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Acids & Bases

Acids & Bases. Models of Acids & Bases. Arrhenius Acids produce H + in aqueous solutions Bases produce OH - in aqueous solutions Limited approach Brønsted -Lowry Acids are proton donors (H + ) Bases are proton acceptors H 3 O + - Hydronium ion. Conjugate Pairs.

jillian-adriano
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Acids & Bases

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  1. Acids & Bases

  2. Models of Acids & Bases • Arrhenius • Acids produce H+ in aqueous solutions • Bases produce OH- in aqueous solutions • Limited approach • Brønsted-Lowry • Acids are proton donors (H+) • Bases are proton acceptors • H3O+ - Hydronium ion

  3. Conjugate Pairs • Conjugate base – what remains after an acid has donated a proton • Cl- is the conjugate base of HCl • Conjugate acid – what is formed when a base accepts a proton • H3O+ is the conjugate acid of H2O • H2O + HCl H3O+ + Cl- • HCl is a stronger acid than H3O+so equilibrium lies far to the right • General reaction : H2O + HA  H3O+ + A-

  4. Acid Dissociation Constant (Ka) • General reaction : H2O + HA  H3O+ + A- • Ka = [H3O+][A-] = [H+][A-] [HA] [HA] • Water is not included • In dilute solutions the [H2O] is high and changes very little • Ka is the equilibrium constant for the reaction in which a proton is removed from the HA to form the conjugate base A-

  5. Remember! A strong acid has a large Ka value

  6. Practice • Write the dissociation reaction for the following acids • Acetic acid • HC2H3O2 • Hydrofluoric acid • HF

  7. Acid Strength

  8. Strong Acids • Acids for which the equilibrium lies far to the right • Almost completely dissociated • Yield weak conjugate bases • Common strong acids • Strong six (or seven) • HCl, H2SO4, HBr, HI, HClO4, HNO3, (HClO3)

  9. Weak Acids • Equilibrium lies far to the left • Mostly together, few ions • Yield relatively strong conjugate bases • Examples – acetic (HC2H3O2) oxalic (H2C2O4)

  10. Terminology • Monoprotic – one acidic proton • Diprotic – two acidic protons • Triprotic – three acidic protons • Oxyacids – acids in which the acidic proton is attached to an oxygen atom • Organic acid – contains the mildly acidic carboxyl group • Generally weak acids • Equilibrium lies far to the left

  11. Water is Amphoteric • Water can act as both an acid and a base • Autoionization of water • H2O + H2O  H3O+ + OH- • Ion-product constant Kw • At 25°C [H+] = [OH-] = 1.0 x 10-7 • Kw = [H+] [OH-] • Kw = (1.0 x 10-7) (1.0 x 10-7) = (1.0 x 10-14)

  12. Solution Characteristics • Neutral solution • [H+] = [OH-] • Acidic solution • [H+] > [OH-] • Basic solution • [H+] < [OH-]

  13. pH Scale Basic – pH > 7 Neutral – pH = 7 Acidic – pH < 7

  14. pH and pOH • pH = -log[H+] • The number of decimal places in the log is equal to the number of significant figures in the original number • [H+] = 1.0 x 10-7 • pH = 7.00 • Keep the two decimal places • pOH = -log[OH-] • pH + pOH = 14

  15. Calculating the pH of Strong Acid Solutions • Determine what species are present • Determine which components are significant • Those present in large amounts • 1.0 M solution HCl • Solution components – H+, Cl-, H2O (notHCl) • Major species – H+, Cl-, H2O • OH- is present in only small amounts - ignored

  16. Calculate the pH of 1.0 M HCl solution • Consider the H+ from HCl • H+ from H2O is negligible • The H+ from HCl will drive the autoionization back according to LeChatelier’s Principle • pH =-log [H+] • pH = -log(1.0) • pH = 0

