1 / 45

Comprehensive Overview of Acids, Bases, and Their Models

This detailed guide delves into the fundamental concepts of acids and bases, including key models such as Arrhenius and Brønsted-Lowry. It covers the behavior of acids and bases in aqueous solutions, emphasizing their roles as proton donors and acceptors. Key topics include acid dissociation constants (Ka), strong and weak acids, and the unique properties of water as an amphoteric substance. The document also provides practical calculations for pH and percent dissociation, ensuring a thorough understanding of these essential chemistry concepts.

jillian-adriano
Télécharger la présentation

Comprehensive Overview of Acids, Bases, and Their Models

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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

More Related