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Acids & Bases

Acids & Bases. Unit 12 Chapter 16. What are Acids & Bases?. The term acid comes from the Latin word acidus , meaning “ sour. ” Sour tasting foods and beverages have at least a little bit of acid in them. Bases can be thought of as the opposite of acids.

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Acids & Bases

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  1. Acids & Bases Unit 12 Chapter 16

  2. What are Acids & Bases? • The term acid comes from the Latin word acidus, meaning “sour.” • Sour tasting foods and beverages have at least a little bit of acid in them. • Bases can be thought of as the opposite of acids. • They tend to taste bitter & feel slippery or “soapy”.

  3. Acids or Bases = Bad day • Acids are Corrosive and dissolve metals via a single displacement reaction. • Acids denature proteins and can be used to “cook” foods (pickling). • Bases are also corrosive (called Caustic) and are very dangerous to body tissue. • Bases saponify the fats & oils in body tissue, thereby causing it to disintegrate.

  4. Acids + Bases = Beach Party! • Oddly enough, when you combine two very dangerous substances, you get two pretty innocuous products. • Acids and bases neutralize each other creating water and a salt. • HCl + NaOH  H2O + NaCl • H2SO4 + Ca(OH)2  2H2O + CaSO4

  5. Definitions • There are several definitions of Acids & Bases, but the first and most common is from 1884 by Svante Arrhenius, a Swedish scientist. • Arrhenius Acids increase the concentration of H+ ions. • Arrhenius Bases increase the concentration of OH- ions. Svante Arrhenius 1859-1927

  6. Two More Blokes • A slightly different definition was given by two different scientists at the same time, Johannes Brønsted from Denmark & Martin Lowry from England in 1923. • Brønsted-Lowry acids donate a proton to another substance. • Brønsted-Lowry bases accept a proton from another substance. Johannes Brønsted 1879-1947 Martin Lowry 1874-1936

  7. Not to be Outdone! • Not totally happy with either definition, G.N. Lewis (remember him?), publishes his own definition in 1923. • A Lewis Acid accepts a lone pair to complete its valence shell. • Lewis Bases donate a lone pair to complete the valence shell of another substance. G.N. Lewis 1875-1946

  8. Comparing the Three • Since H+ is a proton, the Arrhenius & Brønsted-Lowry definitions are typically identical (exceptions are limited to narrow types of chemistry). • HA = generic symbol for an Acid • BOH = generic symbol for a Base • Lewis’ definition is much more broad.

  9. That Crazy Lewis… • All ions are Lewis Acids/Bases: • Ca+2 + SO4-2 CaSO4 • I2 + I-  I3- • Some aren’t even ions… Lewis Acid Lewis Base Lewis Acid Lewis Base Adduct

  10. It Just Tears ‘em Apart • Acids and Bases dissociate in water. • When an acid dissociates, it produces H+ ions which immediately attach to a molecule of water forming the hydronium ion, H3O+. • When a base dissociates, it produces the hydroxide ion, OH-.

  11. Polyprotism…Sounds Contagious • Polyprotic acids have more than 1 acidic hydrogen that can be donated. • H3PO4 is triprotic: • H3PO4 H+ + H2PO4- • H2PO4-  H+ + HPO4-2 • HPO4-2  H+ + PO4-3 • H2SO4 & H2CO3 are diprotic • Just for added confusion  polybasic

  12. ‘Cause they work out… • The rate at which an acid or base will dissolve something is related to its concentration. • For acids and bases, the terms Strong and Weak have nothing to do with their dissolving power or concentration.

  13. Strength in Numbers • A strong acid is an acid that ionizes completely in a solvent, while a weak acid dissociates very little. • Weak acids create few hydronium ions, H3O+, in the solution. • Similarly, a strong base is a base that ionizes completely in a solvent, while a weak base dissociates very little, • Weak bases leave few free hydroxide ions, OH-, in the solution.

  14. Strong Acids • HCl • HNO3 • H2SO4 (first proton only) • HBr • HI • HClO4 (All others are weak)

  15. Strong Bases • LiOH • NaOH • KOH • RbOH • CsOH • Ca(OH)2 • Sr(OH)2 • Ba(OH)2 (All others are weak)

  16. Acid Strength & Concentration • The concentration of a strong acid is always 0* • 3.0 M HCl is actually: • 0 M HCl* • 3.0 M H+ • 3.0 M Cl- • Weak acids are the main component present in solution. * At really high concentrations, strong acids do not completely dissociate

  17. Conjugate Acid/Base Pairs HA(aq) + H2O(l) H3O+(aq) + A(aq) acid base conj. acid conj. base • Conjugate base: The anion or substance that remains after a proton is lost from an acid. • Conjugate acid: The cation or substance formed when the proton is transferred to the base. • (The term “conjugate” implies that the substance exists on the products side of a reaction)

  18. Relativity Rules • Strong acids have weak conjugate bases. • Strong bases have weak conjugate acids. • H2SO4 H+ + HSO4- • H2SO4 = strong acid (gives up one H+ readily) • HSO4- = weak base (doesn’t like to accept H+)

  19. Water, Water Every…Where? • The conjugate species of a weak acid or base will hydrolyze water: • Na2CO3 2 Na+(aq) + CO3-2(aq) • CO3-2 is the conjugate base of a weak acid (HCO3-) • CO3-2 + H2O  HCO3- + OH- • Which is the conjugate base of H2CO3 • HCO3- + H2O  H2CO3 + OH- H2O

  20. Water, Water, Everywhere! • Water is the most important solvent on the planet… • It’s also amphoteric: • H2O + HCl  H3O+ + Cl- (acting as a base) • H2O + NH3 NH4+ + OH- (as an acid) • Water is also a weak acid & a weak base • H2O  H+ + OH- (at the same time!!!)

