1 / 50

Reactions in Aqueous Solution

Reactions in Aqueous Solution. Chapter 4. Solution. Solvent. Solute. A solution is a homogenous mixture of 2 or more substances. The solute is(are) the substance(s) present in the smaller amount(s). The solvent is the substance present in the larger amount. H 2 O. Soft drink ( l ).

iyarbrough
Download Presentation

Reactions in Aqueous Solution

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Reactions in Aqueous Solution Chapter 4

  2. Solution Solvent Solute A solution is a homogenous mixture of 2 or more substances The solute is(are) the substance(s) present in the smaller amount(s) The solvent is the substance present in the larger amount H2O Soft drink (l) Sugar, CO2 Air (g) N2 O2, Ar, CH4 Pb Sn Soft Solder (s) 4.1

  3. nonelectrolyte weak electrolyte strong electrolyte An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity. A nonelectrolyte is a substance that, when dissolved, results in a solution that does not conduct electricity. 4.1

  4. What is the difference between these chemicals? • NaCl • HCl • CH3COOH • Sugar

  5. Ionic compounds that dissolve in water always produce solutions that are strong electrolytes. Molecular compounds that dissolve in water may produce nonelectrolytes, weak electrolytes (weak acids or bases), or strong electrolytes (strong acids).

  6. H2O NaCl (s)Na+ (aq) + Cl- (aq) CH3COOHCH3COO- (aq) + H+ (aq) Conduct electricity in solution? Cations (+) and Anions (-) Strong Electrolyte – 100% dissociation Dissociation means breaking up of the compound into cations and anions Example Weak Electrolyte – not completely dissociated Reversible reaction - Chemical equilibrium reversible

  7. H2O C6H12O6 (s) C6H12O6 (aq) Nonelectrolyte does not conduct electricity? No cations (+) and anions (-) in solution 4.1

  8. d- d+ H2O Hydration is the process in which an ion is surrounded by water molecules arranged in a specific manner.

  9. Water’s power as an ionizing solvent results from the distribution of its electrons and its overall shape. A polar molecule is one that has a + end and a - end. Water is one of the most polar molecules.

  10. The Solubility of Ionic Compounds in Water The solubility of ionic compounds in water depends upon the relative strengths of the electrical forces between ions in the ionic compound and the attractive forces between the ions and water molecules in the solvent. There is a tremendous range in the solubility of ionic compounds in water! The solubility of so called “insoluble” compounds may be several orders of magnitude less than ones that are called “soluble” in water, for example: Solubility of NaCl in water at 20oC = 365 g/L Solubility of MgCl2 in water at 20oC = 542.5 g/L Solubility of AlCl3 in water at 20oC = 699 g/L Solubility of PbCl2 in water at 20oC = 9.9 g/L Solubility of AgCl in water at 20oC = 0.009 g/L Solubility of CuCl in water at 20oC = 0.0062 g/L

  11. Chemical Reactions in Aqueous Solution

  12. Precipitation Reactions

  13. precipitate Pb(NO3)2(aq) + 2NaI (aq) PbI2(s) + 2NaNO3(aq) Pb2+ + 2NO3- + 2Na+ + 2I- PbI2 (s) + 2Na+ + 2NO3- Pb2+ + 2I- PbI2 (s) PbI2 Precipitation Reactions Precipitate – insoluble solid that separates from solution molecular equation ionic equation net ionic equation Na+ and NO3- are spectator ions 4.2

  14. AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq) Write the net ionic equation for the reaction of silver nitrate solution with sodium chloride solution. Ag+ + NO3- + Na+ + Cl- AgCl (s) + Na+ + NO3- Ag+ + Cl- AgCl (s) Writing Net Ionic Equations • Write the balanced molecular equation. • Determine precipitate from solubility rules. (page 97 of text) • Write the ionic equation showing the soluble strong electrolytes as ions. • Cancel the spectator ions on both sides of the ionic equation 4.2

  15. Molecular Equation: reactants and products are written as if they were molecular compounds. Total Ionic Equation: all soluble ionic substances (strong electrolytes) are dissociated into ions. Net Ionic Equation: Eliminate spectator ions. (These are ions which are unchanged in the chemical reaction.)

