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Lecture 17

Lecture 17. The Water We Drink (cont.) and Neutralizing the Threat of Acid Rain May 16, 2005. Ionic compounds. • Formed by elements with large differences in electronegativity • Do not form discrete units like molecular compounds • Collection of positive and negative ions

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Lecture 17

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  1. Lecture 17 The Water We Drink (cont.) and Neutralizing the Threat of Acid Rain May 16, 2005

  2. Ionic compounds • Formed by elements with large differences in electronegativity • Do not form discrete units like molecular compounds • Collection of positive and negative ions • Stack together to maximize attractions between opposite ions, minimize repulsions between negative ions • Regular repeating lattices

  3. Ionic and molecular compounds Ionic • High melting and boiling points • Lattices • Do not conduct electricity in solid state, do in liquid state • Metals and nonmetals • NaCl, CaCl2, MgSO4 Molecular • Low melting and boiling points • Discrete units • Do not conduct electricity in solid or liquid state • Usually made of non-metals • H2O, CO2, C6H12O6, CH4

  4. Polyatomic ions • Groups of atoms with “extra” or “too few” electrons PO4 has 32 e- So ion is 3- O P: 5 valence e- O: 6 each, 24 total 29 total valence e- O P O O

  5. Density • Ice is less dense than liquid water. Density: ratio of mass per unit volume grams/cm3 = g/mL Liquid water: 1 g/mL Ice: 0.92 g/mL

  6. Acids • Taste sour • React with carbonates (marble) and bicarbonates (baking soda) to produce CO2 • Produce H+ ions in water - protons • Directly: HCl H+ + Cl- H2SO4 H+ + HSO4- HNO3 H+ + NO3- CH3COOH H+ + CH3COO-

  7. Hydronium ion (H3O+) • Hydrogen ion is much too small and reactive to exist free in a water solution – Hydrogen atom: proton and electron – Hydrogen ion: proton – MUCH smaller than hydrogen atom, very high charge density • Reacts with water to form hydronium ion (H3O+) H+ + H2O H3O+

  8. H+ and H3O+ • In most cases can be used interchangeably • Only difference is whether or not water is explicitly written HCl H+ + Cl- HCl + H2O H3O+ + Cl- • You may see either reaction, but second is a more accurate description of what is actually going on

  9. Base • Produces OH- ions in water • Directly: NaOH Na+ + OH- Mg(OH)2 Mg2+ + 2OH- • Indirectly: NH3 + H2O NH4+ + OH-

  10. Acids and bases are opposites • Neutralize each other when mixed H+ + OH- H2O H3O+ + OH- 2H2O Another example HCl + NaOH NaCl + H2O H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O H+ + OH- H2O

  11. Acidic vs. Basic • A water solution can be acidic or basic or neutral. Acidic solution: [H+] > [OH-] Neutral solution: [H+] = [OH-] Basic solution: [H+] < [OH-] concentration (molarity)

  12. [H+] x [OH-] = Constant • In water solutions, product of H+ and OH-concentrations is a constant – If H+ concentration increases, OH- concentration must decrease – If OH- concentration increases, H+ concentration must decrease • Depends on temperature, but at 25 °C, the constant is 10-14

  13. [H+] x [OH-] = 10-14 • Thus, for a neutral solution: [H+] = [OH-] [H+] x [OH-] = [H+] x [H+] = [H+]2 = 10-14 [H+] = [OH-] = 10-7 M • Acidic solutions: [H+] > 10-7 M • Basic solutions: [OH-] > 10-7 M

  14. H+ concentrations in commonsolutions • Stomach acid: 0.1 M = 10-1 M • Vinegar: 0.001 M = 10-3 M • Pure water: 0.0000001 M = 10-7 M • Baking Soda: 0.000000001 M = 10-9 M • Ammonia: 0.0000000001 M = 10-10 M • Drano: 0.000000000001 M = 10-12 M

  15. pH Convenient way of describing how acidic or basic a solution is. Look at the exponents: • Stomach acid: 1 • Vinegar: 3 • Pure water: 7 • Baking Soda: 9 • Ammonia: 10 • Drano: 12 pH values

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