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Chapter 4

Chapter 4. Aqueous Reactions and Solution Stoichiometry Pg 105. Aqueous Solutions. -Aqueous Solutions are solutions that have water as the dissolving medium. -Many reactions contain substances that have been dissolved in water, making them aqueous solutions. donkistry.tripod.com/Chem.jpg.

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Chapter 4

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  1. Chapter 4 Aqueous Reactions and Solution Stoichiometry Pg 105

  2. Aqueous Solutions -Aqueous Solutions are solutions that have water as the dissolving medium. -Many reactions contain substances that have been dissolved in water, making them aqueous solutions. donkistry.tripod.com/Chem.jpg

  3. 3 Main Major Chemical Reaction Types Involving Aqueous Solutions • Precipitation Reactions • Acid-Base Reactions • Redox Reactions

  4. 4.1 General Properties of Aqueous Solutions • Solutions are homogeneous mixtures • Usually has more solvent than solute. • Solute is the substance being dissolved in the solvent

  5. Electrolytic Properties • Pure water is a bad conductor • The presence of Ions in water makes it into a good conductor • Aqueous solution that conduct electricity such as NaCl(aq) or other ionic compounds are electrolyte. • Solutions that do not form ions like sucrose and other molecular compounds are nonelectrolytes.

  6. Ionic Compounds in Water • Ionic compounds dissolve in water dissociating into component ions (ex. NaCl -> Na+&Cl-) • The polar nature of water makes it a very effective solvent • The polarity helps prevent anions and cations from rejoining.

  7. Molecular Compounds in Water • Structure usually remains unchanged, they usually do not form ions • Acids and a few other compounds like ammonia react with water forming ions making an electrolyte. • Ex. HCl make H+ and Cl- ions

  8. Strong & Weak Electrolytes • Strong Electrolytes = Most ionic compounds and a few molecular compounds. • Weak Electrolytes = Molecular compounds that produce few ions when dissolved • If the chemical reaction goes both ways, breaking into ions, and recombining, than the substance is a weak electrolyte.

  9. HCl(aq) --> H+ + Cl-One arrow means strong electrolyte HC2H3O29(aq) <--> H+ + C2H3O2- Double arrow means weak electrolyte

  10. 4.2 Precipitation Reactions • Precipitation Reaction = Reactions that result in the formation of an insoluble product. • Precipitate = Insoluble solid formed by a reaction in a solution www.wiley.com/.../resources/ch04/mercury_ppt.jpg

  11. Solubility Guidelines for Ionic Compounds • Solubility = Amount of substance that can be dissolved in given amount of solvent • If less than .01 mol dissolves in a liter, substance is insoluble. In these substances intermolecular attraction is stronger than the waters polarity. • Table 4.1 pg 111 (Solubility Guidelines for Common Ionic Compounds in Water) • All ionic compounds with 1A elements or ammonia ions are soluble in water.

  12. Is Sodium Carbonate Soluble(Na2CO3) • Yes. Carbonate is usually insoluble, but when paired with a 1A element, Sodium, the compound becomes soluble. www.csudh.edu/oliver/demos/carbnate/BL1.jpg

  13. Exchange (Metathesis) Reactions • Exchange or Metathesis Reaction = AX+BY --> AY+BX • Precipitation and Acid Base Reactions conform to this pattern

  14. What precipitate forms when BaCl2 and K2SO4 are mixed? • BaSO4, SO42- is soluble but Ba2+ is not www.galleries.com/.../Sulfates/BARITE/barite.jpg

  15. Ionic Equations • Molecular Equation = complete chemical formulas of reactants and products • Complete Ionic Equation = All Soluble strong Electrolytes are shown as ions • Spectator ions = ions that are present in the same form on both product and reactant side. These are dropped out to form a Net Ionic Equation.

  16. Steps to Write a Net Ionic Equation • Write a balanced Molecular Equation • Rewrite to show ions that are formed during dissociation or ionization, only the strong electrolytes are written in ionic form • Cancel spectator ions on both sides

  17. Write the net ionic equation for the mixing of CaCl2 and Na2CO3 • CaCl2(aq) + Na2CO3(aq) --> CaCO3(S) + 2NaCl(aq) • Ca2++ 2Cl- + 2Na+ + CO32--->CaCO3(s) + 2Na+ + 2Cl- • Ca2+(aq) + CO32-(aq)-->CaCO3(s)

  18. Acid-Base Reactions • Acids and Bases are common Electrolytes • Are some of the most common compounds we encounter

  19. Acids • Substances that ionize in aqueous solutions upping H+ concentration • Protic refers to amount of H+ ions ionizing. Monoprotic = 1, Diprotic = 2. • Diprotic Acid ionization occurs in two steps, One hydrogen is separated at a time.

