Chapter 12 The Liquids and Solids Around Us: Especially Water

1. Chapter 12The Liquids and Solids Around Us: Especially Water

2. No Gravity, No Spills • Undisturbed by air and gravity, liquids take the shape of a sphere. • The geometrical shape of lowest surface area to volume ratio is the sphere. • Spills would not spread over surfaces, as on Earth • Liquids and solids are held together by cohesive forces. • Without these forces all matter would be gaseous.

3. Each molecule within a liquid is attracted to its neighbors, but thermal energy keeps it moving in a random path. At each point in the path, the molecule is attracted to other neighboring molecules but not strongly enough to hold the molecule in one place. At the surface, there are fewer other molecules with which to interact. The liquid molecules may then be converted to a gas. Evaporation

4. In solids where the arrangement of molecules or atoms is well ordered each molecule or atom vibrates, but its average position remains fixed in the crystalline structure. In solids molecules move less, allowing them to interact more. The immobility of the molecules in ice give the material its properties. Will not flow and cannot be poured Relatively hard and does not easily break apart Crystalline Solids

5. Melting and Boiling • Thermal energy competes with cohesive forces to determine the phase of a substance. • At the melting point, thermal energy in a solid overcomes the forces of attraction between molecules. • At the boiling point, thermal energy in a liquid overcomes the forces of attraction between molecules.

6. Melting and Boiling In the cases of both melting and boiling, no chemical bonds within the molecules have been broken. Temperatures required to break chemical bonds are frequently much higher. The strength of cohesive forces is related to the molecule’s structure and influence melting and boiling points.

7. Forces that Hold Molecules Together • Dispersion forces (London forces) • Dipole-dipole interactions • Hydrogen bonding

8. The weakest cohesive force Present among all atoms and molecules Result from temporary induced dipoles formed when electrons are not distributed evenly in the molecule or atom For chemically similar elements or compounds, the magnitude of dispersion forces is generally proportional to the molar mass of the molecule or atom. Dispersion (London) Forces

9. Dispersion (London) Forces

10. Dispersion (London) Forces

11. Concept Check 12.1 • Which compound has the higher boiling point, CH3Br or CH2Br2.

12. Concept Check 12.1 Solution • Because the molecules are similar, the magnitude of the dispersion force should be proportional to molar mass. CH2Br2 has a higher molar mass than CH3Br, therefore we predict that CH2Br2 will have a higher boiling point than CH3Br.

13. Present in polar molecules that have permanent dipoles Molecules align such that positive ends of dipoles undergo attractive interactions with the negative ends of neighboring dipoles. Consequently, polar molecules have higher boiling points than nonpolar ones even given similar molecular weight. Polar bonds may or may not geometrically cancel within a molecule. Dipole Forces

14. Dipole-dipole attractive forces between polar molecules Bond dipoles cancel: nonpolar molecule Dipole Forces

15. Dipole Forces

16. Concept Check 12.2 • Which of the following molecules are polar? • Cl2 • HF • CH2Cl2 • CH4 • CH3CH2CH2OH • CO2

17. Concept Check 12.2 Solution • Which of the following molecules are polar? • Cl2: nonpolar, no polar bonds • HF: polar, H—F bond polar and not canceled by others. • CH2Cl2: polar, two C—Cl polar bonds not canceled by others. • CH4: nonpolar, no polar bonds • CH3CH2CH2OH: polar, contains polar CO and OH bonds. • CO2: nonpolar, two polar C=O bonds symmetrically opposed, canceling each other.

18. Like Dissolves Like • A polar substance will not mix with a nonpolar substance. • Soap serves as a molecular liaison between polar and nonpolar substances, dissolving nonpolar substances in very polar water.

19. Hydrogen Bonding A cohesive attraction between polar molecules containing hydrogen bonded to either fluorine, oxygen, or nitrogen leads to higher-than-expected boiling points.

20. Smelling Molecules: The Chemistry of Perfume • Perfume is a mixture of molecules with a range of cohesive forces. • The molecules with the weakest cohesive forces are the ones you smell first.

21. Volatility • Liquids that vaporize easily have a high vapor pressure. • Volatile • Weaker cohesive forces • Smaller heats of vaporization, evaporation less endothermic • Liquids that do not vaporize easily have a low vapor pressure. • Nonvolatile • Stronger cohesive forces • Larger heats of vaporization, evaporation more endothermic • Since perfumes have different components, their odors change through the day.

22. Concept Check 12.3 • Order the following molecules according to increasing boiling point: H2O CH4 NH3

23. Concept Check 12.3 Solution • H2O, CH4, and NH3 all have roughly the same molecular mass. • The O-H bond in H2O is the most polar making H2O the most polar molecule. The N—H bond in NH3 is the second most polar bond. Both N—H and O—H bonds can participate in hydrogen bonding. The least polar bond is the C—H bond in CH4, making CH4 a nonpolar molecule. • Boiling points increase as polarity increases in a series of molecules with similar molecular masses. CH4 < NH3 < H2O

24. Chemists Have Solutions • A solution is defined as a homogeneous mixture of two or more substances. • Solvent: The majority component • Solute: The minority component • Concentration • An expression of the amount of solute relative to the amount of solvent

25. Expressions of Concentration • Percent by mass = • Percent by volume = • Molarity = • Parts per million (ppm) =

26. Concept Check 12.4 • 25 g of NaCl is mixed with enough water to make 750. mL of solution. What is the molarity of the resulting sodium chloride solution?

