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Acid-Base Equilibria

Acid-Base Equilibria. REVIEW. Electrolyte : Substances that dissolves in water to produce solutions that conduct electricity Nonelectrolytes : Substances whose aqueous solutions do not conduct electricity Strong and weak relates to the degree of dissociation or ionization

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Acid-Base Equilibria

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  1. Acid-Base Equilibria

  2. REVIEW • Electrolyte: Substances that dissolves in water • to produce solutions that conduct electricity • Nonelectrolytes: Substances whose aqueous • solutions do not conduct electricity • Strongandweakrelates to the degree of • dissociation or ionization • In aqueous solutions protons occasionally exist • as hydrated molecules, H(H2O)n+ (n=1)

  3. Definitions: Arrhenius: An acid is a substance that increases the H+ (or H3O+) concentration in an aqueous solution. HCl + H 2O  H3O+ + Cl- HCl  H+ + Cl- A base is a substance that increases the OH- concentration in an aqueous solution. NaOH(s)  Na+ + OH- What about Na2CO3 ????

  4. Bronsted-Lowry: HCl(aq) + NaOH(aq) → HOH + NaCl Acid = a proton donor in a Reaction Base = a proton acceptor in a Reaction

  5. Lewis: An acid is an electron pair acceptor H+ acid .. H:O:H .. A base is an electron pair donor water .. :O:H- ..

  6. Acid/Base reactions: Produce water and a salt (and sometimes carbon dioxide). Hint: concentrate on the water first. Remember, water has the formula HOH. Complete and balance the following: HCl + KOH  HOH + KCl HCl + Ca(OH)2 2 2HOH + CaCl2 Require equal numbers

  7. 1. Ba(OH)2 + H3PO4 2. HC2H3O2 + NaOH 3. H2SO4 + KOH 4. H2CO3 + NaOH 5. Na2CO3 + HCl 

  8. 6. NH4OH + H2SO4 7. NH3 + HCl  Give a definition of an acid: An acid is a proton donor (H+) Give a definition of a base: A base is a proton acceptor

  9. Conjugate acids and Conjugate bases HCl + KOH  HOH + KCl acid base conj. base conj. acid Na2CO3 + 2HCl  H2CO3 + 2NaCl acid conj. base base conj. acid Na2CO3 + 2HCl  H2O + CO2(g) + 2NaCl acid conj. acid base conj. base

  10. NH3 + HCl  NH4+ + Cl-

  11. What is a strong Acid? An Acid that is 100% ionized in water. Strong Acids: 100% ionized (completely dissociated) in water. HCl + H2O  H3O+ + Cl- often written as: HCl  H+ + Cl-

  12. Strong Acids: 100% ionized (completely dissociated) in water. HCl + H2O  H3O+ + Cl- Strong Acids: Perchloric HClO4 Chloric, HClO3 Hydrobromic, HBr Hydrochloric, HCl Hydroiodic, HI Nitric, HNO3 Sulfuric, H2SO4

  13. What is a strong Base? A base that is completely dissociated in water (highly soluble). NaOH(s)  Na+ + OH- Strong Bases: Group 1A metal hydroxides (LiOH, NaOH, KOH, RbOH, CsOH) Heavy Group 2A metal hydroxides [Ca(OH)2, Sr(OH)2, and Ba(OH)2]

  14. Weak Acids: “The Rest”

  15. Strong Acids: 100% ionized (completely dissociated) in water. HCl + H2O  H3O+ + Cl- Note the “one way arrow”. Weak Acids: Only a small % (dissociated) in water. HC2H3O2 + H2O H3O+ + C2H3O2- Note the “2-way” arrow. Why are they different?

  16. Strong Acids: HCl HCl HCl HCl HCl (H2O) ADD WATER to MOLECULAR ACID

  17. Strong Acids: Cl- H3O+ (H2O) Cl- H3O+ H3O+ Cl- Cl- H3O+ H3O+ Cl- Note: No HCl molecules remain in solution, all have been ionized in water.

  18. Weak Acid Ionization: HC2H3O2 HC2H3O2 (H2O) HC2H3O2 HC2H3O2  HC2H3O2 Add water to MOLECULES of WEAK Acid

  19. Weak Acid Ionization: HC2H3O2 HC2H3O2 H30+ C2H3O2- HC2H3O2 (H2O) HC2H3O2  H30+ C2H3O2- HC2H3O2 HC2H3O2 Note: At any given time only a small portion of the acid molecules are ionized and since reactions are running in BOTH directions the mixture composition stays the same. This gives rise to an Equilbrium expression, Ka

  20. ACID-BASE CONCEPTS • Acid-base reactions may be one of the most • important class of reactions • The most basic of the acid-base concepts is • the Arrhenius theory • Acids are substances that dissociate in water • to produce hydronium ions, H3O+ and bases are • substances that dissociate in water to produce • hydroxide ions, OH-

  21. ACID-BASE STRENGTH • A strongacid is almost completely dissociated • or ionized HCl(aq)  H+(aq) + Cl-(aq) • A weak acid is only partially dissociated or • ionized • So each reaction is an equilibrium with a K

