Answer: Fluorocarbons (compounds that contain both carbon

Rank the contents of the following containers in order of increasing entropy: [Section 13.1] (a) (b) (c)

This figure shows the interaction of a cation with surrounding water molecules. + (a) Which atom of water is associated with the cation? Explain. (b) Which of the following explanations accounts for the fact that the ion-solvent interaction is greater for Li + than for K+? a. Li + is of lower mass than K+. b. The ionization energy of Li is higher than that for K. c. Li + has a smaller ionic radius than K+. d. Li has a lower density than K. e. Li reacts with water more slowly than K. [Section 13.1]

Consider two ionic solids, both composed of singly-charged ions, that have different lattice energies. (a) Will the solids have the same solubility in water? (b) If not, which solid will be more soluble in water, the one with the larger lattice energy or the one with the smaller lattice energy? Assume that solute-solvent interactions are the same for both solids. [Section 13.1]

Are gases always miscible with each other? Explain. [Section 13.1]

Which of the following is the best representation of a saturated solution? Explain your reasoning. [Section 13.2] (a) (b) (c)

The solubility of Xe in water at 1 atm pressure and 20 C is approximately 5 * 10-3 M. (a) Compare this with the solubilities of Ar and Kr in water (Table 13.1). (b) What properties of the rare gas atoms account for the variation in solubility? [Section 13.3]

The structures of vitamins E and B6 are shown below. Predict which is more water soluble and which is more fat soluble. Explain. [Section 13.3] Vitamin B6 Vitamin E

You take a sample of water that is at room temperature and in contact with air and put it under a vacuum. Right away, you see bubbles leave the water, but after a little while, the bubbles stop. As you keep applying the vacuum, more bubbles appear. A friend tells you that the first bubbles were water vapor, and the low pressure had reduced the boiling point of water, causing the water to boil. Another friend tells you that the first bubbles were gas molecules from the air (oxygen, nitrogen, and so forth) that were dissolved in the water. Which friend is mostly likely to be correct? What, then, is responsible for the second batch of bubbles? [Section 13.4]

The figure shows two identical volumetric flasks containing the same solution at two temperatures. (a) Does the molarity of the solution change with the change in temperature? Explain. (b) Does the molality of the solution change with the change in temperature? Explain. [Section 13.4] 25 C 55 C

This portion of a phase diagram shows the vapor-pressure curves of a volatile solvent and of a solution of that solvent containing a nonvolatile solute. (a) Which line represents the solution? (b) What are the normal boiling points of the solvent and the solution? [Section 13.5] 40 1.0 50 60 70 T (C) P (atm)

Suppose you had a balloon made of some highly flexible semipermeable membrane. The balloon is filled completely with a 0.2 M solution of some solute and is submerged in a 0.1 M solution of the same solute: 0.1 M 0.2 M Initially, the volume of solution in the balloon is 0.25 L. Assuming the volume outside the semipermeable membrane is large, as the illustration shows, what would you expect for the solution volume inside the balloon once the system has come to equilibrium through osmosis? [Section 13.5]

The molecule n-octylglucoside, shown here, is widely used in biochemical research as a nonionic detergent for solubilizing large hydrophobic protein molecules. What characteristics of this molecule are important for its use in this way? [Section 13.6]

Indicate whether each statement is true or false: (a) A solute will dissolve in a solvent if solutesolute interactions are stronger than solute-solvent interactions.(b) In making a solution, the enthalpy of mixing is always a positive number. (c) An increase in entropy favors mixing.

Indicate whether each statement is true or false: (a) NaCl dissolves in water but not in benzene 1C6H62 because benzene is denser than water. (b) NaCl dissolves in water but not in benzene because water has a large dipole moment and benzene has zero dipole moment. (c) NaCl dissolves in water but not in benzene because the waterion interactions are stronger than benzeneion interactions.

Indicate the type of solutesolvent interaction (Section 11.2) that should be most important in each of the following solutions: (a) CCl4 in benzene 1C6H62, (b) methanol 1CH3OH2 in water, (c) KBr in water, (d) HCl in acetonitrile 1CH3CN2.

Indicate the principal type of solutesolvent interaction in each of the following solutions and rank the solutions from weakest to strongest solutesolvent interaction: (a) KCl in water, (b) CH2Cl2 in benzene 1C6H62, (c) methanol 1CH3OH2 in water.

An ionic compound has a very negative Hsoln in water. (a) Would you expect it to be very soluble or nearly insoluble in water? (b) Which term would you expect to be the largest negative number: Hsolvent, Hsolute, or Hmix?

When ammonium chloride dissolves in water, the solution becomes colder. (a) Is the solution process exothermic or endothermic? (b) Why does the solution form?

(a) In Equation 13.1, which of the enthalpy terms for dissolving an ionic solid would correspond to the lattice energy?

