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Principles of Reactivity: Chemistry of Acids and Bases

Principles of Reactivity: Chemistry of Acids and Bases. Chapter 17. 17.1 Acids, Bases, and the Equilibrium Concept. Acids are substances that increase the concentration of hydrogen ions in solution. Bases are substances that increases the concentration of hydroxide ions in solution.

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Principles of Reactivity: Chemistry of Acids and Bases

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  1. Principles of Reactivity: Chemistry of Acids and Bases Chapter 17

  2. 17.1 Acids, Bases, and the Equilibrium Concept • Acids are substances that increase the concentration of hydrogen ions in solution. • Bases are substances that increases the concentration of hydroxide ions in solution. • Acids and bases can be either strong or weak electrolytes. • The proton produced by acids combines with water to form hydronium ion, H3O+.

  3. Electrolytes • A strong acid/base completely ionizes in solution. (NaOH, HCl) • Weak electrolytes ionize only to a small extent. (a large amount of acid remains in solution – HC2H3O2) • Equilibrium constants for weak acids are called ionization constants.

  4. Electrolytes • Acids or bases that ionize extensively, with K>1, are referred to as strong acids or bases. • Acids or bases that do not ionize extensively, with K<1, are referred to as weak acids or bases.

  5. 17.2 Bronsted Acids and Bases • Bronsted and Lowry said that an acid is a proton donor, such as nitric acid or ammonium. • A Bronsted base is a proton acceptor such as ammonia. • Monoprotic acids donate one proton (HF, HCl, HNO3) while polyprotic acids are capable of donating two or more (H2SO4).

  6. Bronsted Acids and Bases • Polyprotic bases may accept more than one proton, such as sulfate or phosphate ion. • Molecules that can act as either a Bronsted acid or base are called amphiprotic, such as water!

  7. Practice Problem • Write a balanced equation for the reaction that occurs when H3PO4, phosphoric acid, donates a proton to water to form the dihydrogen phosphate ion. Is the dihydrogen phosphate ion an acid, a base, or amphiprotic? • Write a balanced equation for the reaction that occurs when the cyanide ion, CN-, accepts a proton from water to form HCN. Is CN- a Bronsted acid or base?

  8. Conjugate Acid-Base Pairs • A pair of compounds or ions that differ by the presence of one H+ ion is called a conjugate acid-base pair. • Every reaction between a Bronsted acid and Bronsted base involves H+ transfer and has two conjugate acid-base pairs.

  9. Practice Problem • In the following reaction, identify the acid on the left and its conjugate base on with right. Similarly, identify the base on the left and its conjugate acid on the right. HNO3 + NH3 NH4+ + NO3-

  10. 17.3 Water and the pH Scale • Two water molecules interact with each other to produce a hydronium ion and a hydroxide ion. • This property is called autoionization. Water will still conduct a small amount of electricity since it contains low concentration of hydronium and hydroxide.

  11. Water Ionization Constant • Kw = [H3O+][OH-] = 1.0 x 10-14 at 25 oC • In pure water the two ion concentrations are equal and the water is said to be neutral. • Adding acid increases the hydronium concentration, adding base increases the hydroxide. Both additions would disturb the equilibrium.

  12. Practice Problem • A solution of the strong acid HCl has [HCl] = 4.7 x 10-3 M. What are the concentrations of H3O+ and OH- in this solution? (remember that HCl is 100% ionized in water)

  13. pH Scale • pH = -log[H3O+] • pOH = -log[OH-] • pKw = 14.00 = pH + pOH

  14. Indicators • Approximate pH of solutions can be determined using an acid-base indicator. • Substances that change color in a known pH range are Bronsted acids or bases for which the acid and its conjugate base have different colors. • Modern pH meters are preferable!

  15. Practice Problem • What is the pH of a 0.0012 M NaOH solution? • The pH of a diet soda is 4.32 at 25oC. What are the hydronium and hydroxide ion concentrations in the soda? • If the pH of a solution of the strong base Sr(OH)2 is 10.46, what is the concentration of the Sr(OH)2 in mol/L?

