University:
Bradley University
Course:
CHM 100 | Fundamentals of General Chemistry
Academic year:

2023
Views:

223

Pages:

23
Author:

Remington Spence

Chapter 15 Notes Acids and Bases Sections 15.1 – 15.11 I. Heartburn Heartburn is a painful burning sensation in the esophagus caused by hydrochloric acid (HCl) excreted by the stomach to kill microorganisms and activate enzymes that break down food. The HCl can sometimes back up out of the stomach and into the esophagus resulting in a condition known as acid reflux. Because HCl releases H+, it is the H+ ions that irritate the esophageal lining and create the burning sensation. Heartburn is treated with antacids (bases) such as Tums or Mylanta which act to neutralize the excess acid. THIS IS NOT IN THE NOTE PACKET: - HERE ARE THE 3 RULES FOR NAMING ACIDS - SEE PG. 99 IN TEXTBOOK FOR ADDITIONAL HELP - THERE ARE PRACTICE NAMING QUESTIONS ON PG. 1 OF THE WORKSHEET PACKET Anion Ending -ide -ite -ate Example Acid Name Cl-: chloride hydro- STEM-ic acid stem NO2-: nitrite STEM-ous acid stem NO3-: nitrate STEM-ic acid stem Rhyme to help you remember the 2nd/ 3rd rules Example HCl: hydrochloric acid HNO2: nitrous acid HNO3: nitric acid I ate something icky, I mite be nauseous EXAMPLES: NAME EACH ACID A.) HF: (fluoride) B.) Hydrofluoric acid C.) H2SO4: (sulfate) Carbonic acid D.) HI: (iodide) Hydroiodic acid HNO2: (nitrite) Nitrous acid Sulfuric acid E.) H2CO3: (carbonate) F.) H3PO4: (phosphate) Phosphoric acid I ate something icky, I think I mite be nauseous DO the “Naming Acids” question on the top of pg. 1 WS packet II. The Nature of Acids and Bases • General Properties of ACIDS: - Sour/tart taste - React w/ metals to form H2 gas - Can neutralize bases Some common acids are listed in Table 15.1 on pg. 664 of your text. • General Properties of BASES : - Bitter taste - Slippery feel - Can neutralize acids Some common acids are listed in Table 15.1 on pg. 664 of your text. • ALKALOID: refers to organic bases, typically produced by plants; many are toxic and/or have pharmacological effects (ex. morphine, caffeine, nicotine.) III. Definitions of Acids and Bases “The differences between the various acid-base concepts are not concerned with which is right, but which is most convenient to use in a particular situation.” - James E. Huheey ♦ THE ARRHENIUS DEFINITION: • ACID: produces H+ ions in solution Ex. HCl (aq) → H+ (aq) + Cl- (aq) • BASE: produces OH- ions in solution Ex. NaOH (aq) → Na+ (aq) + OH- (aq) • HYDRONIUM ION (H3O+):- In water, H+ always associates w/ H2O to form H3O+ - often synonymous with H+ ♦ THE BRØNSTED- LOWRY DEFINITION: • ACID: H+ donor HCl donates H+ to H2O Ex. HCl (aq) + H2O → H3O+ (aq) + Cl- (aq) • BASE: H+ acceptor NH3 accepts H+ from H2O Ex. NH3 (aq) + H2O → NH4+ (aq) + OH- (aq) How to tell if a substance can be an “H+ acceptor:” - NH3 is a H+ acceptor because it has a pair of unshared eon N which attracts the H+ ; •• H N H - Other H+ acceptors would be those that have a negative charge, like OH− Note: All Arrhenius acids are also Brønsted-Lowry acids and vice-versa. The main difference between the two definitions is in reference to bases. H • AMPHOTERIC: can act as both acids or bases by either donating or accepting H+; How to tell if it is amphoteric: * H2O, NH3, and species w/ an H & a negative charge, like HSO4− Example: HSO4─ Example: H2O H2O + H+ ↔ H3O+ (base: acceptor) HSO4─ + H+ ↔ H2SO4 (base: acceptor) H2 O ↔ H+ + (acid: donor) HSO4─ ↔ H+ (acid: donor) OH- + SO42- • CONJUGATE ACID-BASE PAIR: Two substances that differ by only one H+ Ex. H2O/ H3O+, NH3/ NH4+ conjugate acids accepting H+ • Bases form _______________________ by _____________ conjugate bases donating • Acids form _______________________ by _____________ H+ Ex. Ex. (donor) (acceptor) acid base NH3 + H2O (donor) acid H2SO4 ↔ conj. acid conj. base NH4+ + OH- (acceptor) base + H2O ↔ conj. acid conj. base HSO4- + H3O+ DO pg. 2/3 of the WS packet * Help w/ 5 a): add an H, one more H, charge increases * Help w/ 6 a): remove an H, one less H, charge decreases IV. Acid Strength and Ionization Constant (Ka) VII. Base Solutions • STRONG ACID: completely ionizes in solution (use one-way arrow) Ex. HCl (aq) + H2O (l)  H3O+ (aq) + Cl- (aq) There are six strong acids (you must know these): HCl HBr HI HNO3 H2SO4 HClO4 POLYPROTIC ACID: an acid that contains > 1 ionizable H+ Ex. H2SO4, H3PO4 The H’s dissociate 1 at a time : Ex: H2SO4  H+ + HSO4͞ • WEAK ACID:only partially ionizes in solution (use two-way arrow) Ex. HF (aq) + H2O (l) ↔ H3O+ (aq) + F- (aq) • ACID IONIZATION CONSTANT (Ka): - quantifies the relative strength of weak acids - ratio of dissociated acid (H3O+ and A-) to undissociated acid (HA) - the larger the value, the stronger the acid Ex. HF (Ka = 3.5 x 10-4) HClO (Ka = 2.9 x 10-8) (See pg. A-12) HF (aq) + H2O (l) ↔ H3O+ (aq) + F͞ (aq) Ka = [H3O+] [F͞ ] [HF] • STRONG BASE: completely ionizes in solution (use one-way arrow Ex. NaOH (aq) → Na+ (aq) + OH¯ (aq) There are six strong bases (you must know these): LiOH NaOH KOH Ca(OH)2 Sr(OH)2 Ba(OH)2 The Group 2 bases yield 2 equivalents of OH- when they dissociate: Ex. Ca(OH)2 Ca2+ + 2 OH- • WEAK BASE: only partially ionizes in solution (two-way arrow) Ex. NH3 (aq) + H2O (l) ↔ NH4+ (aq) + OH− (aq) • BASE IONIZATION CONSTANT (Kb): - quantifies the relative strength of weak bases - the larger the value, the stronger the base Ex. NH3 (Kb = 1.76 x 10-5 ) CO32- (Kb = 1.8 x 10-4) (See pg. A-13)) V. Autoionization of Water and pH We saw earlier that water is amphoteric: it can act as either an acid or a base. Even in pure water, water acts as an acid and a base with itself. • AUTOIONIZATION: Reaction of water dissociating into its ions: H2O (l) + H2O (l) ↔ H3O+ (aq) + OH-(aq) • ION PRODUCT CONSTANT FOR WATER (Kw): the equilibrium constant for the autoionization of water Kw = [H3O+] x [OH-] = 1.0 x 10-14 Notice that Kw is numerically very small, so there is very little H3O+ or OH- in solution; Also, water is neutral because [H3O+] = [OH-] • ACIDIC SOLUTION: Contains more H3O+ than OH− • BASIC SOLUTION: Contains more OH− than H3O+ Summarizing: +] = [OH−] [H O ♦ A neutral solution contains 3 ♦ An acidic solution contains [H3O+] > [OH−] ♦ A basic solution contains [OH−] > [H3O+] ♦ All aqueous solutions contain [H3O+] x [OH−] = 1.0 x 10-14 Calculate either the missing [H3O+] or [OH-]. Label each solution as acidic, basic, or neutral. 1. [H3O+] = 1.0 x 10-7, [OH-] = 1.0 x 10-7 Neutral 2. [H3O+] = 2.5 x 10-4, [OH-] = 4.0 x 10-11 Acidic 3. [OH-] = 4.0 x 10-9, [H3O+] = 2.5 x 10-6 Acidic 4. [OH-] = 3.5 x 10-5, [H3O+] = 2.9 x 10-10 Basic ♦ THE PH SCALE IS A CONVENIENT WAY TO SPECIFY THE ACIDITY OF A SOLUTION • pH: pH = − log [H3O+] [H3O+] = 10^− pH In general: pH Type of Solution [H3O]+ <7 Acidic > 1.0 x 10-7 =7 Neutral = 1.0 x 10-7 >7 Basic < 1.0 x 10-7 • The smaller the pH, the stronger the acid • The larger the pH, the stronger the base Ex. Calculate the pH of each solution and state whether it is acidic or basic: a.) [H+] = 1.4 x10-4 M H+ and H3O+ are Used interchangeably b.) [OH-] = 1.3 x 10-2 M pH = − log [H+] [H3O+] x [OH−] = 1.0 x 10-14 pH = − log [1.4 x 10-4] [H3O+] = 1.0 x 10-14 = 7.69 x 10-13 1.3 x 10-2 pH = 3.85 pH = − log [H3O+] Acidic (pH < 7) pH = − log [7.69 x 10-13] pH = 12.11 Basic (pH > 7) Ex. Calculate the [H3O+] of a solution with a pH of 4.80 [H3O+] = 10^− pH [H3O+] = 10^− 4.80 [H3O+] = 1.58 x 10− 5 Ex. Calculate the [H3O+] and pH of each solution and state whether it is acidic or basic: a.) 0.225 M KOH b.) 0.0015 M Sr(OH)2 SKIP [H3O+] x [OH−] = 1.0 x 10-14 [H3O+] = 1.0 x 10-14 = 4.44 x 10-14 0.225 pH = − log [H3O+] pH = − log [4.44 x 10-14] pH = 13.35 Basic (pH > 7) ANOTHER USEFUL RELATIONSHIP IS: pH + pOH = 14 Ex. Calculate the pH of a solution with a pOH of 11.2 pH + pOH = 14 so pH = 14 - 11.2 pH = 2.8

H+ Acceptor

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