acid The name assigned to the compound depends on its physical state. In the gaseous or pure liquid state, HCl is a molecular compound called hydrogen chloride. When it is dissolved in water, the molecules break apart into H+ and Cl- ions; in this state, the substance is called hydrochloric acid. Acids and Bases 2007-2008

  4. Binary acids (formed by hydrogen and one other element) are named with a “hydro-” prefix and an “-ic” ending on the anion root. Ex- HCl HBr Hydrochloric acid Hydrobromic acid Acids and Bases 2007-2008

  5. The formulas for oxoacids, (acids that contain hydrogen and an anion containing oxygen) are usually written with the H first, followed by the anion, as illustrated in the following examples: H2CO3 HNO3 HClO2 Carbonic acid Nitric acid Chlorous acid If the anion ends in “-ate” then the acid ends in “-ic”, if the anion ends in “-ite”, then the acid ends in “-ous”. Remember, “ic” goes with the higher oxidation state, N has an oxidation state of ___ in HNO3 (nitric acid) and ___ in HNO2 (nitrous acid) +5 +3 Acids and Bases 2007-2008

  6. Acids • Acids have a sour taste; for example, vinegar owes its sourness to acetic acid, and lemons and citrus fruits contain citric acid. • Acids cause color changes in plant dyes; for example, they change the color of blue litmus to red. • Acids react with certain metals to produce hydrogen gas. • Acids react with carbonates and bicarbonates to produce carbon dioxide gas. • Aqueous acid solutions conduct electricity. Acids and Bases 2007-2008

  7. Brønsted Acid • Arrhenius’s definitions of acids are limited in that they apply only to aqueous solutions. Broader definitions were proposed by the Danish chemist Johannes Brønsted. A Brønsted acid is a proton donor. HCl(aq) H+(aq) + Cl- Remember, the H+ ion is really just a proton (a hydrogen atom is one proton and one electron, you pull off the electron and all you are left with is… Acids and Bases 2007-2008

  8. The size of a proton is about 10-15 m, compared to the diameter of 10-10 m for an average atom or ion. Such an exceedingly small charged particle cannot exist as a separate entity in aqueous solution owing to its strong attraction for the negative region of the polar water molecule. Consequently, the proton exists in a hydrated form as H3O+, and is referred to as the hydronium ion H+ + H2O H3O+ Acids and Bases 2007-2008

  9. Since the acidic properties of the proton are unaffected by hydration, we will generally use H+(aq) to represent the hydrated proton. This notation is for convenience only, because H3O+ is closer to reality. Keep in mind that both notations represent the same species in aqueous solution. H+(aq) = H3O+(aq) Acids and Bases 2007-2008

  10. HCl Monoprotic acids each unit of acid yields one hydrogen upon ionization Diprotic acids each unit of an acid gives up two H+ ions Triprotic acids yields three H+ ions upon ionization H2CO3 H3PO4 Acids and Bases 2007-2008

  11. Diprotic acids give up their two H+ ions in separate steps: H2SO4(aq) H+(aq) + HSO4-(aq) HSO4-(aq) H+(aq) + SO4-2(aq) Triprotic acids give up their H+ ions in three separate steps. Is HSO4- a strong or weak acid? Explain Weak, only partially ionizes Acids and Bases 2007-2008

  12. Bases • In another definition formulated by Svante Arrhenius, a base can be described as a substance that yields hydroxide ions (OH-) when dissolved in water. Some examples are NaOH KOH Ba(OH)2 Sodium hydroxide Potassium hydroxide Barium hydroxide Acids and Bases 2007-2008

  13. Ammonia (NH3) is also classified as a common base. At first glance this may seem to be an exception to the definition of a base. Note that as long as a substance yields hydroxide ions when dissolved in water, it need not contain hydroxide ions in its structure to be considered a base. In fact, when ammonia dissolves in water, the following reaction occurs: NH3 + H2O NH4+ + OH- Thus it is properly classified as a base. Acids and Bases 2007-2008

  14. Bases • Bases have a bitter taste • Bases feel slippery; for example, soaps, which contain bases, exhibit this property • Bases cause color changes in plant dyes; for example, they change the color of red litmus to blue • Aqueous base solutions conduct electricity Acids and Bases 2007-2008

  15. Brønsted Base • A Brønsted base is defined by Johannes Brønsted as being a substance capable of accepting a proton. Acids and Bases 2007-2008

  16. Lewis Acids and Bases G.N. Lewis formulated a definition for what is now called a Lewis base – a substance that can donate a pair or electrons. A Lewis acid is a substance that can accept a pair of electrons. Acids and Bases 2007-2008

  17. The significance of the Lewis concept is that it is much more general than other definitions. For example, the reaction between boron trifluoride and ammonia is a Lewis acid-base reaction. Acids and Bases 2007-2008