  21. It’s (still) All Relative • Amphoteric Substances can act as acids & bases • A base is added to a solution of bisulfate: HSO4- + OH- SO4-2 + H2O • An acid is added to a solution of bisulfate: H3O+ + HSO4- H2SO4 + H2O Acid Conjugate Base Base Conjugate Acid

  22. It Can’t Decide • Water will auto-ionize – it has its own equilibrium constant, Kw! • KW = [H+][OH-] = 1x10-14 • It’s so small, that very little ionizes, but, it shows the relationship between [H+] & [OH-] • Adding more [H+] causes [OH-] to go down (& vice versa)

  23. Test Some Grey Matter • What is [H+] and [OH-] in a neutral solution? • [H+][OH-] = 1x10-14 • Neutral  [H+] = [OH-] • [H+]2 = 1x10-14 • [H+] = (1x10-14)0.5 • [H+] = [OH-] = 1x10-7 M

  24. It’s The Power of H! • pH is a measure of low concentrations of acids (or bases). • pH = -log[H+] [H+]=10-pH • Neutral pH = 7. (because [H+] = 1x10-7 M) • Anything less than 7 is acidic. • Above 7 is basic. • Typically, the scale ranges from 0-14; however, negative pH’s are possible as well as pH’s > 14.

  25. The Yin to H’s Yang • pOH is a measure of low concentrations of bases (or acids) • Just like pH, pOH = -log[OH-] • Typically, however, we talk about pH. • pH can only be calculated from [H+], but in a base, we do not necessarily know the hydrogen concentration (it’s really low).

  26. Using Kw • Instead, we take advantage of the Kw equation and see: • -log(1x10-14) = -log([H+][OH-]) • 14 = pH + pOH • If we know the pOH, then pH = 14 - pOH. • Also, [OH-] = 10-pOH

  27. Getting Hit With a Log • pH = -log[H+]. • 0.025 M HCl has a pH of: • pH = -log[0.025] = 1.6 • If the pH is known, [H+] = 10-pH • An acid has a pH of 3.5, what is [H+]? • [H+] = 10-3.5 = 3.2x10-4 M

  28. OH my H! • What is the pOH of a solution with a pH of 6.5? • pOH = 14 - pH = 14 – 6.5 = 7.5 • What is the [OH-] of a solution with a pH of 13? • pOH = 14 – pH = 14 – 13 = 1 • [OH-] = 10-1 = 0.1 M OH-

  29. Indicators • An Indicator is a compound that can reversibly change color depending on the pH of the solution. • Litmus was one of the first pH indicators, used in about 1300 CE by Spanish alchemist, Arnaldus de Villa Nova. • Litmus turns red in acids and blue in bases. Arnaldus de Villa Nova 1235-1311

  30. Back & Forth & Back & … • Indicators change color based on a moving equilibrium: • Ind- + H+ HInd (acid color) • HInd + OH- H2O + Ind- (basic color)

  31. pH Indicators indicate pH • The Transition range of an indicator is the pH range through which it changes color. • Not all indicators change color at the same pH. • Some indicators change well within the pH range of acids, others around neutral, and still others change well within the range of bases.

  32. Universal Indicators • The most common product with a universal indicator is pH paper. • Another easily manufactured universal indicator is the mixture of anythocyanins in red cabbage.

  33. Even-Stevens • Neutralization is the act of Balancing the concentration of H+ with OH-. • At equal concentrations, the majority of acid (H3O+) and base (OH-) combine to form water with a pH of 7.

  34. Tie-tray-shin • Titration is a technique where an unknown solution is reacted with a solution with a known concentration in order to determine the precise concentration of the unknown. • In Acid/Base titrations, an Indicator is used to tell us when the End Point is reached by a color change.

  35. Alright, Stop! • An End Point is the precise moment that the indicator changes color. • The titration stops because no more information can be gleaned from the rxn. • The end point is determined by the pH at which the indicator changes color. • Choice of indicator is, therefore, very important.

  36. Unequivocally Equal… • The Equivalence Point is the precise moment that the number of moles of acid = the number of moles of base. • The pH is, by default, 7 for the titration of a strong acid with a strong base. • For the titration of weak acids & bases, the pH is determined by the dissociation constants.

  37. Buffers! • A buffer is a solution that resists changes in pH. • This is due to the presence of a weak acid and its conjugate base • Or a weak base and its conjugate acid. • This combination allows the solution to absorb H+ and OH- ions without much effect.

  38. Meh, not much effect… • In a carbonic acid buffered system like your blood, H2CO3(aq) and HCO3-(aq) absorb both acid and base: Adding Acid: H2CO3(aq) + HCO3-(aq) + H+(aq) 2 H2CO3(aq) Adding Base: H2CO3(aq) + HCO3-(aq) + OH-(aq) 2 HCO3-(aq) No Free Acid or Base! Buffer System

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