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Chapter 15 Water and Aqueous Systems

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  1. Chapter 15Water and Aqueous Systems

  2. Water • Water makes plant and animal life on Earth possible • Water is present on the earth’s surface. • Water reserves are deep underground. • Ice and snow dominate the polar regions. • Water vapor from evaporation of surface water and steam from geysers and volcanoes is always present in the atmosphere.

  3. Water’s Properties H2O – the oxygen atom forms a covalent bond to each of the hydrogen atoms Because of its greater electronegativity, oxygen attracts the electron pair of the covalent O – H bond to a greater extend than hydrogen. As a result, the Oxygen atom acquires a partial negative charge (δ-) The less electronegative hydrogen atoms acquire partial positive charges (δ+)

  4. Water’s Properties The O – H bonds are highly polar. Polar bond – a covalent bond between atoms in which the electrons are shared unequally. How do the polarities of the two O – H bonds affect the polarity of the molecule? The shape of the molecule is the determining factor.

  5. Water’s Properties The bond angle of water is approximately 105 which give it a bent shape. Polar molecule – a molecule in which one side of the molecule is slightly negative and the opposite side is slightly positive. The water molecule as a whole is polar. Polarity – refers to the net molecular dipole resulting from electronegativity differences between covalently bonded atoms

  6. Water’s Properties In general, polar molecules are attracted to one another by dipole interactions. Dipole interactions – intermolecular forces resulting from the attraction of oppositely charged regions of polar molecules. The negative end of one molecule attracts the positive end of another molecule

  7. Water’s Properties The intermolecular attractions among water molecules result in the formation of hydrogen bonds. Hydrogen bonds – attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom. Many unique and important properties of water, including its high surface tension and low vapor pressure, result from hydrogen bonding.

  8. Surface Tension The surface of water acts like a skin. The skinlike property of water’s surface is explained by water’s ability to form hydrogen bonds. The molecules within the body of the liquid form hydrogen bonds with other molecules that surround them on all sides. The attractive forces on each of these molecules are balanced.

  9. Surface Tension Water molecules at the surface of the liquid experience an unbalanced attraction. Water molecules are hydrogen-bonded on only one side of the drop. As a result, water molecules at the surface tend to be drawn inward. Surface tension – the inward force, or pull that tends to minimize the surface area of a liquid

  10. Spherical Shape of Water Drops All liquids have surface tension, but water’s is higher than most. The surface tension of water tends to hold a drop of liquid in a spherical shape. The drop is not a perfect sphere, because the force of gravity tends to pull it down, causing it to flatten. This is why, on some surfaces, water tends to bead up rather than spread out.

  11. Why A Spherical Shape? Nature tends to find the path of least resistance Moving molecules takes work A spherical shape provides the minimum surface area for a given volume. Molecules expend the least energy possible to move into a spherical arrangement while maximizing their interactions with one another. To adopt any other shape, more work would have to be done. The worked saved is the surface tension.

  12. Surfactants It is possible to decrease the surface tension of water by adding a surfactant. Surfactant – any substance that interferes with the hydrogen bonding between water molecules and thereby reduces the surface tension. Examples of surfactants are soaps and detergents. Adding a detergent to beads of water on a greasy surface reduces the surface tension causing the beads of water to collapse and spread out.

  13. Vapor Pressure Hydrogen bonding also explains water’s unusually low vapor pressure. Vapor pressure is the result of molecules escaping the surface of the liquid & entering the vapor phase. Hydrogen bonds hold water molecules to one another. The tendency to escape is low, thus evaporation is slow. It is a good thing because all the lakes and oceans would tend to evaporate.

  14. Water in the Solid State Water is the solid state exhibits unique properties the same as a liquid. Ice cubes float in your glass of iced tea because solid water has a lower density than liquid water. As a typical liquid cools, it begins to contract and its density increases gradually. Increasing density means the molecules of the liquid move closer together so that a given volume of a liquid contains more molecules and thus more mass.

  15. Water in the Solid State Eventually the liquid solidifies and it continues to cool, having a greater density than the liquid Because the density of a typical solid is greater than that of its liquid, the solid sinks in its own liquid. As water begins to cool, it behaves initially like a typical liquid. It contracts slightly and its density gradually increases.

  16. Water in the Solid State When the temperature of water falls below 4º C, the density of water actually starts to decrease. Below 4º C, water no longer behaves like a typical liquid. Hydrogen bonds hold the water molecules in place in the solid phase. The structure of ice is a regular open framework of water molecules arranges like a honeycomb.

  17. Water in the Solid State Extensive hydrogen bonding in ice holds the water molecules farther apart in a more ordered arrangement than in liquid water. When ice melts, the framework collapses and the water molecules pack closer together, making liquid water more dense than ice.

  18. Water in the Solid State The fact that ice floats has important consequences for organisms. A layer of ice on the top of a pond acts as an insulator for the water beneath, preventing it from freezing solid except in extreme conditions. The liquid water at the bottom is warmer than 0ºC, fish and other aquatic life are better able to survive.

