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Unit 13

Unit 13. Acids and Bases. electrolyte. Properties. ACIDS. BASES. electrolyte. sour taste. bitter taste. sticky feel. slippery feel. turn litmus blue. turn litmus red. react with acids to form water and a salt (ionic compound).

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Unit 13

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  1. Unit 13 Acids and Bases

  2. electrolyte Properties ACIDS BASES • electrolyte • sour taste • bitter taste • sticky feel • slippery feel • turn litmus blue • turn litmus red • react with acids to form water and a salt (ionic compound) • react with bases to form water and a salt (ionic compound)

  3. ACIDS: Most citrus fruits, tea, battery acid, vinegar, milk, soda, apples. BASES: Common household bases include baking soda, lye, ammonia, soap, and antacids. . Examples

  4. Indicators • Indicators are substances that change color in the presence of an acid or a base • Indicators are made up of weak acids or weak bases • Examples of indicators include pH paper, red and blue litmus paper, and phenolphthalein

  5. Acids Affect Indicators: Blue litmus paper turns red in contact with an acid. It remains blue when in contact with a base or neutral solution.

  6. Bases affect indicators: Red litmus paper turns blue in contact with a base. It remains red when in contact with an acid or neutral solution. Phenolphthalein turns pink in a base. It is colorless in an acid or neutral solution.

  7. Definitions • There are 3 definitions used to describe acids and bases: • Arrhenius • BrØnsted-Lowry • Lewis • The most traditional is Arrhenius acids and bases.

  8. H H – + O O Cl Cl H H H H Definitions • Arrhenius - In aqueous solution… • Acidsform hydrogen ions (H+) HCl + H2O H++ Cl– Also called hydronium ions (H3O+) acid

  9. H H – + N O O N H H H H H H H H Definitions • Arrhenius - In aqueous solution… • Basesform hydroxide ions (OH-) NH3+ H2O  NH4+ +OH- base

  10. conjugate base conjugate acid Another common way to refer to hydrogen ions is to call them “protons” Definitions • Brønsted-Lowry • Acidsare proton (H+) donors. • Bases are proton (H+) acceptors. HCl + H2O  Cl– + H3O+ acid base Conjugate acid – particle formed when a base gains a H+ Conjugate base – particle that remains when an acid has donated a H+ • .

  11. Definitions • Lewis • Acidsare electron pair acceptors. • Bases are electron pair donors. Lewis base Lewis acid

  12. White Board Questions ACID Proton H3O+ H+ Blue turns red • When you wafted a substance your nose burned. Would this substance be an acid or a base? • A hydrogen ion (H+) can also be called a _________ or ____________. • Arrhenius acids are compounds that break up in water to give off _____________. • What color litmus paper would you use to test an acid? What color will it turn? 5. If your food tastes bitter, which do you think it could possibly be an acid or a base? BASE

  13. White Board Questions accepts base Arrhenius acid base 6. A BrØnsted-Lowry base _________ hydrogen ions. 7. Phenolphthalein turns pink when it comes in contact with a(n) _________. 8. Which of the scientists defined the typical acid? 9. If you are eating and it has a sour taste, would that be an acid or a base? 10. If a piece of red litmus paper turns blue than it is a(n) ___________.

  14. Naming Acids • Binary acids • Contains 2 different elements: H and another • Always has “hydro-” prefix • Root of other element’s name • Ending “-ic”

  15. Examples of Binary Acids • HI is hydroiodic acid • H2S is hydrosulfuric acid • HBr is hydrobromic acid • HCl is hydrochloric acid

  16. Naming Acids • Ternary Acids - Oxyacids • Contains 3 different elements: H, O, and another • No prefix • Name of polyatomic ion (p. 147) • Ending “–ic” for polyatomic ion ending in “-ate” and “–ous” for ion ending in “-ite”

  17. Examples of Ternary Acids • ClO3 is chlorate so HClO3 is chloric acid • PO4 is phosphate so H3PO4 is phosphoric acid • PO3 is phosphite so H3PO3 is phosphorous acid • NO2 is nitrite HNO2 is nitrous acid • NO3 is nitrate HNO3 is nitric acid

  18. Naming Acids cont. • HC2H3O2 or CH3COOH Name is acetic acid Common name = vinegar

  19. H2SO3 Sulfurous acid HF Hydrofluoric acid H2Se Hydroselenic acid Perchloric acid HClO4 Carbonic acid H2CO3 Hydrobromic acid HBr Practice Naming Acids

  20. Ion Product of Water Self- ionization of water – the simple dissociation of water H2O H+ + OH- Concentration of ea. ion in pure water: [H+] = 1.0 x 10-7M + [OH-] = 1.0 x 10-7M Ion-product constant for water (Kw), Where Kw = 1.0 x 10-14 Kw = [H+] [OH-] Acid [H+] > [OH-] Base [H+] < [OH-] Neutral [H+] = [OH-]

  21. Calculating [H+] and [OH-] • reverse the pH equation • The pH of a solution is 8. Find the [H+] and [OH-] and determine whether it is acidic, basic, or neutral. • basic [H+] = 1 x 10-pH and [OH-] = 1 x 10-pOH [H+] = 1 x 10-8 M [OH-] = 1 x 10-(14-8) M = 1 x 10-6 M