  22. STRONG ACIDS AND BASES • Common strong acids are either monoprotic or • diprotic • HA(aq) + H2O(l)  H3O+(aq) + A-(aq) 100% • Means [HA] = [H3O+] • Similar situation for strong bases • Typical strong bases: Group 1A metal and Ca, • Sr, and Ba hydroxides • Strong acids: HX, HNO3, HClO3, HClO4, H2SO4

  23. WEAK ACIDS • Weak acids and bases are weak electrolytes • For a weak acid, HA • Ka: thedissociationorionizationconstant • For acetic acid

  24. Ka = [H3O+][CH3COO-]/[CH3COOH] • Acids with larger ionization constants ionize or • dissociate to a greater extent than acids with • smaller ionization constants • The larger the value of Ka, the higher [H3O+] • and the stronger is the acid • HIO3; Ka = 1.6 x 10-1 • CH3COOH; Ka = 1.8 x 10-5 • HCN; Ka = 6.2 x 10-10

  25. AUTOIONIZATION OF WATER • Reaction called autoionization of water • K = [H3O+][OH-]/[H2O]2 • K[H2O]2 = [H3O+][OH-] • Kw = [H3O+][OH-] • Kw is called ion product of water • At 25 °C, Kw = 1.0 x 10-14 • Valid also for dilute aqueous solutions

  26. THE pH SCALE • The hydronium ion concentration is a measure • of a solution’s acidity • Usually small numbers • The pH scale is used express acidity and • basicity • pH = -log[H3O+]so[H3O+] = 10-pH • Note that as pH increases [H3O+] decreases • Value of Kw varies with temperature

  27. “Acidity” depends on the concentration of • H3O+ ions • Acidic: [H3O+] > [OH-] • Basic: [H3O+] < [OH-] • Neutral: [H3O+] = [OH-] • Notice that neutral does NOT necessarily • mean pH 7 • pH is usually quoted with the same number of • significant digits as the concentration

  28. pH can be measured using a pH meter or by an • acid-base indicator • An indicator is usually an organic acid that have • different colors in solutions of different pH • Indicators exist that cover the entire pH scale • e.g. bromothymol blue: 6.0 – 7.6 • phenolphthalein: 8.0 – 10

  29. The p-scale can also be applied to ionization • constants • pKa = - log Ka • Thelargerthe value of Ka thesmallerthe value of pKa and thestrongerthe acid • From the acid-dissociation constant we can • calculate equilibrium concentrations as well as • pH

  30. BRØNSTED-LOWRY THEORY • An acid is a proton donor • A base is a proton acceptor • An acid-base reaction is the transfer of a • proton from an acid to a base • Reactions can be described in terms of what • are called conjugate acid-base pairs • - species which differ by a proton • - a charge

  31. - B is a proton acceptor; it is a base - BH+ is a proton donor; it is an acid So A- is the conjugate base of HA and BH+ is the conjugate acid of B • HF/F- and H3O+/H2O are conjugate pairs

  32. Relationship of Ka and concentration • In dilute solutions of weak acids, the assumption • - is that all the H3O+ is coming from the acid • - the concentration change of a species is small • compared to the initial concentration of that • species • - at equilibrium, [HA]  [HA]init • Always be cautious when using the assumptions • given above

  33. %-dissociation • Amount of acid that dissociates can be • expressed as a percent • In general %-dissociation increases with the • value of Ka • For any weak acid, HA, %-dissociation • increases with dilution

  34. POLYPROTIC ACIDS • Polyprotic acids:Acids that provide more than • one hydronium ion in solution • e.g. H2SO4, H3PO4 • Polyprotic acids ionize in a stepwise manner • Consider sulfuric acid, H2SO4 (1) (2) • In general, Ka1 > Ka2

  35. WEAK BASES • Weak bases undergo equilibria in water • For a general weak base, B • Kb is called the base-dissociation constant

  36. For conjugate acid-base pairs the product of • their equilibrium constants is the ionization • constant for water • Ka x Kb = Kw

  37. FACTORS AFFECTING ACID STRENGTH • The strength of an acid depends on the polarity • of the H-E bond • The polarity of the bond is related to the bond • strength of H-E • The weaker the H-E bond the stronger the • acid • Take the hydrohalic acids: • HX; X = F, Cl, Br, I

  38. HF << HCl < HBr < HI • For binary acid in the same group H-A bond • strength determines acid strength • Same applies to other groups: • H2O < H2S < H2Se • For binary acids in the same row polarity of the • H-E bond determines acid strength • :- acid strength increase with the • electronegativity of E

  39. The oxoacids are also important • e.g. HNO3, H2SO4, HClO4 • The acidity is dictated by factors that affect • the O-H bond strength • The electronegativity of the central element • The oxidation number of the central element

  40. Some small, highly charged metal ions are quite • acidic • :- they are hydrated and transfers a proton • to a water molecule

  41. 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

  42. 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

  43. 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

  44. 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

  45. 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+

  46. 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)

  47. 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

  48. 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-

  49. If the parent Ka < Kb: solutions basic • e.g. NH4CN • Ka(HCN): 4.9 x 10-10 < Kb (NH3): 1.8 x 10-5 • Kb(CN-): 2.0 x 10-10 > Ka(NH4+): 5.6 x 10-10 • So NH4+ hydrolyzes to a greater extent than F- • more OH- is produced than H3O+

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