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Chapter 18

Chapter 18. Acids and Bases Rainbow Connection #2. Ch. 18- Acids and Bases. Acids and bases have a central role in chemistry They affect our daily life Uses: manufacturing processes, environmental issues, functioning of our bodies Acid/ Base Video. Properties of Acids. Have pH  0-6

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Chapter 18

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  1. Chapter 18 Acids and Bases Rainbow Connection #2

  2. Ch. 18- Acids and Bases • Acids and bases have a central role in chemistry • They affect our daily life • Uses: manufacturing processes, environmental issues, functioning of our bodies • Acid/ Base Video

  3. Properties of Acids • Have pH  0-6 • Tart or sour taste • Will conduct electricity • Cause indicators to change color (turns blue litmus red) • Reacts w/ metals (Mg,Zn) to form H2 gas • Neutralize w/ a base forms a salt and H20 • Ex. Citrus foods, tomatoes, vinegar

  4. Acid Formulas (memorize) • HCl ( Hydrochloric Acid) • HNO3 (nitric acid) • H2SO4 (sulfuric acid) • H2CO3 (carbonic acid) • HC2H3O2 (acetic acid) • H3PO4 (phosphoric acid)

  5. Properties of Bases • Also known as alkaline • pH  8-14 • Have a bitter taste, slippery feel • Causes indicators to change color (turns red litmus blue) • Conducts electricity • Reacts w/ acid to neutralize  form a salt and H2O • Ex. Cleaning products, soap, baking soda

  6. Acid/ Base Theories • Definitions have changed over the years as new information has been found • Arrhenius Theory video

  7. Arrhenius Theory (1887) • Applies to a H2O solution • Svante Arrhenius (Swedish) saw that not only do acids/ bases conduct electricity, they ionize (or release charged particles) when dissolved in water • Theory: • Acids- substance that ionize & produced (H+) hydrogen ions in H2O • Bases- ionize to produce (OH-) hydroxide ions in H2O

  8. Ex. HCl  H+(aq) + Cl-(aq) ACID NaOH  Na+(aq) + OH-(aq) BASE

  9. Brønsted- Lowry Theory (1923) (video) • Working independently of each other, Johannes Brønsted (Danish) and Thomas Lowry (English) defined a theory that can be used w/ all solvents not just H2O (they found that substance lost or gained protons) • Acid- in a chemical reaction, this is the substance that loses or donates a proton (H+ ion) • Base- substance that accepts or gains a proton (H+ ion = proton)

  10. HCl + H2O  H3O+ + Cl- Accepts proton (base) Donates (loses) proton – (acid) H3O+  hydronium ion (formed when H2O gains H+ ion)

  11. NH3 + H2O  NH4+ + OH- Accepts proton (base) Donates (loses) proton – (acid) Amphoteric – (H2O) acts as an acid or a base (depending on the situation)

  12. Conjugates • The particles that are formed as products can react again (reversible reaction), they behave like acids and bases (we call these conjugates) • Conjugate video

  13. Conjugate Acid - Particle that forms after the base accepts a proton (H+) from the acid • Conjugate Base – particle that remains after a proton (H+) has been released by the acid

  14. Ex. HNO3 + NaOH  H2O + NaNO3 (H-OH) Acid Conjugate Base Base Conjugate acid

  15. Disappearing ink

  16. Ex. KOH + HBr  KBr + H2O Base Conjugate Acid Acid Conjugate Base

  17. Neutralization Reaction • Occurs when an acid and a base react and there is a complete removal of all of the H+ and OH- ions • Water will be formed w/ a salt in this double displacement reaction • The solution will be neutral in pH • Important in: neutralization of soil, antacids

  18. Salt • Crystalline compound composed of the negative (non-metal) ion of the acid and the positive (metallic ) ion of the base • Salt examples: CaSO4 (plaster board), NaCl, KCl, (NH4)2SO4 (fertilizer)

  19. Examples of Neutralization Reactions • Sodium hydroxide + hydrochloric acid  sodium chloride + water • NaOH + HCl  NaCl + H2O (H-OH)

  20. Potassium hydroxide + sulfuric acid  water + potassium sulfate • KOH + H2SO4  H- OH + K2SO4 2 KOH + H2SO4 2 H-OH + K2SO4

