Acids and Bases: Properties, Definitions, and Reactions

1. Lecture Presentation Chapter 15Acids and Bases Sherril Soman Grand Valley State University

2. Properties of Acids • Sour taste • Ability to dissolve many metals • Ability to neutralize bases • Change blue litmus paper to red

3. Common Acids

4. Structures of Acids • Binary acids have acid hydrogens attached to a nonmetal atom. • HCl, HF

5. Structure of Acids • Oxyacids have acid hydrogens attached to an oxygen atom. • H2SO4, HNO3

6. Structure of Acids • Carboxylic acids have COOH group. • HC2H3O2, H3C6H5O7 • Only the first H in the formula is acidic. • The H is on the COOH.

7. Properties of Bases • Taste bitter • alkaloids = plant product that is alkaline • often poisonous • Feel slippery • Ability to turn red litmus paper blue • Ability to neutralize acids

8. Common Bases

9. Indicators • Indicators are chemicals that change color depending on the solution’s acidity or basicity. • Many vegetable dyes are indicators. • Anthocyanins, red cabbage • Litmus • From Spanish moss • Red in acid, blue in base • Phenolphthalein • Found in laxatives • Red in base, colorless in acid

10. Definitions of Acids and Bases • Arrhenius definition Based on H+ and OH- • Acid = H+ increase in water • Base = OH- increase in water • Brønsted–Lowry definition Based on reactions in which H+ is transferred • Acid = substance that donates H+ • Base = substance that accepts H+ • Lewis definition Based on un-shared electron pairs • Acid = an electron-pair acceptor. • Base is any substance that donate a pair of electrons

11. Arrhenius Theory • Acids: produce H+ ions in aqueous solution. HCl(aq) → H+(aq) + Cl−(aq)

12. Hydronium Ion • The H+ ions produced by the acid are so reactive they cannot exist in water. • H+ ions are protons! • Instead, they react with water molecules to produce complex ions, mainly hydronium ion, H3O+. H+ + H2O  H3O+ • There are also minor amounts of H+ with multiple water molecules, H(H2O)n+.

13. Arrhenius Theory • Bases: produce OH− ions in aqueous solution. NaOH(aq) → Na+(aq) + OH(aq)

14. Arrhenius Acid–Base Reactions • The H+ from the acid combines with the OH− from the base to make a molecule of H2O. • It is often helpful to think of H2O as H—OH. • The cation from the base combines with the anion from the acid to make a salt. acid + base → salt + water HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

15. Problems with Arrhenius Theory • It does not explain why molecular substances, such as NH3, dissolve in water to form basic solutions, even though they do not contain OH– ions. • It does not explain how some ionic compounds, such as Na2CO3 or Na2O, dissolve in water to form basic solutions, even though they do not contain OH– ions. • It does not explain why molecular substances, such as CO2, dissolve in water to form acidic solutions, even though they do not contain H+ ions. • It does not explain acid–base reactions that take place outside aqueous solution.

16. Brønsted–Lowry Acid–Base Theory • It defines acids and bases based on what happens in a reaction. • Any reaction involving H+ (proton) that transfers from one molecule to another is an acid–base reaction, regardless of whether it occurs in aqueous solution or if there is OH− present. • All reactions that fit the Arrhenius definition also fit the Brønsted–Lowry definition.

17. Brønsted–Lowry Theory • The acid is an H+ donor. (proton donor) • The base is an H+ acceptor. (proton acceptor) • Base structure must contain an atom with an unshared pair of electrons. • In a Brønsted–Lowry acid–base reaction, the acid molecule donates an H+ to the base molecule. H–A + :B :A– + H–B+

18. Brønsted–Lowry Acids • Brønsted–Lowry acids are H+ donors. • Any material that has H can potentially be a Brønsted–Lowry acid. • Because of the molecular structure, often one H in the molecule is easier to transfer than others. • When HCl dissolves in water, the HCl is the acid because HCl transfers an H+ to H2O, forming H3O+ ions. • Water acts as base, accepting H+. HCl(aq) + H2O(l) → Cl–(aq) + H3O+(aq) acid base

19. Brønsted–Lowry Bases • Brønsted–Lowry bases are H+ acceptors. • Any material that has atoms with lone pairs can potentially be a Brønsted–Lowry base. • Because of the molecular structure, often one atom in the molecule is more willing to accept H+ transfer than others. • When NH3 dissolves in water, the NH3(aq) is the base because NH3 accepts an H+ from H2O, forming OH–(aq). • Water acts as acid, donating H+. NH3(aq) + H2O(l)  NH4+(aq) + OH–(aq) base acid

20. Amphoteric Substances • Amphoteric substancescan act as either an acid or a base because they have both a transferable H and an atom with lone pair electrons. • Water acts as base, accepting H+ from HCl. HCl(aq) + H2O(l) → Cl–(aq) + H3O+(aq) • Water acts as acid, donating H+ to NH3. NH3(aq) + H2O(l)  NH4+(aq) + OH–(aq)