  16. Homework • After reading sections 17.1-17.3, you should be able to do the following… • P. 801 (1-14)

  17. 17.4 Equilibrium Constants • The lower the pH the stronger the acid. • For a strong base/acid, the [OH-]/[H3O+] concentration is equal to the original base/acid concentration. • For a weak base/acid, the ion concentration is much less than the original acid concentration.

  18. Equilibrium Constants • For the general acid HA, we can write Ka = [H3O+][A-]/[HA] where the Ka is the equilibrium constant for an acid in water. • K is less than 1 for weak acids. Ka increases as acid strength increases.

  19. Equilibrium Constants • For a weak base B in water… Kb = [BH+][OH-]/[B] Ionization Constants for Some Acids and their Conjugate Bases are on page 770. The weaker the acid, the stronger its conjugate base.

  20. Practice Problem • Which is the stronger acid, H2SO4 or H2SO3? • Is benzoic acid, C6H5CO2H, stronger or weaker than acetic acid? • Which has the stronger conjugate base, acetic acid or boric acid? • Which is the stronger base, ammonia or the acetate ion? • Which has the stronger conjugate acid, ammonia or the acetate ion?

  21. Aqueous Solutions of Salts • Anions that are conjugate bases of strong acids are such weak bases that they have no effect on solution pH. • There are numerous basic anions; all are the conjugate bases of weak acids. • Acidic anions arise from polyprotic acids, and are amphiprotic as well. • Alkali metal and alkaline earth cations have no measurable effect on solution pH. • All metal cations are hydrated in water. Only when the cation is a +2 or +3 does the ion act as an acid.

  22. Practice Problem • For each of the following salts in water, predict whether the pH will be greater than, less than, or equal to 7. • KBr • NH4NO3 • AlCl3 • Na2HPO4

  23. pKa • The negative log of the Ka value (pKa) becomes smaller as the acid strength increases. pKa = -logKa

  24. Practice Problem • What is the pKa value for benzoic acid, C6H5CO2H? • Is chloroacetic acid (ClCH2CO2H), pKa = 2.87, a stronger or weaker acid than benzoic acid? • What is the pKa for the conjugate acid of ammonia? Is this acid stronger or weaker than acetic acid?

  25. Relating Ionization Constants • As Ka decreases, Kb increases. The product of the two is equal to the autoionization constant for water. • Kw = KaKb

  26. Practice Problem • Ka for lactic acid, CH3CHOHCO2H, is 1.4 x 10-4. What is Kb for the conjugate base of this acid, CH3CHOHCO2-? Where does this base fit in Table 17.3?

  27. 17.5 Acid-Base Reactions • All proton transfer reactions proceed from the stronger acid and base to the weaker acid and base. (Equilibrium favors the weaker acid and base.)

  28. Practice Problem • Which is the stronger Bronsted acid, HCO3- or NH4+? Which has the stronger conjugate base? • Is a reaction between HCO3- ions and NH3 product- or reactant-favored? HCO3-(aq) + NH3(aq)  CO32-(aq) + NH4+(aq)

  29. Practice Problem • Write the net ionic equation for the possible reaction between acetic acid and sodium hydrogen sulfate, NaHSO4. Does the equilibrium lie to the left or right?

  30. Homework • After reading sections 17.4-17.5, you should be able to do the following… • P. 801 (15-17,23-25,30-31,34-35)

  31. 17.6 Types of Acid-Base Rxns • Strong Acid with a Strong Base • mixing equal quantities will produce a neutral solution • Weak Acid with a Strong Base • mixing equal quantities produces a basic solution; pH depends on Kb for anion produced

  32. Types of Acid-Base Rxns • Strong Acid with a Weak Base • mixing equal quantities produces an acidic solution; pH depends on Ka for the cation produced • Weak Acid with a Weak Base • mixing equal quantities produces a solution in which the pH depends on the Ka and Kb for the cations and anions produced

  33. Practice Problem • Equal molar quantities of HCl and NaCN are mixed. Is the resulting solution acidic, basic, or neutral? • Equal molar quantities of acetic acid and sodium sulfite, Na2SO3, are mixed. Is the resulting solution acidic, basic, or neutral?