(b) Which energy term in this equation is always exothermic?

For the dissolution of LiCl in water, Hsoln = -37 kJ>mol. Which term would you expect to be the largest negative number: Hsolvent, Hsolute, or Hmix?

Two nonpolar organic liquids, hexane 1C6H142 and heptane 1C7H162, are mixed. (a) Do you expect Hsoln to be a large positive number, a large negative number, or close to zero? Explain. (b) Hexane and heptane are miscible with each other in all proportions. In making a solution of them, is the entropy of the system increased, decreased, or close to zero, compared to the separate pure liquids?

The enthalpy of solution of KBr in water is about +198 kJ>mol. Nevertheless, the solubility of KBr in water is relatively high. Why does the solution process occur even though it is endothermic?

The solubility of Cr1NO323 # 9 H2O in water is 208 g per 100 g of water at 15 C. A solution of Cr1NO323 # 9 H2O in water at 35 C is formed by dissolving 324 g in 100 g of water. When this solution is slowly cooled to 15 C, no precipitate forms. (a) What term describes this solution? (b) What action might you take to initiate crystallization? Use molecular-level processes to explain how your suggested procedure works.

The solubility of MnSO4 # H2O in water at 20 C is 70 g per 100 mL of water. (a) Is a 1.22 M solution of MnSO4 # H2O in water at 20 C saturated, supersaturated, or unsaturated? (b) Given a solution of MnSO4 # H2O of unknown concentration, what experiment could you perform to determine whether the new solution is saturated, supersaturated, or unsaturated?

The solubility of MnSO4 # H2O in water at 20 C is 70 g per 100 mL of water. (a) Is a 1.22 M solution of MnSO4 # H2O in water at 20 C saturated, supersaturated, or unsaturated? (b) Given a solution of MnSO4 # H2O of unknown concentration, what experiment could you perform to determine whether the new solution is saturated, supersaturated, or unsaturated?

By referring to Figure 13.15, determine the mass of each of the following salts required to form a saturated solution in 250 g of water at 30 C: (a) KClO3, (b) Pb1NO322, (c) Ce21SO423.

Consider water and glycerol, CH21OH2CH1OH2CH2OH. (a) Would you expect them to be miscible in all proportions? Explain. (b) List the intermolecular attractions that occur between a water molecule and a glycerol molecule.

Oil and water are immiscible. Which is the most likely reason? (a) Oil molecules are denser than water. (b) Oil molecules are composed mostly of carbon and hydrogen. (c) Oil molecules have higher molar masses than water. (d) Oil molecules have higher vapor pressures than water. (e) Oil molecules have higher boiling points than water

Common laboratory solvents include acetone 1CH3COCH32, methanol 1CH3OH2, toluene 1C6H5CH32, and water. Which of these is the best solvent for nonpolar solutes?

Would you expect alanine (an amino acid) to be more soluble in water or in hexane? Explain. Alanine

(a) Would you expect stearic acid, CH31CH2216COOH, to be more soluble in water or in carbon tetrachloride? Explain.

(b) Which would you expect to be more soluble in water, cyclohexane or dioxane? Explain. H2C H2C CH2 CH2 O O Dioxane H2C H2C CH2 CH2 CH2 CH2 Cyclohexane

Ibuprofen, widely used as a pain reliever, has a limited solubility in water, less than 1 mg>mL. Which part of the molecules structure (gray, white, red) contributes to its water solubility? Which part of the molecule (gray, white, red) contributes to its water insolubility? Ibuprofen

Which of the following in each pair is likely to be more soluble in hexane, C6H14: (a) CCl4 or CaCl2, (b) benzene 1C6H62 or glycerol, CH21OH2CH1OH2CH2OH, (c) octanoic acid, CH3CH2CH2CH2CH2CH2CH2COOH, or acetic acid, CH3COOH? Explain your answer in each case

Which of the following in each pair is likely to be more soluble in water: (a) cyclohexane 1C6H122 or glucose 1C6H12O62, (b) propionic acid 1CH3CH2COOH2 or sodium propionate 1CH3CH2COONa2, (c) HCl or ethyl chloride 1CH3CH2Cl2? Explain in each case.

(a) Explain why carbonated beverages must be stored in sealed containers.

(b) Once the beverage has been opened, why does it maintain more carbonation when refrigerated than at room temperature?

Explain why pressure substantially affects the solubility of O2 in water but has little effect on the solubility of NaCl in water.

The Henrys law constant for helium gas in water at 30 C is 3.7 * 10-4 M>atm and the constant for N2 at 30 C is 6.0 * 10-4 M>atm. If the two gases are each present at 1.5 atm pressure, calculate the solubility of each gas.