27. Concept Check 12.4 Solution • The concentration as first presented is in g/mL. Molarity is concentration unit in moles/L, therefore, a series of conversions must be used. • Molar mass NaCl = 58.4 g

28. Concept Check 12.5 • How many grams of NaCl are present in 1.3 L of a 0.57 M NaCl solution?

29. Concept Check 12.5 Solution • Molarity has units of mol/L and is therefore a conversion factor between moles and liters. Use the volume and the molarity to find moles of NaCl and then use the molar mass of NaCl (58.5 g/mol) to find grams of NaCl:

30. Concept Check 12.6 • A 100.0 g sample of well water is found to contain 0.0092 g of magnesium carbonate (MgCO3). What is the concentration of magnesium carbonate in parts per million?

31. Concept Check 12.6 Solution • Divide mass of MgCO3 (in grams) by the mass of solution (in grams), then multiply by 106 ppm.

32. Concept Check 12.7 • A 15.0 mL sample of blood is found to contain 2.03 mg of calcium. What is the concentration of calcium in milligrams per liter?

33. Concept Check 12.7 Solution • Start with the original concentration as given, then convert from milliliters to liters by multiplying by the conversion factor.

34. Most common liquid on Earth Unusual molecule Highly polar Two H–O bonds Material expands upon freezing No other compound of similar molar mass comes close to the high boiling point of water. Molar mass of 18 grams per mole Boiling point of 100º C Water: An Oddity Among Molecules

35. Ice: Unique Structure • Highly polar • Two O—H bonds that hydrogen bond to other water molecules • Expands when it freezes • Ice is less dense than liquid water and will float. • Ice on bodies of water insulates the marine life beneath.

36. Cells in biological tissue are damaged when frozen due to the expansion of water on freezing. Water dissolves many polar organic and inorganic compounds and is responsible for the flow of nutrients in the body. Unique Properties

37. Water:Where Is It and How Did It Get There? • Only a tiny fraction of Earth’s water is readily usable, that taken from lakes and streams (0.014%). • The oceans hold 97.4%, and 2.6% is in ice caps, glaciers, and ground water. • Land water comes from oceans through the hydrologic cycle.

38. Water: Pure or Polluted? • Virtually no water is pure; pure water can only be made in a laboratory. • Like most liquids, water is a mixture, containing a number of different elements and compounds. • Some components are harmful and some beneficial. • Hard water • As water runs through soils rich in limestone, calcium and magnesium ions dissolve into it. • Such hard water has no adverse health effects, but scaly deposits on pipes, fixtures, utensils can form. • Hard water decreases the effectiveness of soap and causes bathtub ring.

39. Softeners are charged with sodium ions that exchange with calcium and magnesium ions. Sodium ions increase the risk of high blood pressure. Water Softening

40. Microorganisms that cause diseases like hepatitis, cholera, typhoid, and dysentery come from the dumping of human and animal waste near a drinking water source. Biological Contaminants

41. Chemical Contaminates Pollution sources: • Dumping of wastes by industry into streams and rivers • Dumping of industrial waste into the atmosphere through emissions are returned to water supplies by precipitation • Pesticides and fertilizers are significant contaminants and end up in drinking water supplies.

42. Organic Contaminants • Volatile • Benzene, carbon tetrachloride, others • Nonvolatile • Ethylbenzene, chlorobenzene, TCE, PCB, and dioxins • Both types come from fertilizers, gasoline, pesticides, paints, and solvents. • They increase cancer risk, liver and kidney damage, and CNS damage.

43. Concept Check 12.8 • A water sample near a leaking gasoline storage tank contains 15.0 ppm benzene (C6H6), a carcinogen. If a person consumes 1.50 liters of this water each day, how many grams of benzene have they consumed daily? The density of water sample is 1.00 g/mL.

44. Concept Check 12.8 Solution • First, convert the original concentration of 15 ppm benzene in sample water (15.0 g benzene/106 g sample) to g of benzene/L of solution. • Next, multiply the sample size by the concentration conversion factor previously calculated.

45. Inorganic Contaminants • Asbestos: From natural and human-made supplies • Nitrates: Immediate danger to humans • Diminish hemoglobin’s ability to carry oxygen • Blue baby syndrome • Heavy metals: Mercury, lead • Kidney and CNS damage

46. Concept Check 12.9 • How many milligrams of mercury (Hg) are present in 5.7 × 103 L of water with a mercury concentration of 0.0062 mg/L?

47. Concept Check 12.9 Solution • Multiplying the sample size (5.7 × 103 L of water) by the concentration (a conversion factor) gives the mass (in grams) of Hg present in the sample.

48. Radioactive Contaminants • More common in areas with uranium-rich deposits • Uranium, radium, and radon all come primarily from natural sources. • Ingestion increases cancer risk

49. The Safe Drinking Water Act • Passed in 1974 • Set maximum contaminant levels (MCLs) • Public water suppliers sample and test water. • Responsible for notifying government agencies and YOU of violations • EPA requires water suppliers to actively prevent contamination during transport between treatment and your tap.

50. Public Water Treatment • To meet EPA requirements, providers purify and treat water before delivery.