  34. 17.7 Calculations • You can use an ICE table to calculate K from initial concentrations and measured pH.

  35. Practice Problem • A solution prepared from 0.055 mol of butanoic acid dissolved in sufficient water to give 1.0L of solution has a pH of 2.72. Determine Ka for butanoic acid. The acid ionizes according to the balanced equation CH3CH2CH2CO2H + H2O  H3O+ + CH3CH2CH2CO2-

  36. Equilibrium Constants • Due to the fact that very little ionization occurs in a weak acid, we can assume that the acid concentration at equilibrium is basically the same as the initial acid concentration. • This is valid whenever [HA]0 is greater than or equal to 100Ka.

  37. pH of Weak Acid or Base • We can use ICE tables to calculate the pH of a solution of a weak acid or base and using known equilibrium constants.

  38. Practice Problem • What are the equilibrium concentrations of acetic acid, acetate ion, and H3O+ for a 0.10M solution of acetic acid (Ka = 1.8x10-5)? What is the pH of the solution?

  39. Practice Problem • What are the equilibrium concentrations of HF, fluoride ion, and H3O+ when a 0.015M solution of HF is allowed to come to equilibrium? What is the pH of the solution?

  40. Practice Problem • Sodium hypochlorite, NaOCl, is used as a disinfectant in swimming pools and water treatment plants. What are the concentrations of HOCl and OH- and the pH of a 0.015M solution of NaOCl?

  41. pH After Acid-Base Reaction • In order to calculate pH in a resulting solution, you must write a balanced equation and decide whether the products are acid or basic. • You must then find initial concentrations and can calculate pH by solving an equilibrium problem.

  42. Practice Problem • Calculate the pH after mixing 15mL of 0.12M acetic acid with 15mL of 0.12M NaOH. What are the major species in solution at equilibrium (besides water) and what are their concentrations?

  43. Homework • After reading 17.6 and 17.7, you should be able to do the following… • P. 802 (38-39,42-45,50-55,60-62)

  44. 17.8 Polyprotic Acids and Bases • Acids that can donate more than one proton are polyprotic. • Each step has its own Ka, which becomes progressively smaller due to the increased energy required to remove a proton. • The pH of many polyprotic acids depends primarily on the hydronium ion generated in the first ionization step.

  45. Practice Problem • What is the pH of a 0.10M solution of oxalic acid, H2C2O4? What are the concentrations of H3O+, HC2O4-, and the oxalate ion, C2O42-?

  46. 17.9 Lewis Acids and Bases • This concept is based on the sharing of electron pairs between acid and base. • A Lewis acid is a substance that can accept a pair of electrons from another atom to form a new bond. • A Lewis base is a substance that can donate a pair of electrons to another atom to form a new bond.

  47. Lewis Acids and Bases • The product of an acid-base reaction in the Lewis sense is often called an acid-base adduct. This type of bond is called a coordinate covalent bond. • The formation of hydronium and ammonium are both examples.

  48. Cationic Lewis Acids • All metal cations form hydrated cations in which the metal ion is surrounded by water molecules, such a [Fe(H2O)6]2+. • These structures are called complex ions, or coordination complexes. • Hydroxide ion is a Lewis base and binds readily to metal cations to form metal hydroxides. Metal hydroxides are usually amphoteric.

  49. Molecular Lewis Acids • Acidic oxides such as carbon dioxide and sulfur dioxide. • Due to oxygen’s high electronegativity, electrons are polarized away from the other element which can therefore react with a Lewis base such as hydroxide ion.

  50. Practice Problem • Describe each of the following as a Lewis acid or a Lewis base. (Draw the Lewis dot structure. Are there lone pairs on the central atom? If so, it may be a Lewis base. Does the central atom lack an electron pair? If so it can behave as a Lewis acid. • PH3, BCl3, H2S, HS-

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