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CHEMISTRY Chapter 14 Aqueous Equilibria: Acids and Bases

John E McMurry and Robert C Fay. CHEMISTRY Chapter 14 Aqueous Equilibria: Acids and Bases. Acid–Base Concepts 01. Arrhenius Acid: A substance which dissociates to form hydrogen ions (H + ) in solution. HA( aq )  H + ( aq ) + A – ( aq )

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CHEMISTRY Chapter 14 Aqueous Equilibria: Acids and Bases

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  1. John E McMurry and Robert C Fay CHEMISTRYChapter 14Aqueous Equilibria: Acids and Bases

  2. Acid–Base Concepts 01 • Arrhenius Acid:A substance which dissociates to form hydrogen ions (H+) in solution. HA(aq)  H+(aq) + A–(aq) • Arrhenius Base:A substance that dissociates in, or reacts with water to form hydroxide ions (OH–). • MOH(aq)  M+(aq) + OH–(aq)

  3. Acid–Base Concepts 02 • Brønsted–Lowry Acid:Substance that can donate H+ • Brønsted–Lowry Base:Substance that can accept H+ • Chemical species whose formulas differ only by one proton are said to be conjugate acid–base pairs.

  4. Acid–Base Concepts 03

  5. Acid–Base Concepts 04

  6. Acid–Base Concepts 05 • A Lewis Acid is an electron-pair acceptor. These are generally cations and neutral molecules with vacant valence orbitals, such as Al3+, Cu2+, H+, BF3. • A Lewis Base is an electron-pair donor. These are generally anions and neutral molecules with available pairs of electrons, such as H2O, NH3, O2–. • The bond formed is called a coordinate bond.

  7. Acid–Base Concepts 06

  8. Acid–Base Concepts 07 • Write balanced equations for the dissociation of each of the following Brønsted–Lowry acids. (a) H2SO4 (b) HSO4– (c) H3O+ • Identify the Lewis acid and Lewis base in each of the following reactions: • (a) SnCl4(s) + 2 Cl–(aq)  SnCl62–(aq) • (b) Hg2+(aq) + 4 CN–(aq)  Hg(CN)42–(aq) • (c) Co3+(aq) + 6 NH3(aq)  Co(NH3)63+(aq)

  9. Dissociation of Water 01 • Water can act as an acid or as a base. H2O(l)  H+(aq) + OH–(aq) • This is called the autoionization of water. H2O(l) + H2O(l)H3O+(aq) + OH–(aq)

  10. Dissociation of Water 02 • This equilibrium gives us the ion product constant for water. Kw = Kc = [H+][OH–] = 1.0 x 10–14 • If we know either [H+] or [OH–] then we can determine the other quantity.

  11. Dissociation of Water 03 • The concentration of OH– ions in a certain household ammonia cleaning solution is 0.0025 M. Calculate the concentration of H+ ions. • Calculate the concentration of OH– ions in a HCl solution whose hydrogen ion concentration is 1.3 M.

  12. pH – A Measure of Acidity 01 • The pH of a solution is the negative logarithm of the hydrogen ion concentration (in mol/L). pH = –log [H+], pH + pOH = 14 Acidic solutions: [H+] > 1.0 x 10–7 M, pH < 7.00Basic solutions: [H+] < 1.0 x 10–7 M, pH > 7.00Neutral solutions: [H+] = 1.0 x 10–7 M, pH = 7.00 pH 1pH 7pH 14 strong weak neutralweak strong acid acidbase base The pH scale according to the late Dr. Hubert Alyea, Princeton University

  13. 100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14 14 13 11 9 7 5 3 1 0 10-14 10-13 10-11 10-9 Basic 10-7 10-5 10-3 10-1 100 [H+] pH 0 1 3 5 7 9 11 13 14 Acidic Neutral Basic pOH [OH-]

  14. pH – A Measure of Acidity 02 • Nitric acid (HNO3) is used in the production of fertilizer, dyes, drugs, and explosives. Calculate the pH of a HNO3 solution having a hydrogen ion concentration of 0.76 M. • The pH of a certain orange juice is 3.33. Calculate the H+ ion concentration. • The OH– ion concentration of a blood sample is 2.5 x 10–7 M. What is the pH of the blood?