21. Brønsted–Lowry Acid–Base Reactions • One of the advantages of Brønsted–Lowry theory is that it illustrates reversiblereactions to be as follows: H–A + :B:A– + H–B+ • The original base has an extra H+ after the reaction, so it will act as an acid in the reverse process. • And the original acid has a lone pair of electrons after the reaction, so it will act as a base in the reverse process: :A– + H–B+H–A + :B

22. Conjugate Acid–Base Pairs • In a Brønsted-–Lowry acid–base reaction, • the original base becomes an acid in the reverse reaction. • the original acid becomes a base in the reverse process. • Each reactant and the product it becomes is called a conjugate pair.

23. Conjugate Pairs A base accepts a proton and becomes a conjugate acid. An acid donates a proton and becomes a conjugate base.

24. Arrow Conventions • Chemists commonly use two kinds of arrows in reactions to indicate the degree of completion of the reactions. • A single arrow indicates that all the reactant molecules are converted to product molecules at the end. • A double arrow indicates the reaction stops when only some of the reactant molecules have been converted into products. •  in these notes

25. Strong or Weak • A strong acid is a strong electrolyte. • Practically all the acid molecules ionize. • A strong base is a strong electrolyte. • Practically all the base molecules form OH– ions, either through dissociation or reaction with water. • A weak acid is a weak electrolyte. • Only a small percentage of the molecules ionize, . • A weak base is a weak electrolyte • only a small percentage of the base molecules form OH– ions, either through dissociation or reaction with water, .

26. Strong Acids • Strong acids donate practically all their H’s. • 100% ionized in water • Strong electrolyte • [H3O+] = [strong acid] • [X] means the molarity of X

27. Weak Acids • Weak acids donate a small fraction of their H’s. • Most of the weak acid molecules do not donate H to water. • Much less than 1% ionized in water • [H3O+] << [weak acid]

28. Examples of Strong Acids • Chloric Acid (HClO3) • Periodic (HIO4)

29. Examples of Weak Acids

30. Strengths of Acids and Bases • Commonly, acid or base strength is measured by determining the equilibrium constant of a substance’s reaction with water. HAcid + H2O  Acid− + H3O+ Base: + H2O  HBase+ + OH− • The farther the equilibrium position lies toward the products, the stronger the acid or base. • The position of equilibrium depends on the strength of attraction between the base form and the H+. • Stronger attraction means stronger base or weaker acid.

31. Ionic Attraction and Acid Strength

32. Autoionization of Water • Water is amphoteric; it can act either as an acid or a base. • Therefore, there must be a few ions present. • About 2 out of every 1 billion water molecules form ions through a process called autoionization. H2O Û H+ + OH– H2O + H2O Û H3O+ + OH– • All aqueous solutions contain both H3O+ and OH–. • The concentration of H3O+ and OH– are equal in water. • [H3O+] = [OH–] = 10−7M at 25 °C

33. Ion Product of Water • The product of the H3O+ and OH– concentrations is always the same number. • The number is called the ion productof water and has the symbol Kw. • Also know as the dissociation constant of water • [H3O+] × [OH–] = Kw= 1.00 × 10−14 at 25 °C • If you measure one of the concentrations, you can calculate the other. • As [H3O+] increases the [OH–] must decrease so the product stays constant. • Inversely proportional

34. Acid Ionization Constant, Ka • Acid strength is measured by the size of the equilibrium constant when it reacts with H2O. • The equilibrium constant for this reaction is called the acid ionization constant, Ka. • larger Ka = stronger acid

35. Table 15.5

36. Acidic and Basic Solutions • All aqueous solutions contain both H3O+ and OH– ions. • Neutral solutions have equal [H3O+] and [OH–]. • [H3O+] = [OH–] = 1.00 × 10−7 • Acidic solutions have a larger [H3O+] than [OH–]. • [H3O+] > 1.00 × 10−7; [OH–] < 1.00 × 10−7 • Basic solutions have a larger [OH–] than [H3O+]. • [H3O+] < 1.00 × 10−7; [OH–] > 1.00 × 10−7

37. Measuring Acidity: pH • The acidity or basicity of a solution is often expressed as pH. • pH = −log[H3O+] • Exponent on 10 with a positive sign • pHwater = −log[10−7] = 7 • Need to know the [H3O+] concentration to find pH • pH < 7 is acidic; pH > 7 isbasic. pH = 7 is neutral. • [H3O+] = 10−pH

38. Sig., Figs., and Logs • When you take the log of a number written in scientific notation, the digits before the decimal point come from the exponent on 10, and the digits after the decimal point come from the decimal part of the number. log(2.0 x 106) = log(106) + log(2.0) = 6 + 0.30303… = 6.30303... • Because the part of the scientific notation number that determines the significant figures is the decimal part, the sig. figs. are the digits after the decimal point in the log. log(2.0 × 106) = 6.30

39. What Does the pH Number Imply? • The lower the pH, the more acidic the solution; the higher the pH, the more basic the solution. • 1 pH unit corresponds to a factor of 10 difference in acidity. • Normal range of pH is 0 to 14. • pH 0 is [H3O+] = 1 M; pH 14 is [OH–] = 1 M. • pH can be negative (very acidic) or larger than 14 (very alkaline).