  17. Calculating the pH of Weak Acid Solns • List the major species in the solution • Choose the species that can produce H+ and write the balanced equations for the reactions producing H+ • Using the values of the equilibrium constants for the reactions you have written decide which equilibrium will dominate in producing H+ • Write the equilibrium expression for the dominant equilibrium • List the initial concentrations of the species participating in the dominant equilibrium • Define the change needed to achieve equilibrium; that is define x • Write the equilibrium concentrations in terms of x • Substitute the equilibrium concentrations into the equilibrium expression • Solve for x the “easy” way (by assuming [HA]0 – x ≈ [HA]0) • Use the 5% rule to verify whether the approximation is valid • Calculate [H+] and pH

  18. Example • Calculate the pH of 1.0 M solution of HF. Ka= 7.2 x 10-4

  19. Mixture of Weak Acids • In a mixture of weak acids, if one acid has a relatively higher Ka value, it will be the focus of the solution in order to calculate pH • Same process as before

  20. Example • Calculate the pH of a solution that contains 1.00 M HCN (Ka = 6.2 x 10-10) and 5.00 M HNO2 (Ka = 4.0 x 10-4). Also calculate the concentration of CN- in this solution at equilibrium.

  21. Percent Dissociation • Percent dissociation = amount dissociated x 100% initial concentration • Percent dissociation = percent ionization • Specifies how much of the weak acid has dissociated • For a given weak acid, the percent dissociation increases as the acid becomes more dilute • More dilute (water) = more dissociation • Percent dissociation of acetic acid is significantly greater in 0.10 M solution than in a 1.00 M solution

  22. Example • Calculate the percent dissociation of acetic acid (Ka = 1.8 x 10-5) in each of the following • 1.00 M Solution • 0.10 M Solution

  23. Bases Strong Bases Weak Bases • Group 1A metal hydroxides • LiOH, NaOH, KOH, RbOH, CsOH • Group 2A metal hydroxides • Ca(OH)2, Ba(OH)2, Sr(OH)2 • Less soluble than 1A hydroxides • Use an antacids • Ammonia (NH4+) and other covalent bases • Bases don’t have to contain OH- • React with water to increase [OH-] • Proton acceptors • Compounds with low values of Kb

  24. Bases • Reaction with waterB + H2O  BH+ + OH- • Equilibrium constant = Kb • Kb = [BH+][OH-] [B] • Kb is the equilibrium constant for the reaction of a base to form conjugate acid and OH-

  25. Calculating the pH of Strong Bases • Kw = [H+][OH-] = 1.0 x 10-14 • If [OH-] is known, [H+] can be calculated • From [H+], can calculate pH • Calculate the pH of 0.05 MNaOH

  26. Calculation of pOH • pKw = 14 = pH + pOH • pOH = 14-pH • Same process as a weak acid Calculating pH of Weak Bases

  27. Polyprotic Acids

  28. Polyprotic Acids • Dissociates in a stepwise manner • One proton at a time • Most are weak acids • H2CO3 H+ + HCO3- Ka1 = 4.3 x 10-7 • HCO3-  H+ + CO32- Ka2 = 5.6 x 10-11 • Conjugate base becomes the acid in the 2nd step • Ka1 > Ka2 > Ka3

  29. Example – Phosphoric Acid • Calculate the pH of 5.0 M H3PO4solution, and equilibrium concentrations of H3PO4, H2PO4-, HPO42-, PO43- • Ka1= 7.5 x 10 -3 • Ka2 = 6.2 x 10-8 • Ka3 = 4.8 x 10-13 • Only the first dissociation makes a significant contribution of H+ based on Ka values

  30. Sulfuric Acid • Strong acid in its first dissociation • Weak acid in its second • Second step cannot be ignored for dilute solutions • Calculate the pH of a 1.0 M H2SO4 solution.

