Properties of Solution

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Properties of Solution

  1. 1. CHEMISTRY PROPERTIES OF SOLUTION __RaNNY ROLINda RUSMaN__
  2. 2. homogeneous mixture of two or more substances solute solvent substance in a large amount substance in a small amount N2 gas phase (air) O2 Ag solid phase (alloys) Au H2O liquid phase (sea water) NaCl
  3. 3. 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-
  4. 4. solute water (H2O) solutes solution conducts electricity solution does not conduct electricity EXP3 electrolytes non-electrolytes
  5. 5. electrolytes non-electrolytes solution conducts electricity solution does not conduct electricity
  6. 6. non-electrolyte methanol sugar ethanol water weak electrolyte CH3COOH HCOOH HF EXP5 strong electrolyte ionic compounds (NaCl, KF) NaOH HCl H2SO4 dark medium bright
  7. 7. concentration
  8. 8. 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%
  9. 9. 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
  10. 10. Saturn solvent H2/He solute CH4, PH3
  11. 11. Europa solvent H 2O solute MgSO4
  12. 12. Triton solvent N2 solute CH4
  13. 13. ELECTROLYTES methanol sugar ethanol water ionic compounds CH3COOH (NaCl, KF) HCOOH NaOH HF HCl H2SO4
  14. 14. migrating negative and positive charges Kohlrausch NaCl
  15. 15. 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 polyatomic ions do NOT dissociate EXP5
  16. 16. δO δ+ H H δ+
  17. 17. SOLVATION cations anions
  18. 18. SOLVATION non-electrolyte
  19. 19. 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)
  20. 20. 1.properties of aqueous solutions 2. reactions in aqueous solutions a) precipitation reactions b) acid-base reactions (proton transfer) c) redox reactions (electron transfer)
  21. 21. 2.1. PRECIPITATION REACTIONS solution 1 solution 2 solution 1 + solution 2
  22. 22. 2.1. PRECIPITATION REACTIONS formation of an insoluble product (precipitate) NaCl(aq) + AgNO3(aq) AgCl(s) + NaNO3(aq) EXP 6
  23. 23. insoluble compounds 1.M+ compounds (M = H, Li, Na, K, Rb, Cs, NH4) 2. A- compounds (A = NO3, HCO3, ClO3, Cl, Br, I) (AgX, PbX2) 3. SO42(Ag, Ca, Sr, Ba, Hg, Pb) 4. CO32-, PO43-, CrO42-, S2-
  24. 24. balanced molecular equation NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq) (table to determine which compound precipitates)
  25. 25. 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
  26. 26. 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
  27. 27. 1.properties of aqueous solutions 2. reactions in aqueous solutions a) precipitation reactions b) acid-base reactions (proton transfer) c) redox reactions (electron transfer)
  28. 28. ACIDS HAc → H+ (aq) + Ac- (aq) ionization HCl (g) → H+ (aq) + Cl- (aq) BASES MOH → M+ (aq) + OH- (aq) Arrhenius (1883) NaOH (s) → Na+ (aq) + OH- (aq)
  29. 29. Litmus Paper acid red Säure base blue Base EXP7
  30. 30. ACIDS BASES and NEUTRALIZE EACH OTHER HAc (aq) + MOH (aq) → MAc (aq) + H2O HCl (aq) + NaOH (aq) → NaCl (aq) + H2O acid + base salt + water
  31. 31. Na+ H+ ≈ 10-10 m ≈ 10-15 m
  32. 32. 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
  33. 33. (aq) acid (l) (aq) base hydronium ion (aq)
  34. 34. cation hydronium ion
  35. 35. 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) EXP8 3. acids react with carbonates to water and carbon dioxide 2 HCl(aq) + CaCO3(s) → CaCl2(aq) + [H2CO3] H2CO3 → H2O(l) + CO2(g) 4. some acids are hygroscopic H2SO4 (conc) EXP9
  36. 36. BASES 1. bases have a bitter taste 2. bases feel slippery soap 3. aqueous bases and acids conduct electricity
  37. 37. 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)
  38. 38. ACIDS proton donors HAc → H+ (aq) + Ac- (aq) BASES proton acceptor Bronsted (1932) B + H+ (aq) → BH+ (aq)
  39. 39. strong electrolyte HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq) HNO3(aq) + H2O(l) → H3O+(aq) + NO3-(aq) weak electrolyte CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq) NH3(aq) + H2O(l) NH4+ + OH- donor versus acceptor
  40. 40. 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
  41. 41. monoprotic acids HF, HCl, HBr, HNO3, CH3COOH diprotic acid H2SO4 → H+(aq) + HSO4-(aq) HSO4-(aq) H+(aq) + SO42-(aq) triprotic acid H3PO4 H+(aq) + H2PO4-(aq) H2PO4-(aq) H+(aq) + HPO42-(aq) HPO42-(aq) H+(aq) + PO43-(aq) EXP10
  42. 42. CHEMICAL PROPOERTIES 1. Non-metal oxides react with water to form an acid (acetic anhydrides) SO3 (g ) + H 2O → H 2SO 4 (aq) sulfuric acid N 2 O5 (g ) + H 2O → 2HNO3 (aq) nitric acid + H+H CO 2 (g )2O 2O → H 2CO 3 (aq) carbonic acid + H 2O + H 2O Cl2O7, SO2, Br2O5
  43. 43. CHEMICAL PROPERTIES 2. Soluble metal oxides react with water to form a base (base anhydrides) CaO(s ) + H 22O→ Ca(OH) 2 (aq) calcium hydroxide + HO Na 2 O(s ) + H 2O → 2NaOH(aq) + H 2O MgO, Al2O3 sodium hydroxide