The partial pressure of O2 in air at sea level is 0.21 atm. Using the data in Table 13.1, together with Henrys law, calculate the molar concentration of O2 in the surface water of a mountain lake saturated with air at 20 C and an atmospheric pressure of 650 torr.

(a) Calculate the mass percentage of Na2SO4 in a solution containing 10.6 g of Na2SO4 in 483 g of water.

(b) An ore contains 2.86 g of silver per ton of ore. What is the concentration of silver in ppm?

(a) What is the mass percentage of iodine in a solution containing 0.035 mol I2 in 125 g of CCl4?

(b) Seawater contains 0.0079 g of Sr 2+ per kilogram of water. What is the concentration of Sr 2+ in ppm?

A solution is made containing 14.6 g of CH3OH in 184 g of H2O. Calculate (a) the mole fraction of CH3OH, (b) the mass percent of CH3OH, (c) the molality of CH3OH.

A solution is made containing 20.8 g of phenol 1C6H5OH2 in 425 g of ethanol 1CH3CH2OH2. Calculate (a) the mole fraction of phenol, (b) the mass percent of phenol, (c) the molality of phenol.

Calculate the molarity of the following aqueous solutions: (a) 0.540 g of Mg1NO322 in 250.0 mL of solution, (b) 22.4 g of LiClO4 # 3 H2O in 125 mL of solution, (c) 25.0 mL of 3.50 M HNO3 diluted to 0.250 L.

What is the molarity of each of the following solutions: (a) 15.0 g of Al21SO423 in 0.250 mL solution, (b) 5.25 g of Mn1NO322 # 2 H2O in 175 mL of solution, (c) 35.0 mL of 9.00 M H2SO4 diluted to 0.500 L?

Calculate the molality of each of the following solutions: (a) 8.66 g of benzene 1C6H62 dissolved in 23.6 g of carbon tetrachloride 1CCl42, (b) 4.80 g of NaCl dissolved in 0.350 L of water

(a) What is the molality of a solution formed by dissolving 1.12 mol of KCl in 16.0 mol of water?

(b) How many grams of sulfur 1S82 must be dissolved in 100.0 g of naphthalene 1C10H82 to make a 0.12 m solution?

A sulfuric acid solution containing 571.6 g of H2SO4 per liter of solution has a density of 1.329 g>cm3 . Calculate A sulfuric acid solution containing 571.6 g of H2SO4 perliter of solution has a density of 1.329 g>cm3 .

Ascorbic acid 1vitamin C, C6H8O62 is a water-soluble vitamin. A solution containing 80.5 g of ascorbic acid dissolved in 210 g of water has a density of 1.22 g>mL at 55 C. Calculate (a) the mass percentage, (b) the mole fraction, (c) the molality, (d) the molarity of ascorbic acid in this solution.

The density of acetonitrile 1CH3CN2 is 0.786 g>mL and the density of methanol 1CH3OH2 is 0.791 g>mL. A solution is made by dissolving 22.5 mL of CH3OH in 98.7 mL of CH3CN. (a) What is the mole fraction of methanol in the solution? (b) What is the molality of the solution? (c) Assuming that the volumes are additive, what is the molarity of CH3OH in the solution?

The density of toluene 1C7H82 is 0.867 g>mL, and the density of thiophene 1C4H4S2 is 1.065 g>mL. A solution is made by dissolving 8.10 g of thiophene in 250.0 mL of toluene. (a) Calculate the mole fraction of thiophene in the solution. (b) Calculate the molality of thiophene in the solution. (c) Assuming that the volumes of the solute and solvent are additive, what is the molarity of thiophene in the solution?

Calculate the number of moles of solute present in each of the following aqueous solutions: (a) 600 mL of 0.250 M SrBr2, (b) 86.4 g of 0.180 m KCl, (c) 124.0 g of a solution that is 6.45% glucose 1C6H12O62 by mass.

Calculate the number of moles of solute present in each of the following solutions: (a) 255 mL of 1.50 M HNO31aq2, (b) 50.0 mg of an aqueous solution that is 1.50 m NaCl, (c) 75.0 g of an aqueous solution that is 1.50% sucrose 1C12H22O112 by mass.

Describe how you would prepare each of the following aqueous solutions, starting with solid KBr: (a) 0.75 L of 1.5 * 10-2 M KBr, (b) 125 g of 0.180 m KBr, (c) 1.85 L of a solution that is 12.0% KBr by mass (the density of the solution is 1.10 g>mL), (d) a 0.150 M solution of KBr that contains just enough KBr to precipitate 16.0 g of AgBr from a solution containing 0.480 mol of AgNO3.

