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Elektrolit dan Nonelek

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Elektrolit dan Nonelek

  1. 1. CHEMISTRY 161 Chapter 4
  2. 2. CHEMICAL REACTIONS 2 HgO(s) → 2Hg(l) + O2(g) aq 1. properties of aqueous solutions 2. reactions in aqueous solutions a) precipitation reactions b) acid-base reactions (proton transfer) c) redox reactions (electron transfer)
  3. 3. 1.PROPERTIES OF AQUEOUS SOLUTIONS homogeneous mixture of two or more substances solvent solute substance in a large amount substance in a small amount N2 gas phase O2 (air) Ag solid phase Au (alloys) H2O liquid phase NaCl (sea water)
  4. 4. EXP1 iodine in ethyl alcohol (C2H5OH) does not conduct electricity (molecular solid) I2 EXP2 table salt in water (H2O) does conduct electricity (ionic solid) Na+Cl-
  5. 5. AQUEOUS SOLUTION solute water (H2O) solutes solution conducts electricity solution does not conduct electricity EXP3 electrolytes non-electrolytes
  6. 6. electrolytes non-electrolytes solution conducts electricity solution does not conduct electricity
  7. 7. non-electrolyte weak electrolyte strong electrolyte methanol sugar ethanol water ionic compounds (NaCl, KF) NaOH HCl H2SO4 CH3COOH HCOOH HF EXP5 dark medium bright
  8. 8. SOLUTION concentration
  9. 9. SOLUTION percentage concentration % = g [solute] / g solvent X 100 12 g of sodium chloride are solved in 150 g of water. Calculate the percentage concentration 8 %
  10. 10. SOLUTION solubility of a solute number of grams of solute that can dissolve in 100 grams of solvent at a given temperature 36.0 g NaCl can be dissolve in 100 g of water at 293 K
  11. 11. GAS PHASE SOLUTION Saturn solvent H2/He solute CH4, PH3
  12. 12. LIQUID SOLUTION Europa solvent H2O solute MgSO4
  13. 13. SOLID SOLUTION Triton solvent N2 solute CH4
  14. 14. methanol sugar ethanol water ELECTROLYTES ionic compounds (NaCl, KF) NaOH HCl H2SO4 CH3COOH HCOOH HF
  15. 15. migrating negative and positive charges Kohlrausch NaCl
  16. 16. DISSOCIATION ‘breaking apart’ NaCl (s) → Na+ (aq) + Cl- (aq) NaOH (s) → Na+ (aq) + OH- (aq) HCl (g) → H+ (aq) + Cl- (aq) Ca(NO3)2 (s) → Ca2+(aq) + 2 NO3 - (aq) strong electrolytes are fully dissociated EXP5 polyatomic ions do NOT dissociate
  17. 17. δ- O H H δ+ δ+
  18. 18. SOLVATION cations anions
  19. 19. SOLVATION non-electrolyte
  20. 20. NaCl (s) → Na+ (aq) + Cl- (aq) strong electrolytes are fully dissociated ←→ CH3COOH (aq) H+ (aq) + CH3COO- (aq) weak electrolytes are not fully dissociated reversible reaction (chemical equilibrium)
  21. 21. CHEMICAL REACTIONS 1.properties of aqueous solutions 2. reactions in aqueous solutions a) precipitation reactions b) acid-base reactions (proton transfer) c) redox reactions (electron transfer)
  22. 22. 2.1. PRECIPITATION REACTIONS solution 1 solution 2 solution 1 + solution 2
  23. 23. 2.1. PRECIPITATION REACTIONS formation of an insoluble product (precipitate) NaCl(aq) + AgNO3(aq) AgCl(s) + NaNO3(aq) EXP 6
  24. 24. insoluble compounds 1.M+ compounds (M = H, Li, Na, K, Rb, Cs, NH4) 2. A- compounds (A = NO3, HCO3, ClO3, Cl, Br, I) (AgX, PbX) 23. SO2- 4 (Ag, Ca, Sr, Ba, Hg, Pb) 4. CO2-, PO3-, CrO3 4 4 2-, S2- (Ag, Ca, Sr, Ba, Hg, Pb)
  25. 25. balanced molecular equation NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq) (table to determine which compound precipitates)
