1 / 70

HCl (aq) hydrochloric acid HBr (aq) hydrobromic acid HI (aq) hydroiodic acid

Binary acids – use the prefix “hydro” + the name of the negative (second ion), changing the ending to “-ic acid”. Oxy-acids – Name the polyatomic ion, change the ending to either “-ic acid” or “-ous acid” according to the “ate = ic, ite = ous” rule. HCl (aq) hydrochloric acid

benito
Download Presentation

HCl (aq) hydrochloric acid HBr (aq) hydrobromic acid HI (aq) hydroiodic acid

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. Binary acids – use the prefix “hydro” + the name of the negative (second ion), changing the ending to “-ic acid”.Oxy-acids – Name the polyatomic ion, change the ending to either “-ic acid” or “-ous acid” according to the “ate = ic, ite = ous” rule. • HCl (aq) hydrochloric acid • HBr (aq) hydrobromic acid • HI (aq) hydroiodic acid • HF (aq) hydrofluoric acid • H2S (aq) hydrosulfuric acid • H2SO4 sulfuric acid • H2SO3 sulfurous acid • H3PO4 phosphoric acid • H3PO3 phosphorous acid • HNO3 nitric acid

  2. Acids and Bases

  3. What is an Acid? Base? • What do you think and acid is? • What do you think a base is?

  4. Properties of Acids and Bases! • Acids have specific properties • Sour taste • pH beow 7.00 S.U. • Turn litmus paper red • Dissolve metals • Bases have similar properties • Bitter taste • pH above 7.00 S.U. • Feel slippery • Turn litmus paper blue

  5. What else can they do? • Acids and bases neutralize each other • In other words, they raise/lower the pH to a non-acid/non-base • Acids and Bases can both destroy human tissue

  6. How do they do that? • Acids and bases are electrolytes • The ions that they produce give them their properties

  7. Strength of Acids • A strong acid is one that dissolves completely in solution. • A weak acid is one that does not dissolve completely. • Why would an acid not dissolve completely?

  8. Do you know any acids or bases? • List some common acids and bases • Vinegar is Acetic Acid • Ammonia is a Base • Sodium Bicarbonate (Baking soda) is a Base • Batteries contain Sulfuric Acid • Your stomach contains Hydrochloric Acid!

  9. There are more than one • There are 3 different types of acids and bases • Arrhenius • Brönsted-Lowry • Lewis

  10. Arrhenius Acids and Bases • An Arrhenius Acid is a substance that releases H3O+ ions in an aqueous solution • HCl H2O H3O+ + Cl- • An Arrhenius Base is a substance that releases –OH in an aqueous solution • NaOH H2O Na+ + OH-

  11. Brönsted-Lowry Acids and Bases • A Brönsted-Lowry acid is any species that can donate a proton (H+ ion) to another species; a proton donor • A Brönsted-Lowry base is any speices that can accept a proton (H+ ion) from another species; a proton acceptor

  12. HCl + H2O  H3O+ + Cl- • Which is the Brönsted-Lowry acid/base? • Acid is the proton donator • Base is the proton acceptor • HCl is the acid • H2O is the base

  13. Water: Acid or Base? • Take a look at these two reactions 1. NH3 + H2O  NH4++ OH- 2. HCl + H2O  H3O++ Cl- • In reaction #1, which is the B-L acid? Base? • Water is the acid and ammonia is the base • In reaction #2? • HCl is the acid and water is the base

  14. Amphiprotic • A compound that can be both a proton donor OR a proton acceptor in separate reactions is called amphiprotic or amphoteric

  15. Lewis Acids and Bases • A Lewis Acid is any species that can accept a pair of electrons from another species; electron pair acceptor • A Lewis Base is any species that can donate a pair of electrons from another species; electron pair donor

  16. Remember Lewis Structures! • Let’s look at a few examples H+ H+ is an electron pair acceptor :Ö-H- OH- is an electron pair donor :Ö-H Lewis acid-base product .. H ..

  17. Compare the Definitions

  18. Acid-Base Equilibria • In a chemical reaction involving both acids and bases, dynamic equilibrium can still be achieved • A species that is an acid in the forward direction, can become the base in the reverse direction

  19. HNO2 + H2O  H3O+ + NO2- acid base acid base • Let’s take a look at this reaction • HNO2 donates a proton in the forward reaction and is therefore a B-L acid • H2O accepts a proton in the forward reaction and is the B-L base • In the reverse direction, H3O+ donates a proton; B-L acid • And NO2- accepts a proton; B-L base

  20. Conjugate Pairs • When an acid donates a proton, it becomes a conjugate base • When a base accepts a proton, it becomes a conjugate acid • Why? • The acid-base reaction can be a reversible reaction and so when you move in the opposite direction, the compounds are redefined as acids and/or bases

  21. Conjugates • For example, look at the following reaction HCl + H2O  H3O+ + Cl- • Which reactant is the acid? The base? • Now, in the reverse direction, which product is the acid? The base? • Now, what are the conjugate pairs?

  22. Conjugates, some more • Look at this reaction: NH3 + H2O  NH4+ + OH- • Which reactant is the acid? The base? • Now, in the reverse direction, which product is the acid? The base? • Now, what are the conjugate pairs?

