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Importance of Chemical Reactions in Water: Acids and Bases Overview

Explore why studying chemical reactions in water is crucial, including properties of acids and bases, nomenclature, common substances, and acid-base strength models.

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Importance of Chemical Reactions in Water: Acids and Bases Overview

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  1. Why study chemical reactions in water? 3/4 of the earth surface covered by water Life, as we know it, evolved in water Reactions occur quickly in water Water is relatively cheap (cents per gallon) Reactions in water can be grouped into categories: Reactions of acids and bases Reactions that form precipitates Reactions involving oxidation and reduction

  2. Acids and Bases

  3. Properties of Acids • Aqueous solution have sour taste • Change the color of acid / base indicators • Litmus paper for example turns red • React with active metals to release H2 gas 2HCl + Mg  MgCl2 + H2

  4. Properties of Acids • React with bases to produce salts and H2O (neutralization) • HCL + NaOH NaCl + H2O • Conduct electricity - when put in water they form ions, they are considered electrolytes

  5. Acid Nomenclature • Binary Acid - contain H and 1 other element. • Prefix “hydro” • Root “name of 2nd element” • Suffix “ic” • Ex. Hydrochloric Acid HCl Hydroiodic Acid - HI

  6. Oxyacid • H, O and a 3rd element (mostly non-metal) Root - name the oxy ion, replace suffix • Suffix “ic” for “ate” “ous” for “ite” ions Carbonic Acid H2CO3 Sulfuric Acid H2SO4 2 H+ + SO42- sulfate Sulfurous Acid H2SO3 2 H+ + SO32- sulfite

  7. Common household substances that contain acids and bases. Vinegar is a dilute solution of acetic acid. Drain cleaners contain strong bases such as sodium hydroxide.

  8. Common Industrial Acids • H2 SO4 - sulfuric Acid - most produced chemical in the world. Metallurgy, fertilizer, paper, paint, dyes • HNO3 nitric acid- suffocating odor, stains, protein yellow - explosives, rubber, plastics • H3PO4 - phosphoric acid - fertilizer, flavors beverages, cleaning agent • HCl - hydrochloric acid - found in stomach, cleaning agent, food processing, dilute forms called “muriatic acid” in stores. • CH3COOH - Acetic Acid- foul smelling, vinegar - 4% - 6% acetic acid : plastics, food additives.

  9. Bases • Aqueous solutions that taste bitter • Change color of acid/ base indicators • Litmus paper turns blue in the presence of a base • Dilute aqueous solutions feel slippery

  10. Bases • React with acids to produce salts and water (neutralization) HCL + NaOH NaCl + H2O • Conduct electricity – when put in water they form ions, they are considered electrolytes

  11. Strong Acid • ionizes completely in water - strong electrolyte. • Remember covalent molecules ionize! Complete Ionization equation example: HBr + H2O H3O+ + Br- Ex. 6 strong acids that will always ionize completely • 3 strong oxyacids: H2SO4, HClO4, HNO3 , • 3 strong binary acids: HCl, HBr, HI

  12. Weak Acid • partially ionizes - weak electrolyte. • Incomplete ionization equation example HCN + H2O  H3O + + CN – *Remember the arrow for weak ionization equations mean that the molecule does not form ions as readily and will stay in the molecular form more often Ex. H3PO4, HF, CH3COOH, H2CO3, H2S

  13. Strong Base • strong electrolyte - completely dissociate. H2O NaOH  Na+ + OH- Ex. NaOH , KOH, LiOH (Strong bases are generally metallic ions in the first and second column on the periodic table that combine with OH- )

  14. Weak Base • Weak electrolyte- partially ionizes; Most of it stays in its molecular form • Notice this weak base ionizes! • Ex. NH3 + H2O  NH4+ + OH- • The arrow is showing the direction of the equation

  15. Acid/Base Strength Strong Acids HNO3 HClO4 H2SO4 HCl HBr HI Strong Bases NaOH KOH LiOH Ba(OH)2 Ca(OH)2 (slightly soluble) Sr(OH)2(slightly soluble)

  16. Models of Acids and Bases • Arrhenius Concept: Acids produce H+ ions in solution, bases produce OH ions in solution HCl H+ + Cl- NaOH  Na+ + OH – • Brønsted-Lowry: Acids are H+ donors, bases are proton acceptors. HCl + H2O  Cl + H3O+ acid base NH3 + H2O  NH4+ + OH – base acid Lewis Concept: Acids are electron pair accepters, bases are electron pair donors (we will not need to learn about this concept!)