  18. Strength of Acids and Bases Strong acids are strong electrolytes, which, for practical purposes, are assumed to ionize completely in water. That means that at equilibrium, solutions of strong acids will not contain any nonionized acid molecules. Like strong acids, strong bases are all strong electrolytes that ionize completely in water. Acids and Bases 2007-2008

  19. The most common strong acids are HClO4, HCl, HNO3 and H2SO4. Hydroxides of alkali metals and alkaline Earth metals are strong bases (like NaOH, KOH and Ba(OH)2). Other strong acids and strong bases are listed on your Relative Strengths of Acids and Bases Reference Sheet. Acids and Bases 2007-2008

  20. The strength of an acid is measured by its tendency to ionize: HX  H+ + X- The strength of the H-X bond influences the extent to which an acid undergoes ionization. The stronger the bond (the higher the bond dissociation energy in kJ/mol), the more difficult it is for the HX molecule to break up and hence the weaker the acid. Acids and Bases 2007-2008

  21. Bond Dissociation Energies for Hydrogen Halides and Acid Strengths weak strong strong strong Acids and Bases 2007-2008

  22. The Strength of Oxoacids Oxoacids contain hydrogen, oxygen, and one other element Z, which occupies a central position. To compare oxoacid strength, it is convenient to separate the oxoacids into two groups. Acids and Bases 2007-2008

  23. Oxoacids having different central atoms that are from the same group of the periodic table and that have the same oxidation number. Within this group, acid strength increases with increasing electronegativity of the central atom. HClO3 > HBrO3 Acids and Bases 2007-2008

  24. Oxoacids having the same central atom but different numbers of attached groups. Within this group, acid strength increases as the oxidation number of the central atom increases. HClO4 > HClO3 > HClO2 > HClO Which of the following should be a stronger acid, sulfurous or sulfuric? H2SO4 > H2SO3 Acids and Bases 2007-2008

  25. You are going to remember these trends because… • HCl is a strong acid and HF (with the higher bond dissociation energy) isn’t. • HClO4 is a strong acid and HClO3 (where the Cl has a lower oxidation number because it has fewer oxygen atoms attached to it) isn’t. Acids and Bases 2007-2008

  26. Note: H3O+ is the strongest acid that can exist in aqueous solutions. Acids stronger than H3O+ react with water to produce H3O+ and their conjugate bases. Thus, HCl, which is a stronger acid than H3O+, reacts with water completely to form H3O+ and Cl-. HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq) Acids and Bases 2007-2008

  27. The OH- ion is the strongest base that can exist in aqueous solution. Bases stronger than OH- react with water to produce OH- and their conjugate acids. For example, the oxide ion, (O-2) is a stronger base than OH-, so it reacts with water completely as follows: O-2(aq) + H2O(l)  2OH-(aq) For this reason, the oxide ion does not exist in aqueous solutions. Acids and Bases 2007-2008

  28. Amphoteric Compounds As you could see from the previous two examples, water will act as either an acid or a base, depending on the strength of the acid or base with which it is reacting. Any species that can react as either an acid or a base is described as amphoteric. NH3(g) + H2O(l)  OH-(aq) + NH4+(aq) H2SO4(aq) + H2O(l)  H3O+(aq) + HSO4-(aq) Proton acceptor (base) Proton donor (acid) Acids and Bases 2007-2008

  29. Acid Proton (H+) donor Conjugate acid Results from the addition of a proton to a Bronsted base Conjugate base Remains when one proton has been removedfrom the acid Base Proton (H+) acceptor An extension of the Brønsted definition of acids and bases is the concept of the conjugate acid-base pair CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq) A conjugate acid-base pair is defined as an acid and its conjugate base (what’s left after the H+ was removed from the acid) or a base and its conjugate acid (substance formed by the addition of the H+ to the base). ** Because the acid and base are always stronger than the conjugate acid and conjugate base, the direction of the reaction proceeds from acid/base  conjugate acid/conjugate base.

  30. Conjugate Acid Accepted proton (H+) (weaker acid) Base Proton (H+) acceptor (stronger base) Conjugate Base What’s left after H+ was donated by acid (weaker base) Acid Proton donor (Stronger acid) Identify the acid, base, conjugate acid and conjugate base in the following reaction (**Reaction proceeds from stronger to weaker…) NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) Acids and Bases 2007-2008

  31. The Acid-Base Properties of Water Water is a very weak electrolyte and therefore a poor conductor of electricity, but it does undergo ionization to a small extent: H2O(l) H+(aq) + OH-(aq) This reaction is sometimes called the autoionization of water. Acids and Bases 2007-2008

  32. Remember, pure liquids and solids are not listed in the ionization equation, therefore… [H+][OH-] [H2O] In the study of acid-base reactions in aqueous solutions, the hydrogen ion concentration is the key, because it indicates the acidity or alkalinity of the solution. Expressing the hydrogen ion as H+, we can write the equilibrium constant for the autoionization of water as kw = Acids and Bases 2007-2008

