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Unit 3

Unit 3. Acids and Bases. Common Acids. Battery acid. Sulfuric Acid H 2 SO 4 Nitric Acid HNO 3 Phosphoric Acid H 3 PO 4 Hydrochloric Acid HCl Acetic Acid CH 3 COOH Carbonic Acid H 2 CO 3. Used to make fertilizers and explosives. Food flavoring. Stomach acid. Vinegar.

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Unit 3

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  1. Unit 3 Acids and Bases

  2. Common Acids Battery acid Sulfuric Acid H2SO4 Nitric Acid HNO3 Phosphoric Acid H3PO4 Hydrochloric Acid HCl Acetic Acid CH3COOH Carbonic Acid H2CO3 Used to make fertilizers and explosives Food flavoring Stomach acid Vinegar Carbonated water

  3. Sulfuric Acid, H2SO4 Sulfuric acid is the most commonly produced industrial chemical in the world. Uses: petroleum refining, metallurgy, manufacture of fertilizer, many industrial processes: metals, paper, paint, dyes, detergents Sulfuric acid is used in automobile batteries. H2SO4 “oil of vitriol”

  4. Nitric Acid, HNO3 Nitric acid stains proteins yellow (like your skin). Uses: make explosives, fertilizers, rubber, plastics, dyes, and pharmaceuticals. HNO3 “aqua fortis”

  5. Hydrochloric Acid, HCl The stomach produces HCl to aid in the digestion of food. Uses: For ‘pickling’ iron and steel. Pickling is the immersion of metals in acid solution to remove surface impurities. A dilute solution of HCl is called muriatic acid (available in many hardware stores). Muriatic acid is commonly used to adjust pH in swimming pools and in the cleaning of masonry. HCl(g) + H2O(l) HCl(aq) hydrogen chloride water hydrochloric acid

  6. Common Bases Name Formula Common Name Sodium hydroxide NaOH lye or caustic soda Potassium hydroxide KOH lye or caustic potash Magnesium hydroxide Mg(OH)2 milk of magnesia Calcium hydroxide Ca(OH) 2 slaked lime Ammonia water NH3 H2O household ammonia . NH4OH

  7. Naming Bases Bases are ionic compound. They follow the rules for naming ionic compounds and you have to use criss cross method to write the chemical formula. (The “+” ion first followed by the “-” ion.) Ex. 1 sodium hydroxide Na+ OH- NaOH

  8. Naming Bases Ex. 2 KOH K+ OH- Potassium hydroxide

  9. Naming Acids and writing formulas for acids Acids donate hydrogen ions (H+) and their chemical formulas have a hydrogen written at the beginning. Names of acid depend on the negative ion and its ending.

  10. Naming Acids and writing formulas for acids (continued) Anion ends with –ide : Ex. 1 acid formed with anion bromide if anion ends with –ide: To write name use prefix hydro + root of anion+ change –ide ending to –ic + word “acid” NAME: hydrobromic acid

  11. Naming Acids and writing formulas for acids (continued) Anion ends with –ate : Ex. 2 acid formed with anion carbonate if anion ends with –ate: To write name Use root of anion+ change –ate ending to –ic + word “acid” NAME: carbonic acid

  12. Naming Acids and writing formulas for acids (continued) Anion ends with –ite : Ex. 3 acid formed with anion nitrite if anion ends with –ite: To write name Use root of anion+ change –ite ending to –ous + word “acid” NAME: nitrous acid

  13. Naming Acids and writing formulas for acids (continued) All acids have as the positive ion the hydrogen ion (H+) and a negative ion that can be identified by the name of the acid. Ex. 4 hydrochloric acid if name of ion includes prefix hydro and has an -ic ending, then the anion has an -ide ending. From ion sheet determine which ion has root like the anion and an -ide ending. For the example is the chloride ion which formula is Cl- H+ Cl- HCl

  14. Naming Acids and writing formulas for acids (continued) Ex. 4 phosphoric acid if name of ion has an -ic ending, then the anion has an -ate ending. From ion sheet determine which ion has root like the anion and an -ate ending. For the example is the phosphate ion which formula is PO43- H+ PO43- H3 PO4

  15. Naming Acids and writing formulas for acids (continued) Ex. 4 phosphorous acid if name of ion has an -ous ending, then the anion has an -ite ending. From ion sheet determine which ion has root like the anion and an -ite ending. For the example is the phosphite ion which formula is PO33- H+ PO33- H3 PO3

