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Chapter 4

Chapter 4. Solution Chemistry and the Hydrosphere. Earth: The Water Planet. About 70% of the Earth is covered with water, with 97% residing in oceans Earth’s early atmosphere may have been formed from the gases released by volcanic activity.

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Chapter 4

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  1. Chapter 4 Solution Chemistry and the Hydrosphere

  2. Earth: The Water Planet • About 70% of the Earth is covered with water, with 97% residing in oceans • Earth’s early atmosphere may have been formed from the gases released by volcanic activity. • As the Earth cooled the vapor in the atmosphere condensed and rain fell. This filled the depressions in the planet’s crust. We think…

  3. Temperature and Salinity (halinity) in the Oceans

  4. Salinity (halinity)

  5. Temperature and Salinity (halinity) vary with ocean depth

  6. Solution Terms • Solutions are homogeneous mixtures of two or more substances. • Solvent – the substance present in a solution in the greatest proportion (in number of moles). • Solute - the substance dissolved in the solvent.Aqueous Solution (aq) – a solution where water is the solvent.

  7. Solutions can be either Homogeneous or Heterogeneous (like oceans which are turbid). Solute + Solvent = Solution

  8. Solutions • Solutions are homogeneous mixtures of two or more substances. • The solvent is the substance in greatest quantity. • Solutes are the smaller quantity ingredients (usually) dissolved in the mixture.

  9.  104.5o   Aqueous SolutionsWater is the dissolving medium

  10. Some Properties of Water • Water is “bent” or V-shaped. • Water is a molecular compound. • Water is a polar molecule. • Hydration occurs when ionic compounds dissolve in water.

  11. Hydration of a solute Ion (Mg)

  12. How much of a solute dissolves? Concentration – ratio of the quantity of solute to either the mass or volume of the solution or solvent in which the solute is dissolved. Consequently, there are many different concentration terms depending on: 1. What is used to identify the quantity of solute (e.g. moles, mass, volume, etc.) 2. Whether the denominator of the ratio is the solvent or solution. 3. Whether mass or volume is used as the unit in the denominator.

  13. Concentration of Solutions • Molarity (M) = moles of solute per volume of solution in liters:

  14. Other Common Units of Concentration • Molality (m): moles solute/kg solvent • ppm: parts per million; mg solute/kg solution • ppb: mg solute/kg solution • ppt: ng solute/kg solution • % by weight: (grams of solute/total g solution) x 100% • Mole fraction: mole solute/total moles in solution

  15. Problem What is the molarity of an aqueous solution prepared by adding 36.5 g of barium chloride (208.233g/mol) to enough water to make 750.0 mL of solution?

  16. Determining the Number of Moles of Solute • # moles Solute = (Molarity)(# Liters of Solution) • n = M x volume (in Liters)

  17. How many grams of aluminum nitrate (212.996 g/mol) are required to make 500.0 mL of a 0.0525 M aqueous solution? Problem

  18. Problem What is the molarity of nitrate ions in a 0.0525 M solution of aluminum nitrate?

  19. Dilution of Concentrated Solutions • In dilution, a volume of stock solution is obtained and more solvent is added • The number of moles of solute is constant in a dilution. • # moles solute(stock) = Molaritys x volumes • The dilute solution uses the moles of solute with additional amounts of solvent. • # moles solute(dilute) = Molarityd x volumed Ms x Vs = Md x Vd

  20. Dilutions

  21. Practice Hydrochloric acid is obtained in 12.0 M stock solution. What volume of stock solution is required to make 500.0 mL of a 0.145 M dilute solution?

  22. Electrolytes Strong - conduct current efficiently Examples: Aqueous solutions of NaCl, HNO3,HCl

  23. ElectrolytesWeak- conduct only a small current (vinegar, tap water)

  24. Non Electrolyte A solution in which no ionization occurs. There is no conduction of electrical current. Examples: Aqueous solutions of sugar, ethelyne glycol

  25. Acid-Base Reactions • Bronsted-Lowry acids are proton (H+) donors. • Bronsted-Lowry bases are proton acceptors. • Free hydrogen ions don’t exist in water because they strongly associate with a water molecule to create a hydronium ion (H3O+) (a hydrated proton).

  26. Acid-Base Reactions • A neutralization reaction takes place when an acid reacts with a base and produces a solution of a salt and water. • A salt is made up of the cation characteristic of the base and the anion characteristic of the acid. • Example: HCl + NaOH ---> NaCl + H2O

  27. Strong Acids and Bases • A strong acid or strong base is completely ionized in aqueous solution. • HCl, HBr, HI, HNO3, HClO4 and H2SO4 are all strong acids. All other acids are assumed to be weak acids. • A weak acid or weak base only partially ionizes in aqueous solution. • Amphiprotic substances can behave as either a proton acceptor or a proton donor. Water is an example.