  23. Conjugate Practice • What are the conjugate pairs for the following reactions? • HNO3 + OH-  H2O + NO3- • HNO3 is the acid and NO3- is the conjugate base • OH- is the base and H2O is the conjugate acid • H2SO4(aq) + SO32-(aq)  HSO4-(aq) + HSO3-(aq) • Sulfuric acid is the acid and HSO4- is the conjugate base • Sulfite ion is the base and HSO3- is the conjugate acid

  24. Strength • If a acid is strong, then the conjugate base is weak and vice-versa

  25. Strong/Weak Electrolytes • An electrolyteis a solution that will conduct an electrical current. • Any ionic compound that dissociates in solution will conduct electricity. • If a substance dissociates 100%, it is called a strongelectrolyte. • A weak electrolyte dissociates < 100% (usually < 5%)

  26. Strong Acids • Binary acids: • HCl, HBr, and HI are strong acids. All other binary acids are weak. • Oxy acids: • If the number of oxygen atoms exceed the number of hydrogen atoms by 2 or more, the acid is strong.

  27. Strong Bases • Hydroxides made from metals in group 1 or group 2 are strong bases. • All other hydroxides are weak bases.

  28. Molecular Structure and the Strength of Acids • The strength of an acid depends on a number of factors, such as the properties of the solvent, the temperature, and the molecular structure of the acid. • We compare the strengths of two acids, in the same solvent and at the same temperature. That way we can focus on the structure of the acid.

  29. Binary Hydrides • HA ---> H+ + A- • Two factors influence the extent to which the acid undergoes ionization. • One is the strength of the H-A bond • the stronger the bond, the more difficult it is for the HA molecule to break up and hence the weaker the acid.

  30. Polarity versus Strength • The other factor is the polarity of the H-A bond. • The difference in electronegativities between H and A results in a polar bond. • If the bond is highly polarized, it is a stronger acid.

  31. Hydrohalic Acids • The halogens form a series of binary acids called the hydrohalic acids. • The strengths of the hydrohalic acids increase in the following order: • HF<<HCl<HBr<HI

  32. Other Non-metallic Acid Strengths • In any vertical column (Group) of nonmetallic elements, there is a tendency toward increasing acidity of the hydride with increasing atomic number (as you go down the group). • For example, among the group VIA elements the acid strength increases in the order: • H2O< H2S<H2Se<H2Te

  33. Oxyacids • Many common acids contain one or more O-H bonds.

  34. Practice Problems • HF • weak • H2S • weak • H2SO4 • strong • H2CO3 • weak • Fe(OH)3 • weak 6. Barium hydroxide strong 7. Chloric acid strong 8. Sulfurous acid weak 9. Hypochlorous acid weak 10. Tin(IV) hydroxide weak Ba(OH)2 HClO3 H2SO3 HClO Sn(OH)4

  35. Review • In the following equations, name the acid and base. What was your justification? • NH3 + H2O  NH4+ + OH- • 2NaBr + H2SO4  Br2 + SO2 + 2NaOH

  36. Review • What is the difference between a strong acid and a weak acid? • A buffer is a solution made from a weak acid and it’s conjugate base that neutralizes a small amount of acids or bases added to it

  37. Calculating pH

  38. pH • pH is a measure of how acidic, neutral or basic a solution is • The scale for pH is 0-14 • 7 is neutral • 0-7 is acidic • 7-14 is basic

  39. The Numbers mean . . . • pH + pOH = 14 S.U. • pH = -log10[H3O+] • pOH = -log10[OH-]

  40. What is pH? • The term pH refers to a scale that describes how strongly acidic or basic a solution is • pH + pOH = 14 • Acidic = 0-7 • Neutral = 7 • Basic = 7-14

  41. The Kw Constant • Kw = 1.0 x 10-14 • at standard thermodynamic temperature (25oC or 298 K)

  42. Ionization Constants of Acids and Bases • HA(aq) + H2O(l) ↔ H3O+(aq) + A-(aq) • Keq= [H3O+][A-] / [HA][H2O] • In a dilute solution, the concentration of water is constant, so • Keq= [H3O+][A-] / [HA]

  43. Ionization Constants • The Ionization Constant for water; Kw is derived as follows: • First: for the acids HA(aq) + H2O(l)  H3O+(aq) + A-(aq) Keq = Ka = [H3O+][A-] / [HA]

  44. Now the bases . . . B(aq) + H2O(l)  BH+(aq) + OH-(aq) Keq = Kb = [BH+][OH-] / [B]

  45. Combine the two HA(aq) + H2O(l)  H3O+(aq) + A-(aq) Keq = Ka = [H3O+][A-] / [HA] & B(aq) + H2O(l)  BH+(aq) + OH-(aq) Keq = Kb = [BH+][OH-] / [B] = Kw= [H3O+][OH-]

  46. Calculations • Problem: What are the hydronium and hydroxide ion concentrations in a 0.10M HCl? HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq) Kw = [H3O+][OH-] 1.0 x 10-14 = (0.1M) x [OH-] [OH-] = 1.0 x 10-13

  47. Calculating pH • A pH is equal to the -log10 of the Hydronium ion concentration pH = -log10[H3O+] • If an acid has a [H3O+] of 1.0 x 10-2M; then the pH is 2 S.U. • S.U. = standard unit

  48. Calculating [H3O+] • The concentration of Hydronium ions can be determined by the reverse calculation, sort of • If the pH of a solution is 5.05 S.U., then the concentration of [H3O+] is 10-pH. • 5.05 S.U. = 10-5.05 = 8.9 x 10-6 M

  49. Tricks • If your concentration is 1.0 x 10-10 M, then your pH will be 10 S.U.! • The negative exponent becomes the pH • If your concentration is greater than 1.0, then the pH will be less than the negative exponent. • 2.5 x 10-10 M = >10 S.U. = 9.6 • This comes in handy when you are answering multiple choice questions about the pH of a solution.

  50. Practice Problems • What is the pH of a solution if the [H3O+] is 5.0 x 10-3M? • What is the pH of a solution if the [OH-] is 2.0 x 10-3M? • What is the [H3O+] of a solution whose pH is 3.3? • What is the [OH-] of a solution whose pH is 8.1?

More Related