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by Steven S. Zumdahl & Donald J. DeCoste University of Illinois

Introductory Chemistry: A Foundation, 6 th Ed. Introductory Chemistry, 6 th Ed. Basic Chemistry, 6 th Ed. by Steven S. Zumdahl & Donald J. DeCoste University of Illinois. Chapter 16 Acids and Bases. Properties of Acids. Sour taste Turn blue litmus paper red

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by Steven S. Zumdahl & Donald J. DeCoste University of Illinois

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  1. Introductory Chemistry: A Foundation, 6th Ed. Introductory Chemistry, 6th Ed. Basic Chemistry, 6th Ed. by Steven S. Zumdahl& Donald J. DeCoste University of Illinois

  2. Chapter 16Acids and Bases

  3. Properties of Acids • Sour taste • Turn blue litmus paper red • Change color of vegetable dyes (red cabbage juice) • React with “active” metals • Like Al, Zn, Fe, but not Cu, Ag or Au Zn + 2 HCl ZnCl2 + H2 • Corrosive • React with carbonates, producing CO2 • Marble, baking soda, chalk CaCO3 + 2 HCl CaCl2 + CO2 + H2O • React with bases to form ionic salts, and often water

  4. Properties of Bases • Also known as alkalis • Bitter Taste • Feel slippery • Change color of vegetable dyes • Different color than acid • Turn red litmus blue • React with acids to form ionic salts, and often water • Neutralization

  5. Arrhenius Theory • Acids ionize in water to H+ ions and anions • Bases ionize in water to OH- ions and cations • Neutralization reaction involves H+ combining with OH- to make water • H+ ions are protons

  6. Arrhenius Theory (cont.) • Definition only good in water solution • Definition does not explain why ammonia solutions turn litmus blue • Basic without OH- ions

  7. Brønsted-Lowry Theory • H+ transfer reaction • Since H+ is a proton, also known as proton transfer reactions Acids are proton donors, bases are proton acceptors • In the reaction, a proton from the acid molecule is transferred to the base molecule • Products are called the conjugate acid and conjugate base

  8. Brønsted-Lowry Theory (cont.)

  9. Brønsted-Lowry Theory (cont.) H-A + :B  A- + H-B+ A- is the conjugate base, H-B+ is the conjugate acid • Conjugate acid-base pair is either the original acid and its conjugate base or the original base and its conjugate acid • H-A and A- are a conjugate acid-base pair • :B and H-B+ are a conjugate acid-base pair

  10. Example #1: • Determine what species you will get if you remove 1 H+1 from the acid. • Conjugate base will have one more negative charge than the original acid H3PO4  H+ + H2PO4- Write the conjugate base for the acid H3PO4

  11. Self- check p 490 • Which of the following represent conjugate acid base pairs? • A. H2O, H3O+ • B. OH-, HNO3 • C. H2SO4, SO42- • D. HC2H3O2, C2H3O2-

  12. Brønsted-Lowery Theory (cont.) • In this theory, instead of the acid, HA, dissociating into H+(aq) and A- (aq), the acid donates its H to a water molecule HA + H2O  A- + H3O+ A-1 is the conjugate base, H3O+ is the conjugate acid

  13. Brønsted-Lowry Theory (cont.) • H3O+is calledthe hydronium ion • In this theory, substances that do not have OH- ions can act as a base if they can accept a H+1 from water. H2O + :B  OH- + H-B+ :B is acting here as a base.

  14. Strength of Acids & Bases • The stronger the acid, the more willing it is to donate H+

  15. Strength of Acids & Bases (cont.) • Strong bases will react completely with water to form hydroxide: CO3-2 + H2OHCO3- + OH- • Only small fraction of weak base molecules pull H+ off water: HCO3- + H2OH2CO3 + OH-

  16. Multiprotic Acids • Monoprotic acids have 1 acid H, diprotic 2, etc. • In oxyacids only the H on the O is acidic • In strong multiprotic acids, like H2SO4, only the first H is strong; transferring the second H is usually weak H2SO4 + H2O  H3O+ + HSO4- HSO4- + H2O  H3O+ + SO4-2

  17. Water As an Acid and a Base • Amphoteric substances can act as either an acid or a base. • Water as an acid, NH3 + H2O NH4+ + OH- • Water as a base, HCl + H2O  H3O+ + Cl- • Water can even react with itself: H2O + H2O H3O+ + OH-

  18. Autoionization of Water • Water is an extremely weak electrolyte. • Therefore there must be a few ions present H2O + H2O  H3O+ + OH-

  19. Acid Nomenclature • Acids • Compounds that form H+ in water. • Formulas usually begin with ‘H’. • In order to be an acid instead of a gas, binary acids must be aqueous (dissolved in water) • Ternary acids are ALL aqueous • Examples: • HCl (aq) – hydrochloric acid • HNO3 – nitric acid • H2SO4 – sulfuric acid

  20. Acid nomenclature • If anion ending is –ide (Binary compound), the acid name is hydro(stem)ic acid • If ternary compounds- • -Ate ending: (stem)ic acid • -ite ending: (stem)ous acid