Describe how you would prepare each of the following aqueous solutions: (a) 1.50 L of 0.110 M 1NH422SO4 solution, starting with solid 1NH422SO4; (b) 225 g of a solution that is 0.65 m in Na2CO3, starting with the solid solute; (c) 1.20 L of a solution that is 15.0% Pb1NO322 by mass (the density of the solution is 1.16 g>mL), starting with solid solute; (d) a 0.50 M solution of HCl that would just neutralize 5.5 g of Ba1OH22 starting with 6.0 M HCl.

Commercial aqueous nitric acid has a density of 1.42 g>mL and is 16 M. Calculate the percent HNO3 by mass in the solution.

Commercial concentrated aqueous ammonia is 28% NH3 by mass and has a density of 0.90 g>mL. What is the molarity of this solution?

Brass is a substitutional alloy consisting of a solution of copper and zinc. A particular sample of red brass consisting of 80.0% Cu and 20.0% Zn by mass has a density of 8750 kg>m3 . (a) What is the molality of Zn in the solid solution? (b) What is the molarity of Zn in the solution?

Caffeine 1C8H10N4O22 is a stimulant found in coffee and tea. If a solution of caffeine in the solvent chloroform 1CHCl32 has a concentration of 0.0500 m, calculate (a) the percentage of caffeine by mass, (b) the mole fraction of caffeine in the solution. Caffeine

During a persons typical breathing cycle, the CO2 concentration in the expired air rises to a peak of 4.6% by volume. (a) Calculate the partial pressure of the CO2 in the expired air at its peak, assuming 1 atm pressure and a body temperature of 37 C. (b) What is the molarity of the CO2 in the expired air at its peak, assuming a body temperature of 37 C?

Breathing air that contains 4.0% by volume CO2 over time causes rapid breathing, throbbing headache, and nausea, among other symptoms. What is the concentration of CO2 in such air in terms of (a) mol percentage, (b) molarity, assuming 1 atm pressure and a body temperature of 37 C?

You make a solution of a nonvolatile solute with a liquid solvent. Indicate whether each of the following statements is true or false. (a) The freezing point of the solution is higher than that of the pure solvent. (b) The freezing point of the solution is lower than that of the pure solvent. (c) The boiling point of the solution is higher than that of the pure solvent. (d) The boiling point of the solution is lower than that of the pure solvent.

You make a solution of a nonvolatile solute with a liquid solvent. Indicate if each of the following statements is true or false. (a) The freezing point of the solution is unchanged by addition of the solvent. (b) The solid that forms as the solution freezes is nearly pure solute. (c) The freezing point of the solution is independent of the concentration of the solute. (d) The boiling point of the solution increases in proportion to the concentration of the solute. (e) At any temperature, the vapor pressure of the solvent over the solution is lower than what it would be for the pure solvent.

Consider two solutions, one formed by adding 10 g of glucose 1C6H12O62 to 1 L of water and the other formed by adding 10 g of sucrose 1C12H22O112 to 1 L of water. Calculate the vapor pressure for each solution at 20 C; the vapor pressure of pure water at this temperature is 17.5 torr

(a) What is an ideal solution?

(b) The vapor pressure of pure water at 60 C is 149 torr. The vapor pressure of water over a solution at 60 C containing equal numbers of moles of water and ethylene glycol (a nonvolatile solute) is 67 torr. Is the solution ideal according to Raoults law? Explain.

(a) Calculate the vapor pressure of water above a solution prepared by adding 22.5 g of lactose 1C12H22O112 to 200.0 g of water at 338 K. (Vapor-pressure data for water are given in Appendix B.)

(b) Calculate the mass of propylene glycol 1C3H8O22 that must be added to 0.340 kg of water to reduce the vapor pressure by 2.88 torr at 40 C.

(a) Calculate the vapor pressure of water above a solution prepared by dissolving 28.5 g of glycerin 1C3H8O32 in 125 g of water at 343 K. (The vapor pressure of water is given in Appendix B.)

(b) Calculate the mass of ethylene glycol 1C2H6O22 that must be added to 1.00 kg of ethanol 1C2H5OH2 to reduce its vapor pressure by 10.0 torr at 35 C. The vapor pressure of pure ethanol at 35 C is 1.00 * 102 torr.

At 63.5 C, the vapor pressure of H2O is 175 torr, and that of ethanol 1C2H5OH2 is 400 torr. A solution is made by mixing equal masses of H2O and C2H5OH. (a) What is the mole fraction of ethanol in the solution? (b) Assuming ideal-solution behavior, what is the vapor pressure of the solution at 63.5 C? (c) What is the mole fraction of ethanol in the vapor above the solution?

At 20 C, the vapor pressure of benzene 1C6H62 is 75 torr, and that of toluene 1C7H82 is 22 torr. Assume that benzene and toluene form an ideal solution. (a) What is the composition in mole fraction of a solution that has a vapor pressure of 35 torr at 20 C? (b) What is the mole fraction of benzene in the vapor above the solution described in part (a)?