  28. Types of Equations • Molecular Equations have reactants and products written as undissociated (not ionized) molecules. HCl(aq) + NaOH(aq) ---> NaCl(aq) + H2O(l) • Overall Ionic Equations show all the species, both ionic and molecular present in aqueous solution for the reaction. H+ + Cl- + Na+ + OH- --> Na+ + Cl- + H2O

  29. Continued • Strong acids and strong bases are written as the corresponding ions in an overall ionic equation. • Net Ion Equations describe the actual chemical reaction occurring. H+ + OH- ----> H2O • The Na+ and Cl- ions are spectator ions in this reaction, because they are unchanged by the reaction (as is the solvent).

  30. Problem Write a balanced molecular equation and a net ionic equation for the following reactions: a. Solid magnesium hydroxide reacts with a solution of sulfuric acid.

  31. Problem Write a balanced molecular equation and a net ionic equation for the following reactions: a. Ammonia gas reacts with hydrogen chloride gas.

  32. Precipitation reactions: • A solid product (often colloidal) is formed from a reaction in solution. • The General Solubility Rules can be used to predict whether precipitates will form when mixing solutions of ionic compounds. • Rules are summarized on the following two slides. For example: AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) Notice that solid silver chloride is formed (a precipitate).

  33. GENERAL IONIC SOLUBILITY RULES (see Table 4.5, p. 149) Soluble Compounds:Exceptions: 1. All salts of (Na+), (K+) and (NH4+).  2. All (Cl‾ ), (Br‾ )and (I‾) [halide salts]  Halide salts of Ag+, Hg22+, Pb2+3. All (F‾) salts Fluoride salts of Mg2+, Ca2+, Sr2+, Ba2+, Pb2+4. All (NO3 ‾), (ClO3 ‾), (ClO4 ‾ ), Acetates of Ag+ and Hg22+ (C2H3O2 ‾) only moderately soluble 5.All sulfate salts (SO42 ‾) Sr2+, Ba2+, Pb2+, (Ca2+, Ag+ are moderately soluble)  

  34. GENERAL IONIC SOLUBILITY RULESPoorly Soluble Salts:Exceptions:6. All (CO32), (PO43), Na +, K +, NH4 + (CrO42 ), & (C2O42 )  7. All (S2 ) Group 1 & 2 cations and NH4 +8. All (OH ) & (O2 ) Group 1 & NH4 +, (Ca 2+, Sr2+ and Ba 2+ are moderately soluble)  

  35. Determining Whether a Precipitate will Form Does a precipitate form when sodium chloride is mixed with silver nitrate? If so what is the precipitate? CaCl2(aq) ---> Ca+2(aq) + 2Cl-(aq) Na2CO3 ---> 2Na+(aq) + CO3-2(aq)

  36. Net Ionic Equations • Soluble ionic compounds are called strong electrolytes and completely ionize in aqueous solution. • Write the balanced net ionic equation when sodium sulfate reacts with barium acetate.

  37. Types of Solutions • A saturated solution contains the maximum concentration of solute that can dissolve in it (for a given T, V and P). • A supersaturated solution contains more than the quantity of a solute that is predicted to be soluble in a given volume of solution at a given temperature.

  38. A Saturated Solution Example

  39. Supersaturated Solution Sodium acetate precipitates from a supersaturated solution.

  40. Problem What mass of barium sulfate (233.390g/mo) is produced when 100.0 mL of a 0.100 M solution of barium chloride is mixed with 100.0 mL of a 0.100 M solution of iron (III) sulfate?

  41. Rules for Assigning Oxidation States • The oxidation number of elements in a neutral molecule sum to zero or sum to charge of the ion in an ion. • Oxidation state of an atom in an element = 0 • Oxidation state of monatomic ion = charge • Fluorine = 1 in all compounds • Hydrogen = +1, Oxygen = 2 in most compounds (except in peroxides where oxygen = 1) • Unless combined with O, or F the halogens are -1.

  42. SO2 CrO42- NH3 ClO3- SF6 Cl2 Oxygen is -2 and Sulfur is +4 Oxygen is -2 and Chromium is +6 Practice Problems

  43. M X e- X- M+ Oxidation-Reduction Reactions Fe2O3(s) + Al(s)  Fe(l) + Al2O3(s) Oxidized Reduced Loses e- Gains e- Oxidation State Reduction State Increases Increases Reducing Agent Oxidizing Agent

  44. Fe2O3(s) + Al(s)  Fe(l) + Al2O3(s)

  45. Oxidation-Reduction Half Reactions • When copper wire is immersed in a solution of silver nitrate it is oxidized. • Cu ---> Cu2+ + 2 e- • Ag+ + e- ---> Ag • Silver ion is reduced.

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