  15. pH – A Measure of Acidity 03

  16. Strength of Acids and Bases 01 • Strong acids and bases: are strong electrolytes that are assumed to ionize completely in water. • Weak acids and bases: are weak electrolytes that ionize only to a limited extent in water. • Solutions of weak acids and bases contain ionized and non-ionized species.

  17. Strength • Strong acids and bases are strong electrolytes • They fall apart (ionize) completely. • Weak acids don’t completely ionize. • Strength different from concentration • Strong-forms many ions when dissolved • Mg(OH)2 is a strong base- it falls completely apart when dissolved. • But, not much dissolves- not concentrated

  18. Measuring strength • Ionization is reversible. • HA H+ + A- • This makes an equilibrium • Acid dissociation constant = Ka • Ka = [H+ ][A- ] (water is constant) [HA] • Stronger acid = more products (ions), thus a larger Ka

  19. What about bases? • Strong bases dissociate completely. • B + H2O BH+ + OH- • Base dissociation constant = Kb • Kb = [BH+ ][OH-] [B] (we ignore the water) • Stronger base = more dissociated, thus a larger Kb.

  20. Strength vs. Concentration • The words concentrated and dilute tell how much of an acid or base is dissolved in solution - refers to the number of moles of acid or base in a given volume • The words strong and weak refer to the extent of ionization of an acid or base • Is concentrated weak acid possible?

  21. Strength of Acids and Bases 02 HClO4 HI HBr HCl H2SO4 HNO3 H3O+ HSO4– ACIDCONJ. BASE ACIDCONJ. BASE ClO4– I– Br – Cl – HSO4 – NO3 – H2O SO42– HSO4– HF HNO2 HCOOH NH4+ HCN H2O NH3 SO42– F – NO2 – HCOO – NH3 CN – OH – NH2 – IncreasingAcid Strength Increasing Acid Strength

  22. Strength of Acids and Bases 03 • Stronger acid + stronger base weaker acid + weaker base • Predict the direction of the following: • HNO2(aq) + CN–(aq)  HCN(aq) + NO2–(aq) • HF(aq) + NH3(aq)  F–(aq) + NH4+(aq)

  23. Acid Ionization Constants 01 • Acid Ionization Constant: the equilibrium constant for the ionization of an acid.HA(aq) + H2O(l)  H3O+(aq) + A–(aq) • Or simply: HA(aq)  H+(aq) + A–(aq)

  24. Acid Ionization Constants 02 ACIDKaCONJ. BASE Kb HF HNO2 C9H8O4 (aspirin) HCO2H (formic) C6H8O6 (ascorbic) C6H5CO2H (benzoic) CH3CO2H (acetic) HCN C6H5OH (phenol) 7.1 x 10 –4 4.5 x 10 –4 3.0 x 10 –4 1.7 x 10 –4 8.0 x 10 –5 6.5 x 10 –5 1.8 x 10 –5 4.9 x 10 –10 1.3 x 10 –10 F– NO2 – C9H7O4 – HCO2 – C6H7O6 – C6H5CO2 – CH3CO2 – CN – C6H5O – 1.4 x 10 –11 2.2 x 10 –11 3.3 x 10 –11 5.9 x 10 –11 1.3 x 10 –10 1.5 x 10 –10 5.6 x 10 –10 2.0 x 10 –5 7.7 x 10 –5

  25. Strength of Acids and Bases 03 (a) Arrange the three acids in order of increasing value of Ka. (b) Which acid, if any, is a strong acid? (c) Which solution has the highest pH, and which has the lowest?