(a) Does a 0.10 m aqueous solution of NaCl have a higher boiling point, a lower boiling point, or the same boiling point as a 0.10 m aqueous solution of C6H12O6?

(b) The experimental boiling point of the NaCl solution is lower than that calculated assuming that NaCl is completely dissociated in solution. Why is this the case?

Arrange the following aqueous solutions, each 10% by mass in solute, in order of increasing boiling point: glucose 1C6H12O62, sucrose 1C12H22O112, sodium nitrate 1NaNO32.

List the following aqueous solutions in order of increasing boiling point: 0.120 m glucose, 0.050 m LiBr, 0.050 m Zn1NO322.

List the following aqueous solutions in order of decreasing freezing point: 0.040 m glycerin 1C3H8O32, 0.020 m KBr, 0.030 m phenol 1C6H5OH2.

Using data from Table 13.3, calculate the freezing and boiling points of each of the following solutions: (a) 0.22 m glycerol 1C3H8O32 in ethanol, (b) 0.240 mol of naphthalene 1C10H82 in 2.45 mol of chloroform, (c) 1.50 g NaCl in 0.250 kg of water, (d) 2.04 g KBr and 4.82 g glucose 1C6H12O62 in 188 g of water.

Using data from Table 13.3, calculate the freezing and boiling points of each of the following solutions: (a) 0.25 m glucose in ethanol; (b) 20.0 g of decane, C10H22, in 50.0 g CHCl3;(c) 3.50 g NaOH in 175 g of water, (d) 0.45 mol ethylene glycol and 0.15 mol KBr in 150 g H2O.

How many grams of ethylene glycol 1C2H6O22 must be added to 1.00 kg of water to produce a solution that freezes at -5.00 C?

What is the freezing point of an aqueous solution that boils at 105.0 C?

What is the osmotic pressure formed by dissolving 44.2 mg of aspirin 1C9H8O42 in 0.358 L of water at 25 C?

Seawater contains 3.4 g of salts for every liter of solution. Assuming that the solute consists entirely of NaCl (in fact, over 90% of the salt is indeed NaCl), calculate the osmotic pressure of seawater at 20 C.

Adrenaline is the hormone that triggers the release of extra glucose molecules in times of stress or emergency. A solution of 0.64 g of adrenaline in 36.0 g of CCl4 elevates the boiling point by 0.49 C. Calculate the approximate molar mass of adrenaline from this data. Adrenaline

Lauryl alcohol is obtained from coconut oil and is used to make detergents. A solution of 5.00 g of lauryl alcohol in 0.100 kg of benzene freezes at 4.1 C. What is the molar mass of lauryl alcohol from this data?

Lysozyme is an enzyme that breaks bacterial cell walls. A solution containing 0.150 g of this enzyme in 210 mL of solution has an osmotic pressure of 0.953 torr at 25 C. What is the molar mass of lysozyme?

A dilute aqueous solution of an organic compound soluble in water is formed by dissolving 2.35 g of the compound in water to form 0.250 L of solution. The resulting solution has an osmotic pressure of 0.605 atm at 25 C. Assuming that the organic compound is a nonelectrolyte, what is its molar mass?

The osmotic pressure of a 0.010 M aqueous solution of CaCl2 is found to be 0.674 atm at 25 C. (a) Calculate the vant Hoff factor, i, for the solution. (b) How would you expect the value of i to change as the solution becomes more concentrated? Explain.

Based on the data given in Table 13.4, which solution would give the larger freezing-point lowering, a 0.030 m solution of NaCl or a 0.020 m solution of K2SO4? How do you explain the departure from ideal behavior and the differences observed between the two salts?

(a) Do colloids made only of gases exist? Why or why not?

(b) In the 1850s, Michael Faraday prepared ruby-red colloids of gold nanoparticles in water that are still stable today. These brightly colored colloids look like solutions. What experiment(s) could you do to determine whether a given colored preparation is a solution or colloid?

Choose the best answer: A colloidal dispersion of one liquid in another is called (a) a gel, (b) an emulsion, (c) a foam, (d) an aerosol.

An emulsifying agent is a compound that helps stabilize a hydrophobic colloid in a hydrophilic solvent (or a hydrophilic colloid in a hydrophobic solvent). Which of the following choices is the best emulsifying agent? (a) CH3COOH, (b) CH3CH2CH2COOH, (c) CH31CH2211COOH, (d) CH31CH2211COONa.