  44. 44. NAMING ACIDS AND BASES binary acids prefix hydrothe suffix –ic to the stem of the nonmetal name followed by the word acid HCl(g ) hydrogen chloride HCl(aq) hydrochloric acid H 2S(g ) hydrogen sulfide H 2S(aq) hydrosulfuric acid
  45. 45. NAMING ACIDS AND BASES oxo acids acids contain hydrogen, oxygen, plus another element main group 5 HNO3 HNO2 nitric acid nitrous acid H3PO4 H3PO3 phosphoric acid phosphorous acid
  46. 46. main group 6 H2SO4 H2SO3 sulfuric acid sulfurous acid main group 7 HClO4 HClO3 HClO2 HClO perchloric acid chloric acid chlorous acid hypochlorous acid
  47. 47. Europa Venus H2SO4(s) H2SO4(g)
  48. 48. NH3, H2O, H2S CH3COOH HCOOH HF, HCl Orion
  49. 49. 1.properties of aqueous solutions 2. reactions in aqueous solutions a) precipitation reactions b) acid-base reactions (proton transfer) c) redox reactions (electron transfer)
  50. 50. 1. oxidation loss of electrons 2. reduction acceptance of electrons NUMBER OF ELECTRONS MUST BE CONSERVED
  51. 51. Na+Cl- 1. oxidation Na → Na+ + e 2. reduction Cl2 + 2 e → 2 Cl- !!!balance electrons!!! CaO, Al2O3
  52. 52. 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-
  53. 53. 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
  54. 54. OXIDATION NUMBER ionic compounds ↔ molecular compounds NaCl Na+Cl- HF, H2 ? electrons are fully transferred covalent bond charges an atom would have if electrons are transferred completely
  55. 55. H+ + F- HF molecular compound ionic compound H+ oxidation state +1 F- oxidation state -1
  56. 56. H2O 2 H+ + O2- molecular compound ionic compound H+ oxidation state +1 O2- oxidation state -2
  57. 57. H2 molecular compound H+ + Hionic compound OXIDATION NUMBER OF FREE ELEMENTS IS ZERO
  58. 58. OXIDATION NUMBER OF FREE ELEMENTS IS ZERO H2, O2, F2, Cl2, K, Ca, P4, S8
  59. 59. 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-
  60. 60. RULE 3 oxidation number of hydrogen +1 in most compounds (H2O, HF, HCl, NH3) -1 binary compounds with metals (hydrides) (LiH, NaH, CaH2, AlH3)
  61. 61. RULE 4 oxidation number of oxygen -2 in most compounds (H2O, MgO, Al2O3) -1 in peroxide ion (O22-) (H2O2, K2O2, CaO2) -1/2 in superoxide ion (O2-) (LiO2)
  62. 62. 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-)
  63. 63. RULE 6 charges of polyatomic molecules must be integers (NO3-, SO42-) oxidation numbers do not have to be integers -1/2 in superoxide ion (O2-)
  64. 64. 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
  65. 65. oxidizing agents OCl- ????? EXP10 Cl-
  66. 66. reducing agent 2 Na + 2 H2O → H2 + 2 NaOH EXP11/12
  67. 67. K2O PO 34 NO+ SO42- KO2 SO3 NO3NO2 NO KClO4 BrO- SO2 NO2- NO-
  68. 68. 1.redox reactions 2. oxidation versus reduction 3. oxidation numbers versus charges 4. calculation of oxidation numbers
  69. 69. 1.combination reactions A+B→C 2. decomposition reactions C→A+B 3. displacement reactions A + BC → AC + B 4. disproportionation reactions
  70. 70. 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
  71. 71. 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
  72. 72. 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
  73. 73. 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
  74. 74. 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)
  75. 75. even less reactive metals form hydrogen in acids Zn(s) + HCl(aq) → ZnCl2(aq) + H2(g) EXP12
  76. 76. 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
  77. 77. Li K Ba Ca Na Mg Al Zn Cr Fe Cd Co Ni Sn Pb H Cu Hg Ag Pt Au does not like so much to donate electrons likes to donate electrons EXP13
  78. 78. 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)
  79. 79. 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)
  80. 80. 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)
  81. 81. 1.combination reactions A+B→C 2. decomposition reactions C→A+B 3. displacement reactions A + BC → AC + B 4. disproportionation reactions
  82. 82. molar concentration Molarity (M) molarity (M) = moles of solute = liters of solution
  83. 83. How many grams of AgNO3 are needed to prepare 250 mL of 0.0125 M AgNO3 solution? 0.531 g AgNO3
  84. 84. 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
  85. 85. 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
  86. 86. How many grams of iron(II)sulfide have to react with hydrochloric acid to generate 12 g of hydrogen sulfide?
  87. 87. 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 0.0120 mol BaSO 4
  88. 88. 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? Vdil X M dil = Vconcd X M concd
  89. 89. LIMITING REACTANT EXP14 C2H4 + H2O C2H5OH
  90. 90. limiting reactant excess reactant
  91. 91. 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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