  16. Acidsproperties • Conduct electricity (strong acids) • Change blue litmus to red • Have a sour taste • React with bases to neutralize their properties • React with active metals to liberate hydrogen • pH values <7

  17. Acids Acids are defined as: • Substances which ionize to form hydrogen ions (H+) in aqueous solution. (Arrhenius) • Substances that act as proton donors, H+ (Bronsted-Lowry) or as electron-pair acceptors (Lewis) • Examples HCl, H2SO4

  18. Bases properties • Base properties • Conduct electricity (strong bases) • Change red litmus to blue • Have a slippery feeling (like soap) • React with acids to neutralize their properties • pH values >7

  19. Bases Bases are defined as: • Substances which ionize to form hydroxide ions OH(-) in aqueous solution • Substances that act as proton receptors (Bronsted-Lowry) or as electron-pair donors (Lewis) • Examples: NH3OH, NaOH, CaCO3, NaHCO3(baking soda) AMMONIAcleaner

  20. pH 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 Acidic Neutral Basic lemon juice vinegar d water baking soda detergent ammonia pH Scale • The pH scale is used to measure how acidic or basic a liquid is. • pH measures the concentration of hydrogen ions (H+) and hydroxide ions (OH-). • The scale goes from 0 through 14. Distilled water is 7, so is called neutral.

  21. Indicators An indicator is a large organic molecule that works somewhat like a "color dye."

  22. Natural indicators • There are natural indicators for acids and bases, and we may find them in our kitchen or garden! • Red rose flowers • Bougainvillea flowers • Red cabbage • Blue berries

  23. Practice: identifying acid and bases 1) HCl + H2O  H3O+ + Cl– H+ HCl is the acid because it is donating a H+ to H2O H2O is the base because it is accepting a H+ from HCl Cl- is the conjugate base because it’s what’s left after the acid donates a H+ H3O+ is the conjugate acid because it’s what’s left after the base accepts a H+

  24. Practice: identifying acid and bases 2) CN- + H2O  OH- + HCN H+ H2O is the acid because it is donating a H+ to CN- CN- is the base because it is accepting a H+ from H2O OH- is the conjugate base because it’s what’s left after the acid donates a H+ HCNis the conjugate acid because it’s what’s left after the base accepts a H+

  25. Practice: identifying acid and bases H2O + I- 3) OH- + HI  H+ Donates a H+ acid The acid can donate a H+ and its chemical formula has a H at the beginning Classwork :handout conjugate acids and bases

  26. sulfuric acid: H2SO4 2 H1+ + SO42– -- nitric acid: HNO3 H1+ + NO31– -- Common Acids Strong Acids (dissociate ~100%) stomach acid; pickling: cleaning metals w/conc. HCl #1 chemical; (auto) battery acid explosives; fertilizer

  27. acetic acid: CH3COOH H1+ + CH3COO1– -- hydrofluoric acid: HF H1+ + F1– -- citric acid, H3C6H5O7 -- ascorbic acid, H2C6H6O6 -- lactic acid, CH3CHOHCOOH -- Common Acids (cont.) Weak Acids (dissociate very little) vinegar; naturally made by apples used to etch glass lemons or limes; sour candy vitamin C waste product of muscular exertion

  28. H2CO3: beverage carbonation rainwater dissolves limestone (CaCO3) in air CO2 + H2O H2CO3 H2CO3: cave formation H2CO3: natural acidity of lakes carbonic acid, H2CO3 -- carbonated beverages --

  29. 1+ Formation of Hydronium Ions 1+ 1+ + H+ H2O H3O+ hydrogen ion water hydronium ion (a proton)

  30. [H+] pH 10-14 14 10-13 13 10-12 12 10-11 11 10-10 10 10-9 9 10-8 8 10-7 7 10-6 6 10-5 5 10-4 4 10-3 3 10-2 2 10-1 1 100 0 1 M NaOH Ammonia (household cleaner) 7 Acid Base 0 14 Blood Pure water Milk Acidic Neutral Basic Vinegar Lemon juice Stomach acid 1 M HCl pH Scale Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 515

