1 / 67

Understanding Acids and Bases: Theories, Properties, and pH Calculations

This chapter delves into the fundamental concepts of acids and bases, exploring key theories such as the Arrhenius and Bronsted-Lowry theories. It highlights how acids produce H+ ions and bases produce OH- ions in solution, as well as the concept of conjugate acid-base pairs. The chapter further categorizes acids into strong and weak, explains the significance of the acid dissociation constant (Ka), and discusses the role of water as an amphoteric substance. Additionally, it introduces logarithms and pH calculations essential for determining the acidity or basicity of solutions.

tierra
Télécharger la présentation

Understanding Acids and Bases: Theories, Properties, and pH Calculations

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. Chapter 14 Acids and Bases

  2. Acid/Base Theories • Arrhenius Theory • Acids produce H+ ions in solution • Bases produce OH- ions in solution • Downside • Must be in solution and must have those ions • Bronsted-Lowry Theory • Acids are H+ donors (Proton donors) • Bases are H+ acceptors (Proton acceptors)

  3. Vocabulary • H+ is the hydrogen ion • Just a proton • H3O+ is a hydronium ion • It is the way H+ exists in water • Water accepts a hydrogen ion and becomes H3O+ • Either way is fine the first is just easier

  4. Conjugates Acid/Base Pairs • Conjugate Acid is formed when a base gains a proton • Conjugate Base is what remains after an acid donates a proton • Ex – HNO3 + H2O  H3O+ + NO3-

  5. General Form for Acids and Bases HA + B  A- + BH+ • HA is an acid • B is a base • A- is the conjugate base • Just the negative ion of the acid • BH+ is the conjugate acid • Just the base plus a hydrogen

  6. Strong and Weak Acids • Strong acids completely ionize in solution • Nitric, Perchloric, Sulfuric, Hydrochloric, Hydrobromic, Hydroiodic • Weak Conjugate bases • Weak Acids only partially ionize in solution • Every other acid • Equilibrium is established in the ionization • Weak acids have Ka values • Strong Conjugate bases

  7. Acid Dissociation Constant, Ka HA(aq) + H2O  A- + H3O+ • For weak acids equilibrium is established • Equilibruim constant is Ka • Ka=[A-][H3O+]/[HA] • Values tend to be small • Because CB is fairly strong • Strong acids do not have Ka values.

  8. Acid Terms • Monoprotic – One acidic hydrogen • Polyprotic – Many acidic hydrogens • Diprotic – Two acidic hydrogens • Triprotic – Three acidic hydrogens • Oxyacid – Acid that has an acidic hydrogen attached to an oxygen • Organic Acid – Acid that has the acidic hydrogen attached to the carboxyl group

  9. Water As An Acid and Base • Water is amphoteric – Both an acid and base H2O + H2O  H3O+ + OH- • Equilibrium system that always has the same value • Called Kw

  10. Autoionization of Water, Kw H2O + H2O  H3O+ + OH- • Kw = [H3O+] [OH-] • In pure water the products have the same concentration, 1.00x10-7M • Value of Kw = 1.00x10-14 at 25ºC • Concentrations can change is acid or base is added

  11. Acid, Base, or Neutral • If the concentration of H+ = OH- • Neutral • If the concentration of H+ > OH- • Acidic • If the concentration of H+ < OH- • Basic

  12. What is the hydrogen ion concentration when the hydroxide ion concentration is 1.00x10-5M

  13. What is the hydroxide ion concentration when the when the concentration of nitric acid is 0.0010M?

  14. Homework • P. 704 30, 32, 33, 35, 39ab,40ab

  15. Logarithms • The logarithm of a number to a given base (commonly 10) is the power or exponent to which the base must be raised in order to produce the number. • SAY WHAT!

  16. Examples • If your question says log 100 = x It is saying to what power must 10 be raised to equal 100 10x = 100 102 = 100 So x = 2

  17. Examples log 1 = x x = 0 log 10 = x x = 1 log 1000 = x x = 3 log 1x106 = x x = 6

  18. Examples • Logarithms can also be used for numbers smaller than 1 log 0.1 = x x = -1 log 0.01 = x x = -2 log 1x10-5 x = -5

  19. Examples • If they are not easy to calculate you can do it on your calculator log 15 = x You can approximate it between . . . • Type log 15 on your calculator x = 1.18

  20. Examples • If your question says log x = 7 It is saying 10 to the 7th power is what number 107 = 1x107 log x = 3 x = 1000 log x = -3 x = 0.001