  17. Arrhenius Definition • Based on the idea that aqueous solutions of acids or bases produce ions • Acid – increases the H3O+ or H+ concentration in water HCl H+ + Cl- H2SO4 2H+ + SO42- • Base – increases the OH- concentration in water NaOH  Na+ + OH- NH4OH  NH4+ + OH-

  18. Bronsted - Lowry Acids and Bases • Based on whether a substance is a proton acceptor or donor in non-aqueous solutions. • Bronsted-Lowry is a way to study proton transfer!! • Acid – H+ donor Base – H+ acceptor • HCl + NH3NH4+ + Cl - • HCl donated a proton (H +) to NH3. • Proton donor - acid • The NH3 accepted a proton from HCl • proton acceptor- base

  19. Bronsted - Lowry Acids and Bases • Bronsted-Lowry is a way to study proton transfer!! • Acid – H+ donor Base – H+ acceptor • For Example: HCl + NH3 NH4+ + Cl- H2SO4+ 2H2O2 H3O+ + SO42-

  20. Amphoteric • can react as either an acid or base HCl + H2OH3O + + Cl - proton acceptor (water) H2O + NH3NH4+ + OH - proton donor (water)

  21. Conjugate Acids/Bases Conjugate Acid – formed when BL base gains a proton Conjugate Base – formed when BL acid looses a proton

  22. Conjugate Acids/Bases HCl + NH3 NH4+ + Cl- H2SO4+2H2O2 H3O+ + SO42- Acid Base Conj. Acid Conj. Base Acid Base Conj. Acid Conj. Base

  23. Conjugates Strength • The stronger the acid, the weaker its conjugate base; the stronger the base, the weaker its conjugate acid. • Proton transfers favor the production of weaker acids and weaker base. Therefore, CH3COOH + Water  H3O ++ CH3COO - (weak acid)(weak base) (stronger acid)(stronger base) Reactants are favored!!

  24. Monoprotic Acids Ionization– when ions are formed from solute molecules by the action of the solvent. Monoprotic donates 1 hydrogen For example: H2O (l) + HCl (s)  H3O+ (aq) + Cl- (aq) Hydronium ion = H+

  25. Polyprotic Acids/Bases Some acids have more than one ionizable hydrogen and are called polyprotic: diprotic (2 H+), triprotic (3 H+). For example: 2 H2O (l) + H2SO4 (s)  2 H3O+ (aq) + SO42- (aq) Two moles of hydronium ions

  26. Polyprotic Acids Ionizationis in several distinct steps: e.g., H2CO3: carbonic acid H2CO3 + H2O H3O+ + HCO3- HCO3- + H2O H3O+ + CO32- Transfer of 2nd (or 3rd) proton are more difficult than 1st.

  27. Self-ionization of water • a Very weak electrolyte. • Autoionization of water:  • H2O (l) + H2O (l) H3O+(aq) + OH-(aq) hydroniumhydroxide • When protons (H+) are produced in water, they bind to the lone pair e- of water to produce H3O+

  28. Acid/Base Equilibria At 25 oC pure water, has a pH = 7 [H3O+] = [OH-] = 1 x 10-7 We use the [ ] to indicate concentration (which is measured in moles/liter or Molarity!) Kw = ionization constant for water Kw = 1.0 X 10-14 @ 25 oC Note: your book presents the autoionization of water on the reaction: H2O H+ + OH-; [H+] is analogous to [H3O+] !

  29. Acids and Bases H2O = HOH = H+ + OH- AcidsBases HClNaOH HNO3KOH HFNH4OH Remember water can be written like HOH because you are combining a hydrogen ion from the acid and a hydroxide ion from the base to produce water!

  30. The pH Scale • pH in water ranges from 0 to 14. • Kw = 1.00  1014 = [H3O+] [OH] • pH + pOH = 14.00 • As pH rises, pOH falls (sum = 14.00)

  31. The pH Scale • pH is measured in exponential values A pH change from 6 to 5 is 10 times more [H3O+] A pH change from 6 to 4 is 100 times more [H3O+] A pH change from 6 to 3 is 1000 times more [H3O+] • Reminder: as exponential values get larger the value gets smaller • Ex: 1 x 10 -3 = 0.001 1 x 10 -6 = 0.000001

  32. pH Scale • pH – stands for power of hydrogen,related to the concentration of H3O+ ions in solutions. • The more H3O+ ions, the lower the pH. • pH dictates if something is an acid or a base • Ex. pH 3 = acid pH 9=base

  33. The pH scale and pH values of some common substances.

  34. pH scale pH 7 = H3O+ ions and OH- ions are equal pH 0 = many H3O+ ions and few OH- ions pH 14= many OH- ions and few H3O+ ions Good conductor Non- conductor Good conductor

  35. HIn  H+ + In-For phenolphthalein: pH 0 to 8.2 = colorless; then pink, then pH 10 = red Add H+ then shift to left Add OH- then shift to right

  36. Phenolphthalein indicator simple diagram When the H+ leaves the Indicator it changes shape, which causes it to change color b/c it will refract a different wavelength of light • In = indicator When you add an OH- they attract the H+ OH-  In + In- H+ + OH- H+ Pink in the presence of a base H+ H and OH will eventually combine to form water Clear in the presence of an acid This is reversible by adding an acid What makes the slight pink color? Occurs when There are just slightly more OH- ions than H+ ions at the indicators end point so as soon as all of the H+ ions are used up the pink color will show up.