  26. 26. balanced ionic equation 1. NaCl(s) → Na+(aq) + Cl-(aq) 2. AgNO3(s) → Ag+(aq) + NO3 -(aq) 3. Na+(aq) + Cl-(aq) + Ag+(aq)+ NO3 -(aq) → AgCl(s) + Na+ (aq) + NO3 -(aq) spectator ions
  27. 27. Ba(NO3)2 (aq) + Na2SO4 (aq) Ba(NO3)2(aq) + Na3PO4(aq) Cs2CrO4(aq) + Pb(NO3)2(aq) 1. which compound falls out? 2. balanced molecular equation 3. balanced ionic equations 4. identify spectator ions
  28. 28. CHEMICAL REACTIONS 1.properties of aqueous solutions 2. reactions in aqueous solutions a) precipitation reactions b) acid-base reactions (proton transfer) c) redox reactions (electron transfer)
  29. 29. ACIDS AND BASES Arrhenius (1883) ACIDS HAc → H+ (aq) + Ac- (aq) ionization HCl (g) → H+ (aq) + Cl- (aq) BASES MOH → M+ (aq) + OH- (aq) NaOH (s) → Na+ (aq) + OH- (aq)
  30. 30. IDENTIFICATION Litmus Paper acid base red blue Säure Base EXP7
  31. 31. ACIDS AND BASES ACIDS and BASES NEUTRALIZE EACH OTHER HAc (aq) + MOH (aq) → MAc (aq) + H2O HCl (aq) + NaOH (aq) → NaCl (aq) + H2O acid + base salt + water
  32. 32. ACIDS AND BASES Na+ ≈ 10-10 m H+ ≈ 10-15 m
  33. 33. ACIDS AND BASES HCl (g) → H+ (aq) + Cl- (aq) H+(aq) + H2O H3O+(aq) HCl (g) + H2O → H3O+ (aq) + Cl- (aq) one step hydronium ion
  34. 34. (aq) (l) (aq) (aq) acid base hydronium ion
  35. 35. cation hydronium ion
  36. 36. PROPERTIES OF ACIDS 1. acids have a sour taste vinegar – acetic acid lemons – citric acid 2. acids react with some metals to form hydrogen 2 HCl(aq) + Mg(s) → MgCl2(aq) + H2(g) 3. acids react with carbonates to water and carbon dioxide 2 HCl(aq) + CaCO3(s) → CaCl2(aq) + [H2CO3] H2CO3 → H2O(l) + CO2(g) EXP8 EXP9 4. some acids are hygroscopic H2SO4 (conc)
  37. 37. BASES 1. bases have a bitter taste 2. bases feel slippery soap 3. aqueous bases and acids conduct electricity
  38. 38. EXAMPLES KOH(aq) and HF(aq) Mg(OH)2(aq) and HCl(aq) Ba(OH)2(aq) and H2SO4(aq) NaOH(aq) and H3PO4(aq) (stepwise)
  39. 39. Bronsted (1932) ACIDS proton donors HAc → H+ (aq) + Ac- (aq) BASES proton acceptor B + H+ (aq) → BH+ (aq)
  40. 40. electrolyte HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq) HNO3(aq) + H2O(l) → H3O+(aq) + NO3 weak electrolyte CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq) NH3(aq) + H2O(l) NH4 + + OH-strong -(aq) donor versus acceptor
  41. 41. CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq) NH3(aq) + H2O(l) NH4 +(aq)+ OH-(aq) H2O(l) + H2O(l) H3O+(aq) + OH-(aq) water can be either an acid or a base AUTO DISSOCIATION
  42. 42. monoprotic acids HF, HCl, HBr, HNO3, CH3COOH diprotic acid H2SO4 → H+(aq) + HSO4 -(aq) HSO4 -(aq) H+(aq) + SO4 2-(aq) triprotic acid H3PO4 H+(aq) + H2PO4 -(aq) -(aq) H+(aq) + HPO4 H2PO4 2-(aq) HPO4 2-(aq) H+(aq) + PO4 3-(aq) EXP10
  43. 43. CHEMICAL PROPOERTIES 1. Non-metal oxides react with water to form an acid (acetic anhydrides) g + ® aq g + ® aq g + ® aq SO ( ) H H SO ( ) sulfuric acid 3 2 O 2 4 N O ( ) H O 2HNO ( ) nitric acid 2 5 3 CO ( ) H O H CO ( ) carbonic acid 2 2 2 3 Cl2O7, SO2, Br2O5 22 + H2O
  44. 44. CHEMICAL PROPERTIES 2. Soluble metal oxides react with water to form a base (base anhydrides) s aq s aq + H2O + H2O CaO( ) H O Ca(OH) ( ) calcium hydroxide 2 2 Na O( ) H O 2NaOH( ) sodium hydroxide MgO, Al2O3 2 2 ® + ®
  45. 45. NAMING ACIDS AND BASES binary acids prefix hydro- the suffix –ic to the stem of the nonmetal name followed by the word acid g aq hydro ic acid g aq hydro ic acid HCl( ) hydrogen chloride HCl( ) chlor H S( ) hydrogen sulfide H S( ) sulfur 2 2
