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a cenote

a cenote. Introduction. Water in Life on Earth. Water covers almost two thirds of our planet Water has been the key to much of Earth’s evolutionary history Life itself is almost certainly originated from water Our body is about 60% water by mass

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a cenote

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  1. a cenote

  2. Introduction • Water in Life on Earth • Water covers almost two thirds of our planet • Water has been the key to much of Earth’s evolutionary history • Life itself is almost certainly originated from water • Our body is about 60% water by mass • Water possesses many unusual properties essential to supporting life on Earth

  3. Introduction • Water as a Medium • Water has an exceptional ability to dissolve a wide variety of substances • Water on Earth invariably contains a variety of dissolved substances • Water is the medium for most of the chemical reactions that take place within us and around us

  4. Introduction • Water as a Medium Figure 4.1 Cenote formation http://imnews.imbc.com/replay/nwdesk/article/3243453_5780.html

  5. 4.1 General Properties of Aqueous Solutions • Solution, solvent, and solute • Electrolytic Properties • Consider two aqueous solutions: NaCl (aq) & C6H12O6 (aq) • Electrolyte & nonelectrolyte

  6. 4.1 General Properties of Aqueous Solutions • Ionic & molecular compounds in water • Dissolution (solvation) process in water • We can predict the nature of ionic compounds (Na2SO4) in water from the chemical formula • Most molecular compounds are nonelectrolytes • Glucose, CH3OH, and HCl (ionic)

  7. 4.1 General Properties of Aqueous Solutions • Strong and weak electrolytes • Strong electrolytes exist in solution (nearly) completely as ions; NaCl, HCl • Weak electrolytes exist in solution mostly in the form of molecules with only a small fraction in the form of ions; CH3COOH (about 1% ionized) • Dissolution & ionization (CH3COOH & Ba(OH)2 ) • (weak) (strong) • Soluble ionic compounds are strong electrolytes: metal+nonmetal compounds & ammonium containing compounds (NaCl, FeSO4, Al(NO3)3, NH4Br, (NH4)2CO3)

  8. 4.2 Precipitation Reactions • Precipitation reaction

  9. 4.2 Precipitation Reactions • Solubility guidelines for ionic compounds • Solubility: the amount of a substance dissolved in a given quantity of solvent at a given temperature • Some ionic substances are insoluble in water because the attraction between the opposite charges is too great for the water molecules to separate the ions • No general rules to predict solubility of a substance • All ionic compounds containing NO3-, or CH3COO- • All ionic compounds containing the alkali metal ions or, NH4+

  10. 4.2 Precipitation Reactions • Solubility guidelines for ionic compounds

  11. 4.2 Precipitation Reactions • Exchange (Metathesis) Reactions • Ionic Equations • Molecular equations • Ionic equations (complete ionic equations) • Net ionic equations Spectator ions

  12. 4.3 Acid-Base Reactions • Many acids and bases are industrial and household substances • Hydrochloric acid: important industrial chemical & main component of gastric juice in our stomach.

  13. 4.3 Acid-Base Reactions • Acids • Substances that ionize in water to form H+, thereby increasing the concentration of H+ • Proton donors • Monoprotic acids: HCl and HNO3 • Diprotic acids: H2SO4 • CH3COOH (acetic acid)

  14. 4.3 Acid-Base Reactions • Bases • Substances that accept (react with) H+ • Bases produce hydroxide ions (OH-) when they dissolve in water • NaOH, KOH and Ca(OH)2 are common bases • NH3 is also a common base although it does not have OH-

  15. 4.3 Acid-Base Reactions • Strong and weak acids and bases • Strong acids and bases • Acids and bases that are completely ionized in solution • Weak acids and bases • Acids and bases that are partly ionized in solution

  16. 4.3 Acid-Base Reactions • Identifying strong and weak electrolytes

  17. 4.3 Acid-Base Reactions • Neutralization reactions and salts • Differences in the properties of acids and bases • Sour and bitter taste • Color changes with certain dyes • Neutralization reaction occurs when a solution of an acid and a solution of a base are mixed

  18. 4.3 Acid-Base Reactions • Neutralization reactions and salts

  19. 4.3 Acid-Base Reactions • Neutralization reactions and salts Sample Exercise 4.7 Writing Chemical Equations for a Neutralization Reaction For the reaction between aqueous solutions of acetic acid (CH3COOH) and barium hydroxide, Ba(OH)2, write (a) the balanced molecular equation, (b) the complete ionic equation, (c) the net ionic equation.