  31. Acid-Base Properties of Salts

  32. Salts that Produce Neutral Solutions • Salts that consist of the cations of strong bases and the anions of strong acids have no effect on pH, ([H+]), when dissolved in water • Cations of strong bases • Na+, K+(Group 1A) • No affinity for protons • Can’t produce H+ • Anions of strong acids • Cl-, NO3- • Have no affinity for protons • Solutions of NaCl, KCl, NaNO3, and KNO3 are neutral

  33. Salts that Produce Basic Solutions • For any salt whose cation has neutral properties and whose anion is the conjugate base of a weak acid, the aqueous solution will be basic

  34. Salts that Produce Basic Solutions C2H3O2-(aq) + H2O(l)  HC2H3O2(aq) + OH-(aq) • Acetic acid is a weak acid – acetate is the conjugate base • Significant affinity for protons • Will react with the best proton donor available • Water is the only source of protons • Produces hydroxide ions – basic solution

  35. Salts as Weak Bases • For any weak acid and its conjugate base • Kax Kb = Kw • Calculate the pH of a 0.30 M NaF solution. • Ka = 7.2 x 10-4 • Solve for Kb • ICE Chart • Solve for x

  36. Salts that Produce Acidic Solutions • Salts in which the anion is not a base and the cation is the conjugate acid of a weak base produce acidic solutions • NH4Cl • Anion (Cl-) has no affinity for a proton • Cation (NH4+) behaves as a weak acid

  37. Salts that Produce Acidic Solutions • Calculate the pH of a 0.10 M NH4Cl solution. • Kb value for NH3 = 1.8 x 10-5 • NH4+ and H2O can produce H+ • Solve for Ka (NH4+) • ICE Chart • Solve for x • Find pH

  38. Salts that Produce Acidic Solutions • Salt that contains a highly charged metal ion • AlCl3 • Hydrated Al ion [Al(H2O)63+] is a weak acid • High metallic charge polarized the O-H bond in water • Hydrogen in water becomes acidic • Typically – the higher the charge on the metal ion, the stronger the acidity of the hydrated ion

  39. Salts that Produce Acidic Solutions • Calculate the pH of 0.010 M AlCl3 solution • Ka = 1.4 x 10-5 Al(H2O)63+ • Solve like a typical weak acid

  40. Solutions of Salts • Many salts - both ions can affect the pH of the solution • Predict whether the solution will be acidic, basic, or neutral by comparing Ka value of acidic ion and Kb of basic ion • Ka > Kb – solution is acidic • Kb > Ka – solution is basic • Kb= Ka – solution is neutral

  41. Example • Determine whether an aqueous solution will be acidic, basic, or neutral • NH4C2H3O2 • Ions in solution are NH4+ and C2H3O2- • Ka = 5.6 x 10-10 NH4+ • Kb = 5.6 x 10-10 C2H3O2- • Ka = Kb, solution is neutral

  42. Effect of Structure on Acid-Base Properties • Factors determining acid characteristics of molecules with X-H bond • Strength of bonds • Strong bonds are reluctant to break • Polarity of bonds • High bond polarity tends to increase the acidity of hydrogen • Halogens – electronegativity decreases down a group • HF has an unusually strong bond – weak acid

  43. Effect of Structure • Molecules of form H-O-X • If X has high electronegativity, the hydrogen tends to be acidic • If X has low electronegativity, the compound tends to be basic • -OH comes off • The more oxygen atoms around X, the more acidic the compound • Oxygen pulls the electrons

  44. Acid-Base Properties of Oxides • Acidic Oxides (acid anhydrides) • Nonmetal oxides that react with water to form acidic solutions • SO3(g) + H2O(l)  H2SO4(aq) • 2 NO2(g) + H2O(l)  HNO3(aq) + HNO2(aq) • Covalent bonds • Basic Oxides (basic anhydrides) • Metallic oxides of Group 1A and 2A metals react with water to form basic solutions • K2O(s) +H2O(l)  2 KOH(aq) • CaO(s) + H2O(l)  Ca(OH)2(aq) • Ionic bonds

  45. Lewis Acid-Base Model • Lewis acid – electron pair acceptor • H+ can accept an electron pair • NH3 + H+ NH4+ • Lewis base – electron pair donor • OH- can donate an electron pair • OH- + H+ H2O

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