  16. Solubility Rules for Common Ionic Compounds In water at 250C (page 97 of text) 4.2

  17. These precipitation reactions are one type of a class of reactions sometimes called double displacement or metathesis reactions. The neutralization reactions we will study next are also in the class of reactions.

  18. Neutralization Reactions

  19. Acids Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain citric acid. React with certain metals to produce hydrogen gas. React with carbonates and bicarbonates to produce carbon dioxide gas Bases Have a bitter taste. Feel slippery. Many soaps contain bases. 4.3

  20. Arrhenius acid is a substance that produces H+ (H3O+) in water Arrhenius base is a substance that produces OH- in water 4.3

  21. A Brønsted acid must contain at least one ionizable proton! A Brønsted acid is a proton donor A Brønsted base is a proton acceptor base acid acid base 4.3

  22. HCl H+ + Cl- HNO3 H+ + NO3- CH3COOH H+ + CH3COO- H2SO4 H+ + HSO4- HSO4- H+ + SO42- HPO42- H+ + PO43- H2PO4- H+ + HPO42- H3PO4 H+ + H2PO4- Monoprotic acids Strong electrolyte, strong acid Strong electrolyte, strong acid Weak electrolyte, weak acid Diprotic acids Strong electrolyte, strong acid Weak electrolyte, weak acid Triprotic acids Weak electrolyte, weak acid Weak electrolyte, weak acid Weak electrolyte, weak acid 4.3

  23. Identify each as a Bronsted acid or base or both: a. HI b. CH3COO- c. H2PO4-

  24. acid + base salt + water HCl (aq) + NaOH (aq) NaCl (aq) + H2O H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O H+ + OH- H2O Neutralization Reaction 4.3

  25. Writing Balanced Equations forNeutralization Reactions Problem: Write balanced chemical equations (molecular, total ionic, and net ionic) for the following chemical reactions: a) barium hydroxide(aq) and hydroiodic acid(aq) b) lithium hydroxide(aq) and nitric acid(aq) c) sodium hydroxide(aq) and sulfuric acid(aq)

  26. Oxidation-Reduction Reactions

  27. Oxidation-Reduction (redox) reactions involve a transfer of electrons.oxidation: loss of electronsreduction: gain of electronsreducing agent: substance oxidizedoxidizing agent: substance reduced

  28. 2Mg (s) + O2 (g) 2MgO (s) 2Mg 2Mg2+ + 4e- O2 + 4e- 2O2- 2Mg + O2 + 4e- 2Mg2+ + 2O2- + 4e- 2Mg + O2 2MgO Oxidation-Reduction Reactions (electron transfer reactions) Oxidation half-reaction (lose e-) Reduction half-reaction (gain e-) 4.4

  29. Oxidation number: the charge an atom would have if the shared electrons were transferred completely to the atom with the greater attraction for the electrons.An increase in oxidation number indicates oxidation.A decrease in oxidation number indicates reduction.

  30. Oxidation number The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred. • Free elements (uncombined state) have an oxidation number of zero. Na, Be, K, Pb, H2, O2, P4 = 0 • In monatomic ions, the oxidation number is equal to the charge on the ion. Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2 • The oxidation number of oxygen isusually–2. In H2O2 and O22- it is –1. 4.4

  31. Oxidation numbers of all the elements in HCO3-, HNO3,Na3PO4, IF7, KIO3, K2Cr2O7, SO42- ? • The oxidation number of hydrogen is +1except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1. • Group 1 metals are +1, group 2 metals are +2 and fluorine is always –1. 6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion. 4.4

  32. Fig. 4.10

  33. A + B C S + O2 SO2 Oxidized? Reduced? Oxidizing agent? Reducing agent? C A + B 2KClO3 2KCl + 3O2 Types of Oxidation-Reduction Reactions Combination Reaction +4 -2 0 0 Decomposition Reaction +1 +5 -2 +1 -1 0 4.4

  34. CH4 + 2 O2 CO2 + 2 H2O Types of Oxidation-Reduction Reactions Combustion Reaction • when a substance reacts with oxygen (O2), usually with the release of heat and light to produce a flame. All combustion reactions are Redox reactions. Example: Using a Bunsen Burner to heat a beaker of water in lab. Cooking with natural gas.