  20. 4.3 Acid-Ba se Reactions • Acid-base reactions with gas formation • Consider a acid-base reaction: • Carbonic acid is unstable

  21. Heartburn, reflux, indigestion, and sour stomach Histamine-2 (H2) blockers including cimetidine (Tagamet), famotidine (Pepcid), and ranitidine (Zantac), and the proton pump inhibitor (PPI) omeprazole (Prilosec)

  22. 4.4 Oxidation-Reduction Reactions • Oxidation and Reduction • Oxidation – loss of electrons by a substance • Reduction – gain of electrons by a substance

  23. 4.4 Oxidation-Reduction Reactions • Oxidation and Reduction • Ca(s) is oxidized and O2(g) is reduced

  24. 4.4 Oxidation-Reduction Reactions • Oxidation Numbers • The degree of oxidation of an atom in a chemical compound • Assigning oxidation number Elements in their elemental form have an oxidation number of 0. The oxidation number of a monatomic ion is the same as its charge. Nonmetals tend to have negative oxidation numbers, although some are positive in certain compounds or ions - O: -2 (-1 in O2-2), H: +1 (-1 when bonded to metals), F: -1, other halogens: -1 (+N when bonded to oxygen) The sum of the oxidation numbers in a neutral compound is 0. The sum of the oxidation numbers in a polyatomic ion is the charge on the ion.

  25. 4.4 Oxidation-Reduction Reactions • Oxidation Numbers

  26. 4.4 Oxidation-Reduction Reactions • Oxidation of metals by acids and salts • By acids • By salts

  27. 4.4 Oxidation-Reduction Reactions • Oxidation of metals by acids and salts

  28. 4.4 Oxidation-Reduction Reactions • Activity Series • Can we predict whether a certain metal will be oxidized either by an acid or by a particular salt? • It would be unwise to store a solution of nickel nitrate, Ni(NO3)2, in an iron container. • Different metals vary in the ease with which they are oxidized • Zn is oxidized by aqueous solution of Cu2+, but Ag is not. • Activity series of metals in aqueous solution • A list of metals arranged in order of decreasing ease of oxidation

  29. 4.4 Oxidation-Reduction Reactions • Activity Series Any metal on the list can be oxidized by the ions of elements below it

  30. 4.4 Oxidation-Reduction Reactions • Activity Series

  31. 4.4 Oxidation-Reduction Reactions • Activity Series

  32. Why is gold a special metal? - intrinsic beauty and rarity - soft and easy to be formed - the least active metal • Gold can be found in nature as a pure element rather than combined with oxygen or other elements, which account for its early discovery. • Used in jewelry (73%), coins (10%), and electronics (9%).

  33. 4.5 Concentrations of Solutions • Molarity • The concentration of a solution as the number of moles of solute in a liter of solution Figure 4.16 Preparing 0.250 L of a 1.00M solution of CuSO4.

  34. 4.5 Concentrations of Solutions • Molarity FW 142 g/mol

  35. 4.5 Concentrations of Solutions • The concentration of an electrolyte • Some substances produce multiple cations and/or anions(Na2SO4, MgCl2)

  36. 4.5 Concentrations of Solutions • Interconverting M, mol, and Volume FW 142 g/mol

  37. 4.5 Concentrations of Solutions • Dilution • How would you prepare 250.0 mL of 0.100 M cupric sulfate solution from 1.00 M solution?

  38. 4.5 Concentrations of Solutions • Dilution – preparation of solutions Figure 4.17

  39. 4.6 Solution Stoichiometry and Chemical Analysis • Let’s consider a question:

  40. 4.6 Solution Stoichiometry and Chemical Analysis FW 74.1 g/mol

  41. 4.6 Solution Stoichiometry and Chemical Analysis NaOH, FW 40.0 g/mol

  42. 4.6 Solution Stoichiometry and Chemical Analysis • Titrations • Titration is an analytical technique in which one can determine the concentration of a solute in a solution. • Standard solution • Equivalence point • Indicators

  43. 4.6 Solution Stoichiometry and Chemical Analysis • Titrations Figure 4.19

  44. 4.6 Solution Stoichiometry and Chemical Analysis

  45. 4.6 Solution Stoichiometry and Chemical Analysis

  46. FW 212.3 g/mol Ag3PO4, FW 212.3 g/mol

  47. 1.

  48. 2.

  49. 3.

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