  20. Bases • Substances that accept H+ ions, or increases OH- concentration. • Does not need to have an OH- ion, if accepts H+ like NH3 (ammonia is a weak electrolyte)

  21. Strong and Weak Acids and Bases • Strong Acids and Bases are strong electrolytes that completely ionize in solutions • Weak Acids and Bases are electrolytes that partly ionize in solutions • Table 4.2 pg 115 (Common Strong Acids and Bases)

  22. Identifying Strong and Weak Electrolytes • Is the compound ionic, yes -> probably strong electrolyte • Not ionic, is it an acid • Yes, is an acid, if strong, is a strong electrolyte if weak, is a weak electrolyte. • Not an acid, is it NH3 or another molecular base, yes -> weak base, no -> probably nonelectrolyte

  23. Classify HNO3 as a strong, weak, or non Electrolyte • Strong • HNO3 is a strong acid making it a strong electrolyte. mls.jpl.nasa.gov/images/HNO3.jpg

  24. Neutralization Reactions and Salts • When acid and base react together, it is a neutralization reaction. • These reactions form a salt and a water. • Salt = any ionic compound whose cation comes form a base and anion comes from an acid.

  25. Write the net ionic equation for HC2H3O2 and Ba(OH)2 HC2H3O2(aq)+OH-(aq)--> H2O(l) + C2H3O2-(aq)

  26. Acid-Base Reaction with Gas Formation • The sulfide ion and carbonate ion react with acids to form gases • 2HCl(aq) + Na2S(aq) --> H2S(g)+2NaCl(aq) • HCl(aq)+NaHCO3(aq)-->NaCl(aq) + H2O(l)+CO2(g)

  27. 4.4 Oxidation - Reduction Reactions http://k53.pbase.com/o6/13/615013/1/87094194.ExyClUcI.070929_LF240101w.jpg

  28. Oxidation and Reduction • Metals undergoing erosion are losing electrons and forming cations • Loss of electrons is known as oxidization • The gain of electrons by a substance is called reduction

  29. Oxidation Numbers • Oxidization number is the actual charge of the of the atom. • In elemental form the Oxidization number is 0 • Oxidization of monatomic ions equals the charge • Nonmetals are usually negative, Oxygen usually is -2, Hydrogen is +1, Fluorine is -1. • Sum of oxidation numbers in neutral compound is 0 or equal to the charge in a polyatomic ion.

  30. Determine the oxidation stat of sulfur in H2S • -2, Hydrogen is always +1 2H = +2 so S = -2 so that sum of oxidation numbers = 0

  31. Oxidation of Metals by Acids and Salts • Displacement Reactions = ion is solution is displaced or replaced through the oxidation of an element. • A+BX-->AX+B

  32. Write the net ionic equation for the reaction of aluminum and hydrobromic acid • 2Al(S) + 6H+(aq)+6Br--->2Al3+(aq)+6Br-(aq)+3H2(g) • 2Al(s)+6H+(aq)-->2Al3++3H2(g)

  33. The Activity Series • Table 4.5 pg 124 Activity Series of Metals in Aqueous Solution • Is a table of metals arranged in order of decreasing ease of oxidation. • Alkali and Alkaline Earth Metals are at the top. • Metals on the list can be oxidized by any of the ions below them.

  34. Can Pb(NO3)2 can oxidize Zn, Cu, or Fe • Zn and Fe • Refer to table 4.5

  35. 4.5 Concentrations Of Solutions • Concentration = Amount of solute dissolved in a given quantity of solvent or solution.

  36. Molarity • Molarity (M) = (moles of solute)/(Volume of Solution in Liters)

  37. Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulfate in enough water to form 125ml of solution • 23.4g*(1mol Na2SO4/142g Na2SO4) = .165 mols • 125ml*(1L/1000ml)=.125 • .165mols Na2SO4/.125L = 1.32M

  38. Dilution • Adding water to lower the concentration is called dilution • Mi*Vi=Mf*Vf • i = initial f = final M = Molarity V= Volume

  39. How many mL of 3.0 M H2SO4 are required to make 450 mL of .10 M H2SO4 ? • Vi = (MfVf)/Mi • ((.10M)(450mL))/3.0M = 15 mL

  40. 4.6 Stoichiometry and Chemical Analysis www.sciencebuddies.org/.../Chem_img030.jpg

  41. Titrations • Second solution of known concentration is called the standard solution • Combining the standard solution with a solution of unknown concentration to get a chemical reaction is called titration • Equivalence point is where equivalent quantities have been brought together indicators change the color helping us to find this point. • If molar ratio is 1 to 1 you may use the dilution equation. • If not, convert standard solution to mols, then use molar ratio to give you the mols of the unknown,then convert to grams.

  42. How many grams of chloride ion are in the sample of the water if 20.2 mL of .1 M Ag is required to react with all the chloride in the sample? • (20.2 mL solution) * (1L/1000mL solution) * (.1mol Ag+/L solution) = 2.02 * 10-3 • 1mol Ag+ : 1mol Cl- • (2.02*10-3 mol Cl-) * (35.5g Cl-/1 mol Cl-) = 7.17 * 10-2 g Cl-

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