  33. kw = [H+][OH-] • kw is called the ion-product constant, and is the product of the molar concentrations of H+ and OH- ions at a particular temperature. Acids and Bases 2007-2008

  34. In pure water at 25 oC, the concentrations of H+ and OH- ions are equal and found to be [H+] = 1.0 x 10-7 M and [OH-] = 1.0 x 10-7 M. Thus, kw = [H+][OH-] kw = (1.0 x 10-7)(1.0 x 10-7) kw = 1.0 x 10-14 Acids and Bases 2007-2008

  35. Whether we have pure water or an aqueous solution of dissolved species, the following relation ALWAYS holds at 25 oC kw = [H+][OH-] = 1.0 x 10-14 Acids and Bases 2007-2008

  36. Because HCl is a strong acid… HCl  H+ + Cl- Calculate the concentration of OH- ions in an HCl solution whose hydrogen ion concentration is 1.3 M.

  37. Because the concentrations of H+ and OH- ions in aqueous solutions are frequently very small numbers and therefore inconvenient to work with, Soren Sorensen in 1909 proposed a more practical measure called pH. The pH of a solution is defined as the negative logarithm of the hydrogen ion concentration (in mol/L) pH = -log [H+] pH – A Measure of Acidity pH is a dimensionless quantity (it will not have a label) Acids and Bases 2007-2008

  38. Since pH is simply a way to express hydrogen ion concentration, acidic and basic solutions at 25 oC can be distinguished by their pH values, as follows: Acidic solutions: [H+] > 1.0 x 10-7 M, pH < 7.00 In an acidic solution there is an excess of H+ ions; [H+] > [OH-] Basic solutions: [H+] < 1.0 x 10-7 M, pH > 7.00 In a basic solution there is an excess of OH- ions; [OH-] > [H+] Neutral solutions [H+] = 1.0 x 10-7 M, pH = 7.00 Whenever [H+] = [OH-], the aqueous solution is said to be neutral. **Note – when concentration has two significant figures, pH will have two numbers TO THE RIGHT OF THE DECIMAL! Acids and Bases 2007-2008

  39. Calculate the pH of a 1.0 x 10-3 M HCl solution. Acids and Bases 2007-2008

  40. The concentration of H+ ions in a bottle of table wine was 3.2 x 10-4 M right after the cork was removed. Only half of the wine was consumed. The other half, after it had been standing open to the air for a month, was found to have a hydrogen ion concentration equal to 1.0 x 10-3 M. Calculate the pH of the wine on these two occasions. When the wine was first opened After the wine sat for a month Why did the acidity increase? Some of the ethanol converted to acetic acid, a reaction that takes place in the presence of O2. Acids and Bases 2007-2008

  41. Given the pH of a solution, you can figure out the [H+] concentration by using the simple formula [H+] = 10-pH What is the hydrogen ion concentration of an acid with a pH of 3.00? Acids and Bases 2007-2008

  42. A pH meter is commonly used in the laboratory to determine the pH of a solution. Although many pH meters have scales marked with values from 1 to 14, pH values can, in fact, be less than 1 and greater than 14. Acids and Bases 2007-2008

  43. A pOH scale analogous to the pH scale can be devised using the negative logarithm of the hydroxide ion concentration of a solution. Thus we define pOH as pOH = -log[OH-] Acids and Bases 2007-2008

  44. pH + pOH = 14.00 Now consider again the ion-product constant for water: kw = [H+][OH-] = 1.0 x 10-14 Taking the negative logarithm of both sides we obtain -log[H+] + -log[OH-] = -log(1.0 x 10-14) “Logs make adders multiply” -log[H+] + -log[OH-] = 14.00 pH pOH Acids and Bases 2007-2008

  45. In a NaOH solution [OH-] is 2.9 x 10-4 M. Calculate the pH of the solution. First, figure out the pOH… Then use the pOH to figure out the pH… Acids and Bases 2007-2008

  46. Calculate the pH of a 0.0020 M Ba(OH)2 solution. Acids and Bases 2007-2008

  47. To determine the hydroxide ion when given the pOH, you need to use the formula 10-pOH = [OH-] What is the molarity of a NaOH solution that has a pH of 11.30? Acids and Bases 2007-2008

  48. Weak Acids and Acid Ionization Constants Most acids are weak acids, which ionize only to a limited extent in water. At equilibrium, aqueous solutions of weak acids contain a mixture of nonionized acid molecules, H3O+ ions, and the conjugate base. The limited ionization of weak acids is related to the equilibrium constant for ionization, which is represented as ka. Acids and Bases 2007-2008

  49. Consider a weak monoprotic acid, HA. Its ionization in water is represented by HA(aq) + H2O(l) H3O+(aq) + A-(aq) or simply HA(aq) H+(aq) + A-(aq) Acids and Bases 2007-2008

  50. [H+][A-] [HA] ka= Write the equilibrium expression for the ionization of HA. ka, the acid ionization constant, is the equilibrium constant for the ionization of an acid. Acids and Bases 2007-2008