  37. pH Calculation Practice [H3O+] pH AorB pH [H3O+] AorB 1x10-2 2 A 1 1x10-1 A 1x10-4 3 1x10-6 5 1x10-7 7 1x10-8 9 1x10-10 11 1x10-12 13 1x10-14 14 pX = -log [ion]; [ion] = 10-pX

  38. pH Practice [H3O+] pH AorB pH [H3O+] AorB 1x10-2 2 A 1 1x10-1 A 1x10-4 4 A 3 1x10-3 A 1x10-6 6 A 5 1x10-5 A 1x10-7 7 N 7 1x10-7 N 1x10-8 8 B 9 1x10-9 B 1x10-10 10 B 11 1x10-11 B 1x10-12 12 B 13 1x10-13 B 1x10-14 14 B 14 1x10-14 B pX = -log [ion]; [ion] = 10-pX

  39. More pH Practice Item pH [H3O+] AorB blood 7.4 4.0x10-8 N apples 3.1 milk of mag. seawater 8.5 stomach acid 2.0 eggs 7.8 bananas pX = -log [ion]; [ion] = 10-pX 3.2x10-11 .000025

  40. More pH Practice Item pH [H3O+] AorB blood 7.4 4.0x10-8 N apples 3.1 milk of mag. seawater 8.5 stomach acid 2.0 eggs 7.8 bananas pH = -log [H3O+]; [H3O+] = 10-pH 7.9x10-4 A 3.2x10-11 10.5 B 3.2x10-9 B 1.0x10-2 A 1.6x10-8 B .000025 4.6 A

  41. Item pOH [OH-] AorB More pH Practice pX = -log [ion]; [ion] = 10-pX

  42. Item pOH [OH-] A or B More pH Practice pX = -log [ion]; [ion] = 10-pX (for A or B, think pH) 3.98 x 10-11 10.4 OJ A (pH 3.6) B (pH 9.9) 7.9 x 10-5 4.1 Toothpaste Rainwater 3.2 x 10-8 A (pH 6.5) 7.5 Great Salt Lake 3.9 .00012 B (pH 10.1)

  43. Kw [OH-] 1 x 10-4 Kw = [H3O+][OH-] = 1 x 10-14 [H3O+] 1 x 10-10 1 x 10-14 1 x 10-14 1 x 10-9 1 x 10-5 10.08 x 10-15 2.1 x 10-11 4.8 x 10-4 10 x 10-15 round Only one number to left of decimal 1 x 10-14 9.9 x 10-15 6.2 x 10-9 1.6 x 10-6 1 x 10-14

  44. Kw and pOH Practice Item [H3O+] [OH-] pH blood 7.4 4.0x10-8 3.1 apples 7.9x10-4 milk of mag. 10.5 3.2x10-11 seawater 8.5 3.2x10-9 stomach acid 2.0 1.0x10-2 eggs 7.8 1.6x10-8 bananas 4.6 2.5x10-5 pX= -log [ion]; pH + pOH = 14; Kw = [H3O+][OH-] = 1 x 10-14 pOH

  45. Kw and pOH Practice Item [H3O+] [OH-] pH blood 4.0x10-8 2.5x10-7 7.4 apples 7.9x10-4 1.27x10-11 3.1 milk of mag. 3.2x10-11 3.1x10-4 10.5 seawater 8.5 3.2x10-9 3.1x10-6 stomach acid 2.0 1.0x10-2 1.0x10-12 eggs 7.8 1.6x10-8 6.25x10-7 bananas 4.6 2.5x10-5 4.0x10-10 pX= -log [ion]; pH + pOH = 14; Kw = [H3O+][OH-] = 1 x 10-14 pOH

  46. pH + pOH = 14 pX = -log [ion]; [ion] = 10-pX

  47. pH + pOH = 14 pX = -log [ion]; [ion] = 10-pX

  48. pH word problems If the [H3O+] ion concentration is 3.8 x 10-9 M what is the pOH? 2 step problem • Calculate the pH pH= - log [3.8 x 10-9 ] =8.4 2) Subtract the pH from 14 to get the pOH 14 - 8.4 = 5.6

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