Aerosols are important components of the atmosphere. Does the presence of aerosols in the atmosphere increase or decrease the amount of sunlight that arrives at the Earths surface, compared to an aerosol-free atmosphere? Explain your reasoning.

roteins can be precipitated out of aqueous solution by the addition of an electrolyte; this process is called salting out the protein. (a) Do you think that all proteins would be precipitated out to the same extent by the same concentration of the same electrolyte? (b) If a protein has been salted out, are the proteinprotein interactions stronger or weaker than they were before the electrolyte was added? (c) A friend of yours who is taking a biochemistry class says that salting out works because the waters of hydration that surround the protein prefer to surround the electrolyte as the electrolyte is added; therefore, the proteins hydration shell is stripped away, leading to protein precipitation. Another friend of yours in the same biochemistry class says that salting out works because the incoming ions adsorb tightly to the protein, making ion pairs on the protein surface, which end up giving the protein a zero net charge in water and therefore leading to precipitation. Discuss these two hypotheses. What kind of measurements would you need to make to distinguish between these two hypotheses?

Explain how (a) a soap such as sodium stearate stabilizes a colloidal dispersion of oil droplets in water; (b) milk curdles upon addition of an acid.

Butylated hydroxytoluene (BHT) has the following molecular structure: H3C C CH3 CH3 OH CH3 BHT CH3 C CH3 CH3 It is widely used as a preservative in a variety of foods, including dried cereals. Based on its structure, would you expect BHT to be more soluble in water or in hexane 1C6H142? Explain.

A saturated solution of sucrose 1C12H22O112 is made by dissolving excess table sugar in a flask of water. There are 50 g of undissolved sucrose crystals at the bottom of the flask in contact with the saturated solution. The flask is stoppered and set aside. A year later a single large crystal of mass 50 g is at the bottom of the flask. Explain how this experiment provides evidence for a dynamic equilibrium between the saturated solution and the undissolved solute.

Most fish need at least 4 ppm dissolved O2 in water for survival. (a) What is this concentration in mol>L? (b) What partial pressure of O2 above water is needed to obtain 4 ppm O2 in water at 10 C? (The Henrys law constant for O2 at this temperature is 1.71 * 10-3 mol>L@atm.)

The presence of the radioactive gas radon (Rn) in well water presents a possible health hazard in parts of the United States. (a) Assuming that the solubility of radon in water with 1 atm pressure of the gas over the water at 30 C is 7.27 * 10-3 M, what is the Henrys law constant for radon in water at this temperature? (b) A sample consisting of various gases contains 3.5 * 10-6 mole fraction of radon. This gas at a total pressure of 32 atm is shaken with water at 30 C. Calculate the molar concentration of radon in the water

Glucose makes up about 0.10% by mass of human blood. Calculate this concentration in (a) ppm, (b) molality. (c) What further information would you need to determine the molarity of the solution?

The concentration of gold in seawater has been reported to be between 5 ppt (parts per trillion) and 50 ppt. Assuming that seawater contains 13 ppt of gold, calculate the number of grams of gold contained in 1.0 * 103 gal of seawater.

The maximum allowable concentration of lead in drinking water is 9.0 ppb. (a) Calculate the molarity of lead in a 9.0- ppb solution. (b) How many grams of lead are in a swimming pool containing 9.0 ppb lead in 60 m3 of water?

Acetonitrile 1CH3CN2 is a polar organic solvent that dissolves a wide range of solutes, including many salts. The density of a 1.80 M LiBr solution in acetonitrile is 0.826 g>cm3 . Calculate the concentration of the solution in (a) molality, (b) mole fraction of LiBr, (c) mass percentage of CH3CN

A canned heat product used to warm buffet dishes consists of a homogeneous mixture of ethanol 1C2H5OH2 and paraffin, which has an average formula of C24H50. What mass of C2H5OH should be added to 620 kg of the paraffin to produce 8 torr of ethanol vapor pressure at 35 C? The vapor pressure of pure ethanol at 35 C is 100 torr.

A solution contains 0.115 mol H2O and an unknown number of moles of sodium chloride. The vapor pressure of the solution at 30 C is 25.7 torr. The vapor pressure of pure water at this temperature is 31.8 torr. Calculate the number of grams of sodium chloride in the solution. (Hint: Remember that sodium chloride is a strong electrolyte.)

Two beakers are placed in a sealed box at 25 C. One beaker contains 30.0 mL of a 0.050 M aqueous solution of a nonvolatile nonelectrolyte. The other beaker contains 30.0 mL of a 0.035 M aqueous solution of NaCl. The water vapor from the two solutions reaches equilibrium. (a) In which beaker does the solution level rise, and in which one does it fall? (b) What are the volumes in the two beakers when equilibrium is attained, assuming ideal behavior?

A car owner who knows no chemistry has to put antifreeze in his cars radiator. The instructions recommend a mixture of 30% ethylene glycol and 70% water. Thinking he will improve his protection he uses pure ethylene glycol, which is a liquid at room temperature. He is saddened to find that the solution does not provide as much protection as he hoped. The pure ethylene glycol freezes solid in his radiator on a very cold day, while his neighbor, who did use the 30/70 mixture, has no problem. Suggest an explanation.

