Adv chem chapt 4

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  • 1.Advanced Chemistry Types of Chemical Reactions and Solution Stoichiometry

2. The Common Solvent

  • Water

3. Water as a Solvent

  • Bent (V-shaped) molecule with 105 bond angle.
  • Covalent bonds electron sharing between oxygen and hydrogen.
  • Unequal sharing of electrons due to electronegativity differences cause water to be apolar molecule .

4. Water as a Solvent

  • Partial charges () form on the water molecule due to theunequal sharing of electrons.
  • Waters polarity causes it to be a solvent for ionic compounds.

5. Polar Water Molecules Interact with the Positive and Negative Ions of a Salt Assisting in the Dissolving Process 6. Solubility

  • Solubility of ionic substances in water varies greatly.
  • Solubility depends on the relative attractions of the ions and water molecules for each other.
  • Once dissolved, an ionic compound becomeshydrated , therefore ions disperse independent of one another.

7. Solubility

  • Water will dissolve non-ionic substances depending upon their structure.
  • If a polar bond exists within the structure, the molecule can be subject to being water soluble.
  • In general, polar substances dissolve in polar solvents and nonpolar substances dissolve in nonpolar solvents.

8. An Ethanol Molecule Contains a Polar O-H Bond Similar to Those in the Water Molecule 9. The Polar Water Molecule Interacts Strongly with the Polar-O-H bond in Ethanol 10. Strong and Weak Electrolytes

  • Nature of Aqueous Solutions

11. Strong and Weak Electrolytes

  • A solution in which a substance is dissolved in water; the substance is thesoluteand thesolventis the water.
  • A common property for a solution iselectrical conductivity.
    • Its ability to conduct an electric current .

12. Strong and Weak Electrolytes

  • Pure water is not an electrical conductor.
  • Strong electrolytesconduct current very efficiently.
  • Weak electrolytesconduct only a small current.
  • Nonelectrolytesdo not allow current to flow.

13. Electrical Conductivity of Aqueous Solutions 14. Svante Arrhenius

  • 1859-1927
  • Studied the nature of solutions and theorized that conductivity of solutions arose from the presence of ions.
  • Proved that the strength of conductivity is directly related to the number of ions present in solution.

15. Strong Electrolytes

  • Strong electrolytesare substances that are completely ionized when they are dissolved in water.
    • Soluble salts
    • Strong acids
    • Strong bases

16. NaCl Dissolves 17. Acids

  • Arrhenius discovered in his studies of solutions that when acids were dissolved in water they behaved as strong electrolytes.
    • This result was directly related to an acids ability to ionize in water.
  • AnAcidis a substance that produces H +ions when it is dissolved in water.

18. HCl is Completely Ionized 19. Acids AnAcidis a substance that produces H +ions when it is dissolved in water. 20. Acids In conductivity studies, virtually every molecule ionizes. Therefore, strong electrolytes are strong acids. 21. Strong Acids

  • Sulfuric acid, nitric acid and hydrochloric acid are aqueous solutions and should be written in chemical equations as such.
  • Astrong acidis one that completely dissociates into its ions.In aqueous solutions, the HCl molecule does not exist.
  • Sulfuric acid can produce two H +ions per molecule. Only the first H +ion completely dissociates.The anion HSO 4 -remains partially intact.

22. Strong Bases

  • Bases are soluble ionic compounds containing the hydroxide ion (OH - ).
  • Strong basesare strong electrolytes and these compounds ionize completely in water.

23. Strong Bases 24. An Aqueous Solution of Sodium Hydroxide 25. Weak Electrolytes

  • Weak electrolytesare substances that exhibit a small degree of ionization in water.
    • They produce relatively few ions when dissolved in water
    • Most common weak electrolytes are weak acids and weak bases.

26. Weak Acids

  • Formulas for acids are often written with the acidic hydrogen atom or atoms (the hydrogen atoms that will produce H +ions in solution) listed first.If any nonacidic hydrogens are present they are written later in the formula.