  21. Acid Nomenclature Flowchart

  22. Solved examples • HBr – 2 elements-ide, hydrobromic acid • H2CO3- 3 elements- ate, carbonic acid • H2SO3- 3 elements- ite, sulfurous acid • Hydrofluoric acid: 2 elements= HF • Sulfuric acid: 3 elements, –ic= -ate, H2SO4 • Nitrous acid: 3 elements, -ous= -ite, HNO2

  23. Now your turn! • HI (aq) • HCl • H2SO3 • HNO3 • HIO4

  24. Hydrobromic acid • Nitrous acid • Carbonic acid • Phosphoric acid

  25. 1 x 10-14 [OH-] 1 x 10-14 [H+] [H+] = [OH-] = Acidic and Basic Solutions • Acidic solutions have a larger [H+] than [OH-] • Basic solutions have a larger [OH-] than [H+] • Neutral solutions have [H+]=[OH-]= 1 x 10-7 M

  26. Ion product of water

  27. Example #2 Determine the [H+] and [OH-] in a 10.0 M H+ solution

  28. Example #2 (cont.) • Determine the given information and the information you need to find • Given [H+] = 10.0 M, find [OH-]

  29. Example #2 (cont.) Given [H+] = 10.0 M = 1.00 x 101 M Kw = 1.0 x 10-14

  30. Self check p 497 • Calculate [H+] in a solution in which [OH-] = • 2.0X 10-2 M. Is this solution acidic, neutral or basic?

  31. pH & pOH • The acidity/basicity of a solution is often expressed as pH or pOH. • pH = -log[H3O+] pOH = -log[OH-] • pHwater = -log[10-7] = 7 = pOHwater • [H+] = 10-pH [OH-] = 10-pOH

  32. pH scales

  33. pH & pOH (cont.) • pH < 7 is acidic; pH > 7 is basic, pH = 7 is neutral • The lower the pH, the more acidic the solution; the higher the pH, the more basic the solution • 1 pH unit corresponds to a factor of 10 difference in acidity • pOH = 14 - pH

  34. pH of Common Substances

  35. Example #3 Calculate the pH of a solution with a [OH-] = 1.0 x 10-6 M

  36. Example #3 (cont.) • Find the concentration of [H+]

  37. Example #3 (cont.) • Enter the [H+] concentration into your calculator and press the log key • log(1.0 x 10-8) = -8.0 • Change the sign to get the pH • pH = -(-8.0) = 8.0

  38. Example #4 • Enter the [H+] or [OH-] concentration into your calculator and press the log key log(1.0 x 10-3) = -3.0 • Change the sign to get the pOH pOH = -(-3) = 3.0 • Subtract the calculated pH or pOH from 14.00 to get the other value pH = 14.00 – 3.0 = 11.0 Calculate the pH and pOH of a solution with a [OH-] = 1.0 x 10-3 M

  39. Solving concentration from pH or pOH • If you want to calculate [OH-] use pOH; if you want [H+] use pH. It may be necessary to convert one to the other using 14 = [H+] + [OH-] pOH = 14.00 – 7.41 = 6.59 Calculate the [OH-] of a solution with a pH of 7.41

  40. Example #5 (cont.) • Enter the pH or pOH concentration into your calculator • Change the sign of the pH or pOH -pOH = -(6.59) • Press the button(s) on you calculator to take the inverse log or 10x [OH-] = 10-6.59 = 2.6 x 10-7 M

  41. Self check p 499 • Calculate the pH value for each of the following solutions at 25°C. • A. a solution in which [H+] =1.0 X 10-9M. • B. a solution in which [OH-] =1.0 X 10-6M. • P501- A sample of rain in an area with severe air pollution has a pH of 3.5. What is the pOH of this rainwater?

  42. Calculating the pH of a Strong, Monoprotic Acid • A strong acid will dissociate 100% HA  H+ + A- • Therefore the molarity of H+ ions will be the same as the molarity of the acid • Once the H+ molarity is determined, the pH can be determined pH = -log[H+]

  43. Example #6 Calculate the pH of a 0.10 M HNO3 solution.

  44. Example #6 (cont.) • Determine the [H+] from the acid concentration HNO3 H+ + NO3- 0.10 M HNO3 = 0.10 M H+ • Enter the [H+] concentration into your calculator and press the log key log(0.10) = -1.00 • Change the sign to get the pH pH = -(-1.00) = 1.00

  45. Self check p503 • The pH of rainwater in a polluted area was measured to be 3.5. What is the [H+] in this rainwater? • The pOH of a liquid drain cleaner was found to be 10.50. What is the [OH-] for this cleaner? • P505- Calculate the pH of a solution of 5.0 X 10-3 M HCl.

  46. Buffered Solutions • Buffered solutions resist change in pH when an acid or base is added to it. • Used when need to maintain a certain pH in the system • Blood

  47. Buffered Solutions (cont.) • A buffer solution contains a weak acid and its conjugate base. • Buffers work by reacting with added H+ or OH-ions so they do not accumulate and change the pH. • Buffers will only work as long as there are sufficient weak acid and conjugate base molecules present.

  48. Buffered Solutions (cont.)

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