  31. pH of Common Substances water (pure) 7.0 vinegar 2.8 soil 5.5 gastric juice 1.6 carbonated beverage 3.0 drinking water 7.2 bread 5.5 1.0 M NaOH (lye) 14.0 blood 7.4 potato 5.8 orange 3.5 1.0 M HCl 0 milk of magnesia 10.5 urine 6.0 apple juice 3.8 detergents 8.0 - 9.0 bile 8.0 lemon juice 2.2 milk 6.4 tomato 4.2 ammonia 11.0 seawater 8.5 bleach 12.0 coffee 5.0 13 11 12 14 1 6 9 10 0 2 3 4 7 5 8 basic neutral acidic [H+] = [OH-] Timberlake, Chemistry 7th Edition, page 335

  32. Acid – Base Concentrations 10-1 pH = 3 pH = 11 OH- H3O+ pH = 7 10-7 concentration (moles/L) H3O+ OH- OH- H3O+ 10-14 [H3O+]<[OH-] [H3O+]>[OH-] [H3O+]=[OH-] acidic solution neutral solution basic solution Timberlake, Chemistry 7th Edition, page 332

  33. pH Calculations pH [H3O+] pH = -log[H3O+] [H3O+] = 10-pH pH + pOH = 14 [H3O+] [OH-] = 1 x10-14 pOH [OH-] pOH = -log[OH-] [OH-] = 10-pOH

  34. [ ] are used to represent concentration, usually molarity (mol/L) or (mol/dm3). • pH has no units.

  35. Example 1 What is the pH of a solution that contains 1.00x10-4 mol H3O+/dm3? Given [H3O+ ]=1.00x10-4 mol H3O+/dm3 Unknown: pH pH = -log[H3O+] pH = -log[1.00x10-4] pH = 4.00

  36. Example 2 What is the pH and pOH of a solution that contains 0.00350mol H3O+/dm3? Given [H3O+ ]=0.00350mol H3O+/dm3 Unknown: pH and pOH pH = -log[H3O+] pH = -log[0.00350] pH = 2.46 pH + pOH = 14 pOH = 14 – pH= 14 - 2.46 = 11.54

  37. Example 3 Given: pOH= 4.40 Unknown: [H3O+ ] and [OH- ] What is the [H3O+ ] and [OH- ] of a solution if pOH=4.40 [OH-] = 10-pOH [OH-] = 10-4.40 [OH-] = 3.98x10-5 M [H3O+] [OH-] = 1 x10-14 [H3O+] = 1 x10-14 = 1 x10-14 = 2.51x10-10M [OH-] 3.98x10-5

  38. classwork • P 184 #15,16,17

  39. Strengths of acids and bases • Strong acids and bases : ionize completely in aqueous solution. HCl (g) + H2O (l)  H3O+ (aq) + Cl- (aq) (100% ionized) Examples: HCl, HNO3, H2SO4, KOH, NaOH • Weak acids and bases: ionize slightly in aqueous solution. • CH3COOH (aq) + H2O (l) H3O+ (aq) + CH3COOH - (aq) (partially ionized) • Examples: HClO, H3PO4 , NH3

  40. Reactions between acids and bases When and acid and a base react with each other it is called neutralization reaction.

  41. Reactions between acids and bases General formula for acid base reaction: Acid + Base → Salt H2O + NOT JUST NaCl !! “Salt” means any ionic compound formed from an acid/base reaction

  42. Acids and bases reactions • The acids react with bases, forming salts ACID + BASE→ SALT + WATER HCl + NaOH→ NaCl+ H2O H-OH

  43. Titration • Equivalence point: when number of moles of hydrogen ions equals the number of moles of hydroxide ions. • Titration: adding a known amount of a solution of known concentration to determine the concentration of another solution. • Standard solution: solution of known concentration • End point: point at which the indicator changes color • Point of neutralization is the end point of the titration.

  44. Ex. 1 How many milliliters of 0.45M HCl will neutralize 25.0 mL of 1.00M KOH? • mL HCl= 25.0mL x 1.00M KOH 0.45 M HCl • mL HCl= 55.6mL HCl

  45. Ex. 2 What is the molarity of a NaOH solution if 20.0 mL of the solution is neutralized by 28.0mL of 1.00 M HCl? • M NaOH= 1.00 M HCl x 28.0 mL 20.0 mL • M NaOH = 1.40 M • Classwork: p 189 # 23 (a,c,d), 25, 26

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