Calculate the freezing point of a 0.100 m aqueous solution of K2SO4, (a) ignoring interionic attractions, and (b) taking interionic attractions into consideration by using the vant Hoff factor (Table 13.4).

Carbon disulfide 1CS22 boils at 46.30 C and has a density of 1.261 g>mL. (a) When 0.250 mol of a nondissociating solute is dissolved in 400.0 mL of CS2, the solution boils at 47.46 C. What is the molal boiling-point-elevation constant for CS2? (b) When 5.39 g of a nondissociating unknown is dissolved in 50.0 mL of CS2, the solution boils at 47.08 C. What is the molecular weight of the unknown?

A lithium salt used in lubricating grease has the formula LiCnH2n + 1O2. The salt is soluble in water to the extent of 0.036 g per 100 g of water at 25 C. The osmotic pressure of this solution is found to be 57.1 torr. Assuming that molality and molarity in such a dilute solution are the same and that the lithium salt is completely dissociated in the solution, determine an appropriate value of n in the formula for the salt.

Fluorocarbons (compounds that contain both carbon and fluorine) were, until recently, used as refrigerants. The compounds listed in the following table are all gases at 25 C, and their solubilities in water at 25 C and 1 atm fluorocarbon pressure are given as mass percentages. (a) For each fluorocarbon, calculate the molality of a saturated solution. (b) Explain why the molarity of each of the solutions should be very close numerically to the molality. (c) Based on their molecular structures, account for the differences in solubility of the four fluorocarbons. (d) Calculate the Henrys law constant at 25 C for CHClF2, and compare its magnitude to that for N2 16.8 * 10-4 mol>L@atm2. Suggest a reason for the difference in magnitude. Fluorocarbon Solubility (mass %) CF4 0.0015 CClF3 0.009 CCl2F2 0.028 CHClF2 0.30

At ordinary body temperature 137 C2, the solubility of N2 in water at ordinary atmospheric pressure (1.0 atm) is 0.015 g>L. Air is approximately 78 mol % N2. (a) Calculate the number of moles of N2 dissolved per liter of blood, assuming blood is a simple aqueous solution. (b) At a depth of 100 ft in water, the external pressure is 4.0 atm. What is the solubility of N2 from air in blood at this pressure? (c) If a scuba diver suddenly surfaces from this depth, how many milliliters of N2 gas, in the form of tiny bubbles, are released into the bloodstream from each liter of blood?

Consider the following values for enthalpy of vaporization 1kJ>mol2 of several organic substances: CH3C H 30.4 O Acetaldehyde H2C CH2 28.5 O Ethylene oxide CH3CCH3 32.0 O Acetone H2C CH2 CH2 24.7 Cyclopropane (a) Account for the variations in heats of vaporization for these substances, considering their relative intermolecular forces. (b) How would you expect the solubilities of these substances to vary in hexane as solvent? In ethanol? Use intermolecular forces, including hydrogen- bonding interactions where applicable, to explain your responses.

A textbook on chemical thermodynamics states, The heat of solution represents the difference between the lattice energy of the crystalline solid and the solvation energy of the gaseous ions. (a) Draw a simple energy diagram to illustrate this statement. (b) A salt such as NaBr is insoluble in most polar nonaqueous solvents such as acetonitrile 1CH3CN2 or nitromethane 1CH3NO22, but salts of large cations, such as tetramethylammonium bromide 31CH324NBr4, are generally more soluble. Use the thermochemical cycle you drew in part (a) and the factors that determine the lattice energy (Section 8.2) to explain this fact.

(a) A sample of hydrogen gas is generated in a closed container by reacting 2.050 g of zinc metal with 15.0 mL of 1.00 M sulfuric acid. Write the balanced equation for the reaction, and calculate the number of moles of hydrogen formed, assuming that the reaction is complete.

(b) The volume over the solution in the container is 122 mL. Calculate the partial pressure of the hydrogen gas in this volume at 25 C, ignoring any solubility of the gas in the solution. (c) The Henrys law constant for hydrogen in water at 25 C is 7.8 * 10-4 mol>L@atm. Estimate the number of moles of hydrogen gas that remain dissolved in the solution. What fraction of the gas molecules in the system is dissolved in the solution? Was it reasonable to ignore any dissolved hydrogen in part (b)?