  26. Acid Ionization Constants 04 • Initial Change Equilibrium Table: Determine the pH of 0.50M HA solution at 25°C. Ka = 7.1 x 10–4. - + H + A  HA (aq) (aq) (aq) Initial ( M ) : 0.50 0.00 0.00 Change (M): – x + x + x Equilib 0.50 – x x x (M):

  27. Acid Ionization Constants 05 • pH of a Weak Acid (Cont’d): • Substitute new values into equilibrium expression. • If Ka is significantly (>1000 x) smaller than [HA] the expression (0.50 – x) approximates to (0.50). • The equation can now be solved for x and pH. • If Ka is not significantly smaller than [HA] the quadratic equation must be used to solve for x and pH.

  28. Acid Ionization Constants 06 • The Quadratic Equation: • The expression must first be rearranged to: • The values are substituted into the quadratic and solved for a positive solution to x and pH.

  29. Acid Ionization Constants 07 • Calculate the pH of a 0.036 M nitrous acid (HNO2) solution. • What is the pH of a 0.122 M monoprotic acid whose Ka is 5.7 x 10–4? • The pH of a 0.060 M weak monoprotic acid is 3.44. Calculate the Ka of the acid.

  30. Acid Ionization Constants 08 • Percent Dissociation: A measure of the strength of an acid. • Stronger acids have higher percent dissociation. • Percent dissociation of a weak acid decreases as its concentration increases.

  31. Base Ionization Constants 01 • Base Ionization Constant: The equilibrium constant for the ionization of a base. • The ionization of weak bases is treated in the same way as the ionization of weak acids.B(aq) + H2O(l)  BH+(aq) + OH–(aq) • Calculations follow the same procedure as used for a weak acid but [OH–] is calculated, not [H+].

  32. Base Ionization Constants 02 BASEKbCONJ. ACID Ka C2H5NH2 (ethylamine) CH3NH2 (methylamine) C8H10N4O2 (caffeine) NH3 (ammonia) C5H5N(pyridine) C6H5NH2 (aniline) NH2CONH2 (urea) 5.6 x 10 –4 4.4 x 10 –4 4.1 x 10 –4 1.8 x 10 –5 1.7 x 10 –9 3.8 x 10 –10 1.5 x 10 –14 C2H5NH3+ CH3NH3+ C8H11N4O2+ NH4+ C5H6N+ C6H5NH3+ NH2CONH3+ 1.8 x 10 –11 2.3 x 10 –11 2.4 x 10 –11 5.6 x 10 –10 5.9 x 10 –6 2.6 x 10 –5 0.67 Note that the positive charge sits on the nitrogen.

  33. Diprotic & Polyprotic Acids 01 • Diprotic and polyprotic acids yield more than one hydrogen ion per molecule. • One proton is lost at a time. Conjugate base of first step is acid of second step. • Ionization constants decrease as protons are removed.

  34. Diprotic & Polyprotic Acids 02 ACIDKaCONJ. BASE Kb H2SO4 HSO4– C2H2O4 C2HO4– H2SO3 HSO3– H2CO3 HCO3– H2S HS– H3PO4 H2PO4– HPO42– Very Large 1.3 x 10 –2 6.5 x 10 –2 6.1 x 10 –5 1.3 x 10 –2 6.3 x 10 –8 4.2 x 10 –7 4.8 x 10 –11 9.5 x 10 –8 1 x 10 –19 7.5 x 10 –3 6.2 x 10 –8 4.8 x 10 –13 HSO4 – SO4 2– C2HO4– C2O42– HSO3 – SO3 2– HCO3– CO3 2– HS– S 2– H2PO4– HPO42– PO43– Very Small 7.7 x 10 –13 1.5 x 10 –13 1.6 x 10 –10 7.7 x 10 –13 1.6 x 10 –7 2.4 x 10 –8 2.1 x 10 –4 1.1 x 10 –7 1 x 10 –5 1.3 x 10 –12 1.6 x 10 –7 2.1 x 10 –2

  35. Molecular Structure and Acid Strength 01 • The strength of an acid depends on its tendency to ionize. • For general acids of the type H–X: • The stronger the bond, the weaker the acid. • The more polar the bond, the stronger the acid. • For the hydrohalic acids, bond strength plays the key role giving: HF < HCl < HBr < HI

  36. Molecular Structure and Acid Strength 02 • The electrostatic potential maps show all the hydrohalic acids are polar. The variation in polarity is less significant than the bond strength which decreases from 567 kJ/mol for HF to 299 kJ/mol for HI.