  21. Tougher Examples log 234 = x x = 2.37 log x = -3.3 x = 5.0x10-4 -log 9.1x10-5 = x x = 4.0 -log x = 12.1 x = 7.9x10-13

  22. pH • Negative logarithm of the concentration of hydrogen ions in solution pH = -log [H+] • pH means power of Hydrogen • Measures how acidic or basic a solution is • pH scale typically goes from 0 to 14 • pH < 7 Acidic • pH > 7 Basic

  23. Highly acidic = low pH • Highly basic = high pH

  24. Significant Figures and pH • Digits after the decimal are the only ones that are significant in pH values pH = 4.44 2 Significant Figures pH = 10.874 3 Significant Figures • If your [H+] is 0.088 M your pH is 1.06

  25. Determining [H+] in solution • The concentration of a strong acid is equal to the H+ concentration. • 0.010 M HCl has an [H+] of 0.010 M • To obtain the [H+] for weak acids you must use equilibrium • Need Ka data • Discuss hydroxide later

  26. Other Info Turn Kw into a log equation pH + pOH = 14.00

  27. pH’s You MUST Know • When the [H+] is _______ the pH is _____ 0.10 M 1.00 0.0010M 3.00 1.0x10-6M 6.00 1.0x10-10M 10.00

  28. Find the pOH, [H+], [OH-] of lemon juice that has a pH of 2.48

  29. Determine the pH of 0.150M HCl

  30. Determine the pH of 2.3x10-3 M Hydrocyanic Acid HCN. Ka = 6.2x10-10

  31. Homework • Page 705 #’s 45,47,50,52,53,58

  32. Mixtures of Weak Acids • When there are mixtures of weak acids in solution determining the pH could be a difficult problem. • However • The acid with the largest Ka will control the pH of the solution

  33. Solutions of 0.10M HF (Ka = 7.2x10-4) and 0.10M HCN (Ka = 6.2x10-10) are mixed A) Which acid will control the pH of the solution? Why? B) What is the pH of the resulting solution.

  34. Percent Dissociation • Ratio of the concentration of the dissociated ions to the initial concentration • Found just like the doing the 5% check • Can be used to find pH and Ka

  35. A 0.25M solution of HClO is 20.% dissociated. A) What is the pH? B) What is the Ka value of the acid?

  36. Strong Bases • Strong bases are any compound containing the hydroxide ion • Group 1 hydroxides are very soluble • Group 2 less soluble but still strong • Group 2 have two hydroxides per mole • Be Careful

  37. Determine the pH of 0.022M Sr(OH)2

  38. Weak Bases • Organic Bases are weak bases. (Ammonia) • Contain Nitrogen (Amines) • CH3NH2 – Methylamine • Must have a lone pair of electrons • The hydroxide will come from water (CH3)3N + H2O

  39. Cont. • Weak bases have Kb values B + H2O BH+ + OH- • The conjugate acid of a weak base is stronger than water.

  40. Determine the pH of 15.0 M ammonia. The Kb is 1.8x10-5

  41. Homework • Page 705 #’s 62,63a,72,73,77ab,87

  42. Polyprotic Acids • Acids with more than one acidic hydrogen • Dissociate in a “stepwise” process H3PO4 H+ + H2PO4- Ka1 =7.5x10-3 H2PO4-  H+ + HPO4-2 Ka2 =6.2x10-8 HPO4-2  H+ + PO4-3 Ka3 =4.8x10-13 • Ka1 > Ka2 > Ka3 • Successive dissociation do not effect pH (except sulfuric acid)

  43. Determine the pH of 5.0M H3PO4 (Ka=7.5x10-3) and the [H2PO4-], [HPO4-2], [HPO4-3]

  44. Sulfuric Acid • Sulfuric acids has two dissociations • The first is strong • The second is weak • The second dissociation is quite strong, but it is not complete H2SO4 H+ + HSO4- (Strong) HSO4-  H+ + SO4-2 Ka = 1.2x10-2

  45. Continued • When you write sulfuric acid in net ionic equations only use the first equation H2SO4 H+ + HSO4- • The second dissociation only needs to be considered in dilute solutions. • A 1.0M solution of sulfuric acid will have a lower pH than a 1.0M solution of HCl

  46. Acid / Base Properties of Salts • Some salts have acid base properties • Make a solution have a pH below or above 7 • Some have no acid base properties • Make a solution with a pH of 7

More Related