  46. 46. NAMING ACIDS AND BASES oxo acids acids contain hydrogen, oxygen, plus another element main group 5 HNO3 nitric acid HNO2 nitrous acid H3PO4 phosphoric acid H3PO3 phosphorous acid
  47. 47. main group 6 H2SO4 sulfuric acid H2SO3 sulfurous acid main group 7 HClO4 perchloric acid HClO3 chloric acid HClO2 chlorous acid HClO hypochlorous acid
  48. 48. Acids in the Solar System Venus H2SO4(g) Europa H2SO4(s)
  49. 49. Acids in the Interstellar Medium
  50. 50. Orion NH3, H2O, H2S CH3COOH HCOOH HF, HCl
  51. 51. CHEMICAL REACTIONS 1.properties of aqueous solutions 2. reactions in aqueous solutions a) precipitation reactions b) acid-base reactions (proton transfer) c) redox reactions (electron transfer)
  52. 52. KEY CONCEPTS 1. oxidation loss of electrons 2. reduction acceptance of electrons NUMBER OF ELECTRONS MUST BE CONSERVED
  53. 53. 1. oxidation EXAMPLE 2. reduction Na+Cl- Na ® Na+ + e Cl2 + 2 e ® 2 Cl- !!!balance electrons!!! CaO, Al2O3
  54. 54. substance that lost the electrons reduction agent substance that gained the electrons oxidizing agent oxidizing agent is reduced reducing agent is oxidized 2 Na + Cl2 ® 2 Na+Cl-
  55. 55. EXAMPLE 1 solid state reaction of potassium with sulfur to form potassium sulfide EXAMPLE 2 solid state reaction of iron with oxygen to form iron(III)oxide
  56. 56. OXIDATION NUMBER ionic compounds ↔ molecular compounds NaCl HF, H2 Na+Cl- ? electrons are fully transferred covalent bond charges an atom would have if electrons are transferred completely
  57. 57. EXAMPLE 1 HF H+ + F-molecular compound ionic compound H+ oxidation state +1 F- oxidation state -1
  58. 58. H2O 2 H+ + O2- EXAMPLE 2 molecular compound ionic compound H+ oxidation state +1 O2- oxidation state -2
  59. 59. H2 H+ + H-EXAMPLE 3 molecular compound ionic compound OXIDATION NUMBER OF FREE ELEMENTS IS ZERO
  60. 60. RULE 1 OXIDATION NUMBER OF FREE ELEMENTS IS ZERO H2, O2, F2, Cl2, K, Ca, P4, S8
  61. 61. RULE 2 monoatomic ions oxidation number equals the charge of the ion group I M+ group II M2+ group III M3+ (Tl: also +1) group VII (w/ metal) X-
  62. 62. RULE 3 oxidation number of hydrogen +1 in most compounds (H2O, HF, HCl, NH3) -1 binary compounds with metals (hydrides) (LiH, NaH, CaH2, AlH3)
  63. 63. RULE 4 oxidation number of oxygen -2 in most compounds (H2O, MgO, Al2O3) -1 in peroxide ion (O2 2-) (H2O2, K2O2, CaO2) -1/2 in superoxide ion (O2 -) (LiO2)
  64. 64. RULE 5 oxidation numbers of halogens F: -1 (KF) Cl, Br, I: -1 (halides) (NaCl, KBr) Cl, Br, I: positive oxidation numbers if combined with oxygen (ClO4 -)
  65. 65. RULE 6 charges of polyatomic molecules must be integers (NO3 -, SO4 2-) oxidation numbers do not have to be integers -1/2 in superoxide ion (O2 -)
  66. 66. MENUE 1.oxidation states of group I – III metals 2.oxidation state of hydrogen (+1, -1) 3. oxidation states of oxygen (-2, -1, -1/2, +1) 4.oxidation state of halogens 5.remaining atoms
  67. 67. oxidizing agents ????? OCl- Cl- EXP10
  68. 68. reducing agent 2 Na + 2 H2O ® H2 + 2 NaOH EXP11/12
  69. 69. NO+ NO NO 2 3- SO4 NO3 NO - 2 NO- - PO4 2- SO3 SO2 KO2 K2O BrO-KClO 4
  70. 70. REVISION 1.redox reactions 2. oxidation versus reduction 3. oxidation numbers versus charges 4. calculation of oxidation numbers
  71. 71. TYPES OF REDOX REACTIONS 1.combination reactions A + B → C 2. decomposition reactions C → A + B 3. displacement reactions A + BC → AC + B 4. disproportionation reactions