  37. Molecular Structure and Acid Strength 03 • For binary acids in the same group, H–A bond strength decreases with increasing size of A, so acidity increases. • For binary acids in the same row, H–A polarity increases with increasing electronegativity of A, so acidity increases.

  38. Molecular Structure and Acid Strength 04 • For oxoacids bond polarity is more important. If we consider the main element (Y):Y–O–H • If Y is an electronegative element, or in a high oxidation state, the Y–O bond will be more covalent and the O–H bond more polar and the acid stronger.

  39. Molecular Structure and Acid Strength 05 • For oxoacids with different central atoms that are from the same group of the periodic table and that have the same oxidation number, acid strength increases with increasing electronegativity.

  40. Molecular Structure and Acid Strength 06 • For oxoacids having the same central atom but different numbers of attached groups, acid strength increases with increasing central atom oxidation number. • As shown on the next slide, the number of oxygen atoms increases the positive charge on the chlorine which weakens the O–H bond and increases its polarity.

  41. Molecular Structure and Acid Strength 07 • Oxoacids of Chlorine:

  42. Molecular Structure and Acid Strength 08 • Predict the relative strengths of the following groups of oxoacids: • a) HClO, HBrO, and HIO. • b) HNO3 and HNO2. • c) H3PO3 and H3PO4.

  43. SALTS: ACID-BASE PROPERTIES • Salts: an ionic compound that is formed when an acid neutralizes a base • NaOH(aq) + HCl(aq)  NaCl(aq) + H2O • Aqueous solutions of salts can be neutral, acidic or basic • Strong acid + strong base  neutral solutions • Strong acid + weak base  acidic solutions • Strong base + weak acid  basic solutions

  44. When a salt dissolves in water, its constituent ions may react with water – reaction called hydrolysis • NEUTRAL SOLUTIONS • Salts of strong acids and strong bases e.g. NaCl • Because the ions do not hydrolyze

  45. Cl- is the conjugate base of HCl – it is a weak • base • The same argument is made for Na+ • Essentially [H3O+]/[OH-] ratio does not • change • {Group 1A metals, Ca2+, Sr2+, Ba2+} and {I-, Br-, Cl-, NO3-, ClO4-} does not hydrolyze

  46. Acidic Solutions • NH3(aq) + HCl(aq)  NH4Cl(aq) • Salts from weak bases and strong acids • In aqueous solution NH4+ undergo hydrolysis • - the chloride ion does not • The generation of H3O+ from the reaction • makes these solutions acidic

  47. Basic Solutions • Salts from strong bases and weak acids give basic solutions • This basic anions of weak acids hydrolyze to form hydroxide ions • NaCH3COO  CH3COO- + H3O+

  48. SALTS FROM WEAK ACID AND BASES • e.g. NH4CH3COO, ammonium acetate • Aqueous solution of the salts may be basic, • acidic or neutral • The pH depends on the relative Ka and Kb of • the parent acid and base • Consider the case when Ka = Kb • NH4CH3COO(aq)  NH4+(aq) + CH3COO-(aq)

  49. Both ions can undergo hydrolysis • Hydrolysis constant for the acetate ion (Kbh) is equal to the hydrolysis constant for the ammonium ion (Kah) • The same concentration of H3O+ as OH- is • produced  solution is NEUTRAL

  50. If the parent Ka > Kb: solutions acidic e.g. NH4F • Ka(HF): 7.2 x 10-4 > Kb (NH3): 1.8 x 10-5 • Kb(F-): 1.4 x 10-11 < Ka(NH4+): 5.6 x 10-10 • So NH4+ hydrolyzes to a greater extent than F- • more H3O+ is produced than OH-

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