40. pOH • Another way of expressing the acidity/basicity of a solution is pOH. • pOH = −log[OH], [OH] = 10−pOH • pOHwater = −log[10−7] = 7 • Need to know the [OH] concentration to find pOH • pOH < 7 is basic; pOH > 7 is acidic; pOH = 7 is neutral. • pH + pOH = 14.0

41. Relationship between pH and pOH • pH + pOH = 14.00at 25 °C. • You can use pOH to find the pH of a solution. pH= -log [H3O+] pOH= -log [OH-]

42. pK • A way of expressing the strength of an acid or base is pK. • pKa = −log(Ka), Ka = 10−pKa • pKb = −log(Kb), Kb = 10−pKb • The stronger the acid, the smaller the pKa. • Larger Ka = smaller pKa • Because it is the –log • The stronger the base, the smaller the pKb. • Larger Kb = smaller pKb

43. [H3O+] and [OH−] in a Strong Acid or Strong Base Solution • There are two sources of H3O+ in an aqueous solution of a strong acid—the acid and the water. • There are two sources of OH− in an aqueous solution of a strong acid—the base and the water. • For a strong acid or base, the contribution of the water to the total [H3O+] or [OH−] is negligible. • The [H3O+]acid shifts the Kw equilibrium so far that [H3O+]wateris too small to be significant. • Except in very dilute solutions, generally < 1 × 10−4 M (< 0.0001M)

44. Finding pH of a Strong Acid or Strong Base Solution • For a monoprotic strong acid [H3O+] = [HAcid]. • For polyprotic acids, the other ionizations can generally be ignored. • For H2SO4, the second ionization cannot be ignored. H2SO4  2H3O+ + SO4 • 0.10 M HCl has [H3O+] = 0.10 M and pH = 1.00 • For a strong ionic base, [OH−] = (number OH−) × [Base]. • For molecular bases with multiple lone pairs available, only one lone pair accepts an H; the other reactions can generally be ignored. • 0.10 M Ca(OH)2 has [OH−] = 0.20 M and pH = 13.30. Ca(OH)2  Ca2+ + 2 OH-

45. Finding the pH of a Weak Acid • There are also two sources of H3O+ in an aqueous solution of a weak acid—the acid and the water. • However, finding the [H3O+] is complicated by the fact that the acid only undergoes partial ionization. • Calculating the [H3O+] requires solving an equilibrium problem for the reaction that defines the acidity of the acid. HAcid + H2O  Acid + H3O+

46. Percent Ionization • Another way to measure the strength of an acid is to determine the percentage of acid molecules that ionize when dissolved in water; this is called the percent ionization. • The higher the percent ionization, the stronger the acid. • Because [ionized acid]equil = [H3O+]equil

47. Relationship between [H3O+]equilibrium and [HA]initial • Increasing the initial concentration of acid results in increased [H3O+] at equilibrium. • Increasing the initial concentration of acid results in decreased percent ionization. • This means that the increase in [H3O+] concentration is slower than the increase in acid concentration.

48. Why Doesn’t the Increase in H3O+ Keep Up with the Increase in HA? • The reaction for ionization of a weak acid is as follows: HA(aq) + H2O(l)  A−(aq) + H3O+(aq) • According to Le Châtelier’s principle, if we reduce the concentrations of all the (aq) components, the equilibrium should shift to the right to increase the total number of dissolved particles. • We can reduce the (aq) concentrations by using a more dilute initial acid concentration. • The result will be a larger [H3O+] in the dilute solution compared to the initial acid concentration. • This will result in a larger percent ionization.

49. Finding the pH of Mixtures of Acids • Generally, you can ignore the contribution of the weaker acid to the [H3O+]equil. • For a mixture of a strong acid with a weak acid, the complete ionization of the strong acid provides more than enough [H3O+] to shift the weak acid equilibrium to the left so far that the weak acid’s added [H3O+] is negligible. • For mixtures of weak acids, you generally only need to consider the stronger for the same reasons, as long as one is significantly stronger than the other and their concentrations are similar.

50. Strong Bases • The stronger the base, the more willing it is to accept H. • Use water as the standard acid. • For ionic bases, practically all units are dissociated into OH– or accept H’s. • Strong electrolyte • Multi-OH strong bases completely dissociated