27. Weak Acids

  • In Acetic acid, only 1% of its molecules ionize.
  • The double arrow indicates that the reaction can occur in either direction.
  • Acetic acid is aweak electrolyteand therefore aweak acidbecause it dissociates (ionizes) only to a slight extent in aqueous solutions.

28. Acetic Acid (HC 2 H 3 O 2 ) 29. Weak Bases

  • The most common weak base is NH 3 .
  • In an aqueous solution, ammonia results in a basic solution.

30. The Reaction of NH 3in Water 31. Nonelectrolytes

  • Nonelectrolytes are substances that dissolve in water but do not produce any ions.

32. The Composition of Solutions

  • Concentration:

33. Solutions

  • Most chemical reactions take place in the environment of solutions.In order to perform stoichiometric calculations in solutions, one must know two things.
    • The nature of the reaction; which depends on the exact forms the chemicals take when dissolved.
    • The amounts of the chemicals present in the solutions, usually expressed as concentrations.

34. Concentration

  • Molarity (M) is moles of solute per volume of solution in liters:
  • Example: 1.0M= 1.0 molar = 1.0moles solute/1liter of solution

35. Example

  • Calculate the molarity of a solution prepared by dissolving 11.5g of solid NaOH in enough water to make 1.50L of solution.

.192 M NaOH 36. Example

  • Calculate the molarity of a solution prepared by dissolving 1.56g of gaseous HCl in enough water to make 26.8 ml of solution.

1.60M HCl 37. Example

  • Give the concentration of each type of ion in 0.50M Co(NO 3 ) 2 .

Co 2+= 0.50 M Co 2+ NO 3 -= 1.0 M NO 3 - 38. Example

  • Calculate the number of moles of Cl -ions in 1.75L of 1.0 x 10 -3M ZnCl 2 .

3.5 x 10 -3mol Cl - 39. Dilution

  • The process of changing the molarity of a solution from a more concentrated solution to a lesser concentrated solution.
    • Moles of solute after dilution = moles of solute before dilution.
    • M x V= moles
    • M 1 V 1 =M 2 V 2

40. Example

  • What volume of 16M sulfuric acid must be used to prepare 1.5L of 0.10 M H 2 SO 4solution?

9.4 ml solution 41. Steps Involved in the Preparation of a Standard Aqueous Solution 42. Types of Chemical Reactions 43. Types of Solution Reactions

  • Most solution reactions can be put into three types of reactions:
    • Precipitation Reactions
    • Acid-Base Reactions
    • Oxidation-Reduction Reactions

44. Precipitation Reactions

  • When two solutions are mixed, an insoluble substance sometimes forms; that is, a solid forms and separates from the solution.
    • The solid that forms is called a precipitate.

45. Precipitation Reaction Example A more accurate representation is: 46. Reactant Solutions 47. Solution Post-Reaction 48. Precipitation Reaction Example

  • We look at all the possible combinations of the ions to check for compounds that form solids.
    • K 2 CrO 4
    • KNO 3
    • BaCrO 4
    • Ba(NO 3 ) 2

49. Precipitation Reaction Example Two of these combinations are the reactants and can be ruled out:

    • K 2 CrO 4
    • KNO 3
    • BaCrO 4
    • Ba(NO 3 ) 2

50. Solubility

  • Predicting the identity of a solid product in a precipitation reaction requires knowledge of the solubilities of common ionic substances.
    • Slightly soluble the tiny amount of solid that dissolves is not noticeable. The solid appears insoluble to the naked eye.
    • Insoluble and slightly soluble are often used interchangeably.

51. Simple Rules for the Solubility of Salts in Water 52. Precipitation Reaction Example Two of these combinations are the reactants and can be ruled out:

    • K 2 CrO 4
    • KNO 3
    • BaCrO 4
    • Ba(NO 3 ) 2

53. Precipitation Reactions

  • Precipitation reactions move forward due to the decrease in energy state of the compound.Bonds forming