  19. Water in the Solid State Ice melts at 0ºC. This is a high melting temperature for a molecule with such a low molar mass. A considerable amount of energy is required to return water in the solid state to the liquid state. The heat absorbed when 1g of water changes from a solid to a liquid is 334 J.

  20. Question What causes the high surface tension and low vapor pressure of water? Water molecules are hydrogen-bonded to each other, but not to air molecules. Net attraction is inward, minimizing the water surface area. Hydrogen bonding makes it more difficult for water molecules to escape from the liquid phase to the vapor phase.

  21. Questions How would you describe the structure of ice? Honeycomb-like structure of hydrogen-bonded water molecules. What effect does a surfactant have on the surface tension of water? Surfactants lower the surface tension by interfering with hydrogen bonding.

  22. Questions What are two factors that determine how spherical a drop of liquid will be? Surface tension of a liquid tend to hold a drop of liquid in a spherical shape; gravity tends to flatten the drop. The molecules water (H2O) and methane (CH4) have similar masses, but methane changes from a gas to a liquid at -161 C. Water becomes a gas at 100ºC. What could account for the difference? Water has intermolecular hydrogen bonding between its molecules; methane does not.

  23. Questions Why is the surface tension of water so high compared to that of other liquids? Water molecules form a large number of hydrogen bonds, in addition to dipole-dipole forces between molecules. Why does water form a meniscus in a narrow tube? Water molecules have a greater attraction to the molecules on the surface of the glass than they do to each other.

  24. End of Section 15.1

  25. Solvents and Solutes Water dissolves so many of the substances that it comes in contact with that you won’t find chemically pure water in nature. Even the tap water you drink is a solution that contains varying amounts of dissolved minerals and gases. Aqueous solution – water that contains dissolved substances. Solvent – the dissolving medium Solute – the dissolved particles

  26. Solvents and Solutes Solutions are homogeneous mixtures. They are also stable mixtures. Example: salt (NaCl) does not settle out of the solution when allowed to stand. (provided other conditions, like temperature remain constant) Solute particles can be atoms, ions, or molecules and their average diameter are usually less than 1nm. If you filter a solution through filter paper, both the solute and the solvent pass through the filter.

  27. Solvents and Solutes Ionic compounds and polar covalent molecules dissolve most readily in water. Ionic compounds – composed of a positive and negative ion (ex: metal and non metal) Polar covalent molecules – electrons are shared equally between atoms (covalent) and one side of the molecule is slightly negative and the opposite side is slightly positive. Nonpolar covalent molecules, such as methane and compounds found in oil, grease & gasoline, do not dissolve in water.

  28. The Solution Process Water molecules are in constant motion because of their kinetic energy. When a crystal of NaCl is place in water, the water molecules collide with it. Since the water molecule is polar, the partial positive charge on the H+ attracts the negative solute ion Cl- The partial negative charge on the O2- attracts the positive solute ion Na+

  29. Solvation As individual solute ions break away from the crystal, the negatively (Cl-) and positively (Na+) charged ions become surrounded by solvent molecules and the ionic crystal dissolves. Solvation – the process by which the positive and negative ions on an ionic solid become surrounded by solvent molecules.

  30. Insoluble Ionic Compounds In some ionic compounds, the attractions among the ions in the crystals are stronger than the attractions exerted by water. These compounds cannot be solvated to any significant extent and are therefore nearly insoluble. Barium sulfate (BaSO4) and calcium carbonate (CaCO3) – nearly insoluble ionic compounds

  31. The Solution Process As a rule, polar solvents such as water dissolve ionic compounds and polar compounds. Nonpolar solvents such as gasoline dissolve nonpolar compounds. Like dissolves like

  32. Questions What must happen for an ionic solid to dissolve? The molecules of the solvent must be able to overcome the attractive forces between ions that hold the solid together. What part of a water molecules is attracted to a negatively charges solute ion? The hydrogen atoms

  33. Reminders Solutions are homogeneous mixtures containing a solvent and one or more solutes. Usually the solvent is defined as the component in the system that is present in the greatest amount. A water-soluble solute can be a solid, liquid, or a gas. Anion is a negatively charged atom or group of atoms Cation is a positively charges atom or group of atoms.

  34. Electrolytes & Nonelectrolytes Electrolyte – compound that conducts electric current when it is in an aqueous solution or in the molten state. All ionic compounds are electrolytes because they dissociate into ions. NaCl Na+ + Cl- Nonelectrolyte – compound that does not conduct electric current in aqueous solutions or in the molten state Many molecular compounds are nonelectrolyes because they are not composed of ions.

  35. Electrolytes & Nonelectrolytes Some polar molecular compounds are nonelectrolytes in the pure state, but become electrolytes when they dissolve in water. This occurs because they ionize in solution. Ex: neither ammonia or hydrogen chloride is an electrolyte in the pure state. NH3 + H2O NH4+ + OH- HCl + H2O H3O+ + Cl- Both conduct electricity in aqueous solutions because ions form.