  21. Aluminum hydroxide + acetic acid 

  22. Titration

  23. Titration • Used for a convenient method to determine the concentration or molarity of an acid or base • Uses 2 burets (long glass tubes used to measure volume)- Buret + .05 ml

  24. Definition • Analytical method in which a standard solution is used to determine the concentration of another solution • Standard solution- one in which the concentration is known

  25. Process • Using a given amount of acid in a flask, add phenolphthalein and titrate to the end point with the base (making sure to mark down all the volumes) • The whole reaction is a neutralization reaction • Use an indicator to see the endpoint in which complete neutralization occurs (wait for phenolphthalein stays light pink for 30 sec)

  26. Past endpoint Endpoint

  27. Titration Process

  28. Calculations • Reminder: M= moles • liters • So: Base (titrated soln)  • Moles (base) = volume (base) x Molarity (base) Acid (standard soln)  Moles (acid) = volume (acid) x Molarity (acid)

  29. Look at balanced equation and find the moles of the standard and the moles of the titrated unknown (look at coefficients) • Most times it is a 1:1 ratio • 1 NaOH + 1 HCl  NaCl + H2O (H-OH) • So: 1 mole (base) = 1 mole (acid) • Therefore: MA VA = MBVB

  30. Problem • A titration of 15.00 ml of HCl, required 38.57 ml of a 0.152 M base NaOH. Calculate the molarity of the HCl (acid). • MA = ?

  31. MA VA = MBVB • MA = M BVB • VA • MA = (.152M) (38.57 ml) • (15.00ml) • MA = .391 M

  32. Titration Calculation

  33. ID- A, B, CA, CB • CaCO3 + HCl  CaCl2 + H2CO3 • KOH + H3PO4 K3PO4 + H2O

  34. phosphoric acid + calcium hydroxide calcium phosphate + water • HBr + Al(OH)3 • Sulfuric acid + potassium hydroxide 

  35. Indicators • Used to find out if things are acidic or basic • Def: weak organic acids or bases whose colors differ from the colors of their conjugate acids or bases base acid

  36. Needs of Indicators • Solution being tested needs to be colorless • You need to be able to distinguish the color change • Need several indicators to cover entire pH range (0-14) • Liquid solution

  37. Indicator Examples • Litmus  red turns blue = base, blue turns red = acid • Phenolphthalein hot pink > pH 10 • Bromothymol blue  blue= base (8), Green= neutral, yellow = acid (6) • Universal Indicator  Rainbow (all pH’s 4-10) • (bromothymol blues)

  38. How Chemists use Indicators

  39. Ionization of H2O • Pure water can self ionize, it also acts as an acid or a base (amphoteric) • So: H2O(aq) H+(aq) + OH-(aq) • Experiments have shown that the concentration of [H+] = 1 X 10 –7M and [OH-] = 1 x 10 –7 M in pure water • [ ] means concentration

  40. Ion Product Constant for water • Equilibrium expression from the multiplication of the concentrations of the products • [H+] [OH-] =? • (1 x 10 –7)(1 x 10 –7) = 1.0 x 10 -14 • [H+] [OH-] = 1.0 x 10 -14

  41. If [H+] = 1.5 x 10 -6 M, what is the [OH-] = ? • [H+] [OH-] = 1.0 x 10 -14 • (1.5 x 10 -6) [OH-] = 1.0 x 10 -14 • [OH-] = 1.0 x 10 -14 • 1.5 x 10 -6 • [OH- ] = 6.7 x 10-9 M

  42. pH Concept (video) • Acidity scale developed by Soren Sorenson base on the “power of the hydrogen” • pH – measure of the hydrogen ion concentration of the solution • Equal to the negative logarithm of the hydrogen ion concentration 2. pH = - log [H+]

  43. Ex. [H+] = 1.5 x 10-8 pH= ? Graphing Calc. (-) log 1.5 (2nd) EE -8 Regular calc. 1.5 EE -8 log +/- pH= 7.8

  44. To go from pH to [H+]concentration 3. [H+] = antilog (-pH) Antilog = 10x key

  45. pH= 3.5 • [H+] = ? [H+] = antilog (-pH) [H+] = antilog (-3.5) Graphing = 2nd log -3.5 Regular 3.5 +/- 2nd log (to put in sci.not. Use 2nd #5) [H+] = 3.2 x 10 -4 M

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