  22. Examples 1. If the [H+]in a solution is 1.0 x 10-5M, is the solution acidic, basic or neutral? 1.0 x 10-5 M What is the concentration of the [OH-]? Use the ion-product constant for water (Kw): Kw = [H+] [OH-] 1.0 x 10-14 = [1.0 x 10-5] [OH-] 1.0 x 10-14 = [OH-] 1.0 x 10-5 1.0 x 10-(14-5) pH 5 = acidic 1.0 x 10-9 M

  23. Examples 2. If the pH is 9, what is the concentration of the hydroxide ion? Kw = [H+] [OH-] 1.0 x 10-14 M = [1.0 x 10-9M] [OH-] 1.0 x 10-5 M = [OH-] 14 = pH + pOH 14 = 9 + pOH 5 = pOH 3. If the pOH is 4, what is the concentration of the hydrogen ion? Kw = [H+] [OH-] 1.0 x 10-14 M = [H+] [1.0 x 10-4 M] 1.0 x 10-10 M= [H+] 14 = pH + pOH 14 = pH + 4 10 = pH

  24. Examples 4. A solution has a pH of 4. Calculate the pOH, [H+] and [OH-]. Is it acidic, basic, or neutral? 14= pH + pOH 14= 4 + pOH 10= pOH • Acidic since pH is 4

  25. Practice Problems: Classify each solution as acidic, basic or neutral. 1. [H+] = 1.0 x 10-10 M 2. [H+] = 0.001M 3. [OH-] = 1.0 x 10-7 M 4. [OH-] = 1.0 x 10-4 M Basic pH 10 1.0 x 10-3 acid pH 3 Neutral 14=pH+4 base pH 10

  26. White Board PracticeFill in the chart. 1.0 X 10 -8 1.0 X 10 -6 6 2 12 1.0 X 10 -2 1.0 X 10 -4 4 1.0 X 10 -10 11 3 1.0 X 10 -11

  27. Fill in the chart. 1.0 X 10 -8 1.0 X 10 -6 6 2 12 1.0 X 10 -2 1.0 X 10 -4 4 1.0 X 10 -10 11 3 1.0 X 10 -11 9 1.0 X 10 -5 1.0 X 10 -9 13 1 1.0 X 10 -13

  28. - + Strength or Concentration • Strong Acid/Base • Ionize completely in water • strong electrolyte Acids HCl HNO3 H2SO4 HBr HI HClO4 Bases NaOH KOH Ca(OH)2 Ba(OH)2

  29. - + Strength or Concentration • Weak Acid/Base • ionize partially in water • weak electrolyte Acids HF CH3COOH H3PO4 H2CO3 HCN Base NH3

  30. - - + + Strength or Concentration • How strong or weak an acid or base is, depends on its degree of ionization.

  31. pouvoir hydrogène (Fr.) “hydrogen power” pH Scale 14 0 7 INCREASING BASICITY INCREASING ACIDITY NEUTRAL pH is the negative logarithm of the hydrogen ion concentration pH = -log[H+]

  32. The pH Scale

  33. pH Scale pH of Common Substances

  34. pH formulas pH = -log[H+] pOH = -log[OH-] pH + pOH = 14

  35. Neutralization • Chemical reaction between an acid and a base. • Products are a salt (ionic compound) and water.

  36. Neutralization ACID + BASE  SALT + WATER HCl + NaOH  NaCl + H2O strong strong neutral HC2H3O2 + NaOH  NaC2H3O2 + H2O weak strong basic • Salts can be neutral, acidic, or basic. • Neutralization does not mean pH = 7.

  37. standard solution unknown solution Titration • Titration • Analytical method in which a standard solution is used to determine the concentration of an unknown solution.

  38. Titration cont. • Equivalence point (endpoint) • Point at which equal amounts of H+and OH- have been added. • Determined by… • indicator color change • dramatic change in pH

  39. Titration formula moles H+ = moles OH- MVn = MVn M: Molarity V: volume n: # of H+ ions in the acid or OH- ions in the base

  40. Titration example • 42.5 mL of 1.3M KOH are required to neutralize 50.0 mL of H2SO4. Find the molarity of H2SO4. H2SO4 M = ? V = 50.0 mL n = 2 KOH M = 1.3M V = 42.5 mL n = 1 MVn = MVn M(50.0mL)(2)=(1.3M)(42.5mL)(1) M= 55.25 100 M = 0.55M H2SO4

  41. Review of Acid and Base Definitions • Arrhenius Most specific/exclusive definition Created by Svante Arrhenius, Swedish Acid: compound that creates H+ in an aqueous solution HNO3 H+ + NO3- Base: compound that creates OH- in an aqueous solution NaOH  Na+ + OH-

  42. Review of Acid and Base Definitions • Bronsted-Lowry More general definition than Arrhenius definition Most commonly used definition Created by 2 scientists around the same time (1923) Acid: Molecule or ion that is a proton (H+) donor HCl + H2O  H3O+ + Cl- Base: Molecule or ion that is a proton (H+) acceptor NH3 + H2O  NH4+ + OH-

  43. Review of Acid and Base Definitions • Lewis Most general definition Defined by electrons and bonding rather than H+ Created by the same scientist who electron-dot diagrams are named after Acid: atom, ion, or molecule that accepts an electron pair to form a covalent bond NH3 + Ag+ [Ag(NH3)2]+ Base: atom, ion, or molecule that donates an electron pair to form a covalent bond BF3 + F-  BF4-

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