  36. Strong Electrolytes Not all electrolytes conduct an electric current to the same degree. Strong Electrolyte – a solution that is a good conductor of electricity because a large portion of the solute exists as ions. Strong Acids HCl, HBr, HI, HNO3, HClO3, HClO4, and H2SO4 Strong Bases NaOH, KOH, LiOH, Ba(OH)2, and Ca(OH)2 Salts NaCl, KBr, MgCl2 …

  37. Electrolytes & Nonelectrolytes Weak electrolyte – solution that conducts electricity poorly because only a fraction of the solute exists as ions. Weak Acids HF, HC2H3O2 (acetic acid), H2CO3 (carbonic acid), H3PO4 (phosphoric acid) ….. Weak Bases NH3 (ammonia), C5H5N (pyridine), and several more, all containing "N"

  38. Electrolytes & Nonelectrolytes A solution conducts electricity if it contain ions. Electrolytes are excreted through the skin via sweat, and they must be replenished or cramps and heat stroke may occur. Sports drinks are a good source of electrolytes; they contain Na+, K+ and Ca+

  39. Hydrates When an aqueous solution of copper(II) sulfate (CuSO4) is allowed to evaporate, deep blue crystals of copper(II) sulfate pentahydrate are deposited. The chemical formula for this compound is CuSO4·5H2O Water of Hydration or Water of Crystallization – the water contained in a crystal.

  40. Hydrates Hydrate – a compound that contains water of hydration When writing the formula of a hydrate, use a dot to connect the formula of the compound and the number of water molecules per formula unit. CuSO4·5H2O Crystals of copper(II) sulfate pentahydrate always contain five molecules of water for each copper and sulfate ion pair.

  41. Hydrates When CuSO4·5H2O is heated above 100ºC, the crystals lose their water of hydration and crumble to a white anhydrous powder of CuSO4. If anhydrous CuSO4 is treated with water, the blue CuSO4·5H2O is regenerated. + heat CuSO4·5H2O CuSO4(s)+5 H2O(g) - heat Each hydrate contains a fixed quantity of water and has a definite composition.

  42. Efflorescent Hydrates The forces holding the water molecules in hydrates are not very strong, so the water is easily lost and regained. Because the water molecules are held by weak forces, it is often possible to estimate the vapor pressure of the hydrates. If a hydrate has a vapor pressure higher than the pressure of water vapor in the air, the hydrate will lose its water of hydration – effloresce.

  43. Hygroscopic Hydrates Hydrated salts that have a low vapor pressure remove water from moist air to form higher hydrates. These hydrates and other compounds that remove moisture from air are called hygroscopic. CaCl2· H2O CaCl2· 2H2O Calcium chloride monohydrate spontaneously absorbs a second molecule of water when exposed to moist air.

  44. Hygroscopic Hydrates CaCl2· H2O is used a a desiccant in the laboratory. Desiccant – a substance used to absorb moisture from the air and create a dry atmosphere. Desiccants can be added to a sealed container to keep substances inside the container dry. Desiccants can be added to liquid solvents to keep them dry. When a desiccant has absorbed all the water it can hold, it can be returned to its anhydrous state by heating.

  45. Percent Water • Suppose you are measuring a mass of Na2CO3 for a chemical reaction and you want to use the hydrate form of the compound because it is less expensive than the anhydrous compound. • To determine what percent of the hydrate is water. • Determine the mass of the number of moles of water in one mole of hydrate. • Determine the total mass of the hydrate. • % Water = mass of water / mass of hydrate x 100%

  46. Question What is the percent by mass of water in Na2CO3· 10H2O? Mass of 10 moles of H2O = 18 x 10 = 180 g Mass of Na2CO3· 10H2O = 46 + 12 + 48 + 180 = 286g % Water = mass of water / mass of hydrate x 100% % Water = 180 / 286 = 62.9 %

  47. Question What is the percent by mass of water in CuSO4· 5H2O? Mass of 5 moles of H2O = 18 x 5 = 90 g. Mass of CuSO4· 5H2O = 63.5 + 32.1 + 64 + 90 = 249.6 % Water = mass of water / mass of hydrate x 100% % Water = 90 / 249.6 = 36.1 %

  48. Question What is the percent by mass of water in CaCl2· 6H2O? Mass of 6 moles of H2O = 18 x 6 = 108 g Mass of CaCl2· 6H2O = 40.1 + 71 + 108 = 219.1g % Water = mass of water / mass of hydrate x 100% % Water = 108 / 219.1 = 49.3 %

  49. Questions In the formation of a solution, how does the solvent differ from the solute? The dissolving medium is the solvent, and the dissolved particles are the solute Describe what happens to the solute and the solvent when an ionic compound dissolves in water. As individual solute ions break away from the crystal, the negatively and positively charged ions become surrounded by solvent molecules and the ionic crystal dissolves.

  50. Questions Why are all ionic compounds electrolytes? Because they dissociate into ions. How do you write the formula for a hydrate? Use a dot to connect the formula of the compound with the number of water molecules per formula unit. Why are all ionic compounds electrolytes? Because they dissociate into ions.