  35. A + BC AC + B Sr + 2H2O Sr(OH)2 + H2 TiCl4 + 2Mg Ti + 2MgCl2 Cl2 + 2KBr 2KCl + Br2 Types of Oxidation-Reduction Reactions (Single) Displacement Reaction +1 +2 0 0 Hydrogen Displacement +4 0 0 +2 Metal Displacement -1 0 0 -1 Halogen Displacement Oxidized? Reduced? Oxidizing agent? Reducing agent? 4.4

  36. M + BC MC + B Ca + 2H2O Ca(OH)2 + H2 Pb + 2H2O Pb(OH)2 + H2 The Activity Series for Metals (page 113 of text) (Single) Displacement Reaction Reaction will occur only if M is above B. 4.4

  37. Ni (s) + 2AgNO3 (aq) Ni(NO3)2 (aq) + 2Ag (s) Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s) Cu2+ + 2e- Cu Nickel wire reacts with silver nitrate to form silver metal. What is the oxidizing agent in the reaction? Zn Zn2+ + 2e- Ni Ni2+ + 2e- Ag+ + 1e- Ag Zn is the reducing agent Zn is oxidized Cu2+is reduced Cu2+ is the oxidizing agent Ni is oxidized Ni is the reducing agent. Ag+is reduced Ag+ is the oxidizing agent 4.4

  38. 4.4

  39. For displacement reactions involving halogens the reactivity is F2> Cl2>Br2>I2 Cl2(aq) + NaBr(aq)  ? Cl2(aq) + NaF(aq)  ?

  40. Cl2 + 2OH- ClO- + Cl- + H2O Chlorine Chemistry Types of Oxidation-Reduction Reactions Disproportionation Reaction Element is simultaneously oxidized and reduced. +1 -1 0 4.4

  41. Ca2+ + CO32- CaCO3(s) NH3 + H+ NH4+ Classify the following reactions. Zn + 2HCl ZnCl2 + H2 Ca + F2 CaF2 Precipitation Acid-Base Redox (H2 Displacement) Redox (Combination) 4.4

  42. Solutions and Concentration

  43. What mass of KI is required to make 500. mL of a 2.80 M KI solution? moles of solute M = molarity = liters of solution volume KI moles KI grams KI 1 L 2.80 mol KI 166 g KI x x x 1000 mL 1 L soln 1 mol KI Solution Stoichiometry The concentration of a solution is the amount of solute present in a given quantity of solvent or solution. M KI M KI 500. mL = 232 g KI 4.5

  44. 4.5

  45. 500 mL of solution contains 0.730 moles of glucose. What is the molarity? What volume of 2.00 M CuSO4 will contain 10.0 g of CuSO4?

  46. Moles of solute before dilution (i) Moles of solute after dilution (f) = Dilution Add Solvent = MfVf MiVi Dilution is the procedure for preparing a less concentrated solution from a more concentrated solution. 4.5

  47. Volumetric Analysis - Titrations In a titration a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete. Equivalence point – the point at which the reaction is complete Indicator – substance that changes color at (or near) the equivalence point Slowly add base to unknown acid UNTIL the indicator changes color 4.7

  48. WRITE THE CHEMICAL EQUATION! What volume of a 1.420 M NaOH solution is Required to titrate 25.00 mL of a 4.50 M H2SO4 solution? H2SO4 + 2NaOH 2H2O + Na2SO4 M M rx volume acid moles acid moles base volume base base acid coef. 4.50 mol H2SO4 2 mol NaOH 1000 ml soln x x x 1000 mL soln 1 mol H2SO4 1.420 mol NaOH 25.00 mL = 158 mL 4.7

  49. How would you prepare 60.0 mL of 0.20 M HNO3 from a stock solution of 4.00 M HNO3? MfVf 0.20 M x 60.0 mL Vi = = 4.00 M Mi MiVi = MfVf Mi = 4.00 Vi = ? L Mf = 0.20 Vf = 0.06 L = 3.0 mL 3.0 mL of acid to 60 mL of solution 4.5

More Related