  72. 72. 1.combination reactions A + B → C two or more compounds combine to form a single product S8(s) + O2(g) → SO2(g) 1. oxidation numbers 2. balancing charges
  73. 73. MENUE 1.oxidation states of group I – III metals 2.oxidation state of hydrogen (+1, -1) 3. oxidation states of oxygen (-2, -1, -1/2, +1) 4.oxidation state of halogens 5.remaining atoms
  74. 74. 2. decomposition reactions C → A + B breakdown of one compound into two or more compounds HgO(s) → Hg(l) + O2(g) KClO3(s) → KCl(s) + O2(g) 1. oxidation numbers 2. balancing charges
  75. 75. 3. displacement reactions A + BC → AC + B an ion or atom in a compound is replaced by an ion or atom of another element 3.1. Hydrogen displacement 3.2. Metal displacement 3.3. Halogen displacement
  76. 76. 3.1. Hydrogen displacement group I and some group II metals (Ca, Sr, Ba) react with water to form hydrogen Na(s) + H2O(l) → NaOH + H2(g) less reactive metals form hydrogen and the oxide in water (group III, transition metals) Al(s) + H2O(l) → Al2O3(s) + H2(g)
  77. 77. 3.1. Hydrogen displacement even less reactive metals form hydrogen in acids Zn(s) + HCl(aq) → ZnCl2(aq) + H2(g) EXP12
  78. 78. activity series of metals Li K Ba Ca Na Mg Al Zn Cr Fe Cd Co Ni Sn Pb H Cu Hg Ag Pt Au displace H from water displace H from steam displace H from acids
  79. 79. Li K Ba Ca Na Mg Al Zn Cr Fe Cd Co Ni Sn Pb H Cu Hg Ag Pt Au likes to donate electrons does not like so much to donate electrons EXP13
  80. 80. 3.2. Metal displacement V2O5(s) + 5 Ca(s) → 2 V(s) + 5 CaO(s) TiCl4(g) + 2 Mg (l) → Ti(s) + 2 MgCl2(l)
  81. 81. 3.3. Halogen displacement F2 > Cl2 > Br2 > I2 reactivity (‘likes’ electrons) 0 +1 -1 +1 -1 0 Cl2(g) + 2 KBr(aq) → 2 KCl(aq) + Br2(l) Br2(g) + 2 KI(aq) → 2 KBr(aq) + I2(s)
  82. 82. 4. disproportionation reactions an element in one oxidation state is oxidized and reduced at the same time H2O2(aq) → 2 H2O(l) + O2(g) Cl2(g) + 2 OH-(aq) → ClO-(aq) + Cl-(aq) + H2O(l)
  83. 83. SUMMARY 1.combination reactions A + B → C 2. decomposition reactions C → A + B 3. displacement reactions A + BC → AC + B 4. disproportionation reactions
  84. 84. STOCHIOMETRY (CONCENTRATION) molar concentration Molarity (M) molarity (M) == moles of solute liters of solution
  85. 85. How many grams of AgNO3 are needed to prepare 250 mL of 0.0125 M AgNO3 solution? 3 0.531 g AgNO
  86. 86. How many mL of 0.124 M NaOH are required to react completely with 15.4 mL of 0.108 M H2SO4? 2 NaOH + H2SO4 Na2SO4 + 2H2O 26.8 mL NaOH
  87. 87. How many mL of 0.124 M NaOH are required to react completely with 20.1 mL of 0.2 M HCl? NaOH + HCl NaCl + H2O
  88. 88. How many grams of iron(II)sulfide have to react with hydrochloric acid to generate 12 g of hydrogen sulfide?
  89. 89. How many moles of BaSO4 will form if 20.0 mL of 0.600 M BaCl2 is mixed with 30.0 mL of 0.500 M MgSO4? BaCl2 + MgSO4 BaSO4 + MgCl2 This is a limiting reagent problem 4 0.0120 mol BaSO
  90. 90. How many ml of a 1.5 M HCl will be used to neutralize a 0.2 M Ba(OH)2 solution? How many ml of a 1.5 M HCl will be used to prepare 500 ml of a 0.1 M HCl? dil dil V X M = Vconcd X Mconcd
  91. 91. LIMITING REACTANT C2H4 + H2O C2H5OH EXP14
  92. 92. limiting reactant excess reactant
  93. 93. How many grams of NO can form when 30.0 g NH3 and 40.0 g O2 react according to 4 NH3 + 5 O2 4 NO + 6 H2O

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