The following table presents the solubilities of several gases in water at 25 C under a total pressure of gas and water vapor of 1 atm. (a) What volume of CH41g2 under standard conditions of temperature and pressure is contained in 4.0 L of a saturated solution at 25 C? (b) Explain the variation in solubility among the hydrocarbons listed (the first three compounds), based on their molecular structures and intermolecular forces. (c) Compare the solubilities of O2, N2, and NO, and account for the variations based on molecular structures and intermolecular forces. (d) Account for the much larger values observed for H2S and SO2 as compared with the other gases listed.(e) Find several pairs of substances with the same or nearly the same molecular masses (for example, C2H4 and N2), and use intermolecular interactions to explain the differences in their solubilities. Gas Solubility (mM) CH4 1methane2 1.3 C2H6 1ethane2 1.8 C2H4 1ethylene2 4.7 N2 0.6 O2 1.2 NO 1.9 H2S 99 SO2 1476

A small cube of lithium 1density = 0.535 g/cm3 2 measuring 1.0 mm on each edge is added to 0.500 L of water. The following reaction occurs: 2 Li1s2 + 2 H2O1l2 2 LiOH1aq2 + H21g2 What is the freezing point of the resultant solution, assuming that the reaction goes to completion?

At 35 C the vapor pressure of acetone, 1CH322CO, is 360 torr, and that of chloroform, CHCl3, is 300 torr. Acetone and chloroform can form very weak hydrogen bonds between one another; the chlorines on the carbon give the carbon a sufficient partial positive charge to enable this behavior: CH3 CH3 C H O C Cl Cl Cl A solution composed of an equal number of moles of acetone and chloroform has a vapor pressure of 250 torr at 35 C. (a) What would be the vapor pressure of the solution if it exhibited ideal behavior? (b) Use the existence of hydrogen bonds between acetone and chloroform molecules to explain the deviation from ideal behavior. (c) Based on the behavior of the solution, predict whether the mixing of acetone and chloroform is an exothermic 1Hsoln 6 02 or endothermic 1Hsoln 7 02 process. (d) Would you expect the same vaporpressure behavior for acetone and chloromethane (CH3Cl)? Explain

Compounds like sodium stearate, called surfactants in general, can form structures known as micelles in water, once the solution concentration reaches the value known as the critical micelle concentration (cmc). Micelles contain dozens to hundreds of molecules. The cmc depends on the substance, the solvent, and the temperature. Surfactant tail Surfactant head Surfactant monomers Micelle cmc At and above the cmc, the properties of the solution vary drastically. (a) The turbidity (the amount of light scattering) of solutions increases dramatically at the cmc. Suggest an explanation. (b) The ionic conductivity of the solution dramatically changes at the cmc. Suggest an explanation. (c) Chemists have developed fluorescent dyes that glow brightly only when the dye molecules are in a hydrophobic environment. Predict how the intensity of such fluorescence would relate to the concentration of sodium stearate as the sodium stearate concentration approaches and then increases past the cmc.

Based on Figure 13.18, you might think that the reason volatile solvent molecules in a solution are less likely to escape to the gas phase, compared to the pure solvent, is because the solute molecules are physically blocking the solvent molecules from leaving at the surface. This is a common misconception. Design an experiment to test the hypothesis that solute blocking of solvent vaporization is not the reason that solutions have lower vapor pressures than pure solvents.

Visualizing Concepts Rank the contents of the following containers in order of increasing entropy: [Section 13.1]

This figure shows the interaction of a cation with surrounding water molecules. (a) Which atom of water is associated with the cation? Explain. (b) Which of the following explanations accounts for the fact that the ion-solvent interaction is greater for \(\mathrm{Li}^{+}\) than for \(\mathrm{K}^{+}\)? a. \(\mathrm{Li}^{+}\) is of lower mass than \(\mathrm{K}^{+}\). b. The ionization energy of Li is higher than that for K. c. \(\mathrm{Li}^{+}\) has a smaller ionic radius than \(\mathrm{K}^{+}\). d. Li has a lower density than K. e. Li reacts with water more slowly than K. [Section 13.1] Equation Transcription: Text Transcription: Li^+ K^+ Li^+ K^+ Li^+ K^+

Problem 3E Visualizing Concepts Consider two ionic solids, both composed of singly-charged ions, that have different lattice energies, (a) Will the solids have the same solubility in water? (b) If not, which solid will be more soluble in water, the one with the larger lattice energy or the one with the smaller lattice energy? Assume that solute-solvent interactions are the same for both solids. [Section]

Problem 4E Visualizing Concepts Are gases always miscible with each other? Explain. [Section]

Which of the following is the best representation of a saturated solution? Explain your reasoning. [Section 13.2]

The solubility of Xe in water at 1 atm pressure and \(20\ ^{\circ} \mathrm{C}\) is approximately \(5 \times 10^{-3}\ M\). (a) Compare this with the solubilities of Ar and Kr in water (Table 13.1). (b) What properties of the rare gas atoms account for the variation in solubility? [Section 13.3] Equation Transcription: Text Transcription: 20 ^{o}C 5x10^-3 M

The structures of vitamins E and \(\mathrm{B}_{6}\) are shown below. Predict which is more water soluble and which is more fat soluble. Explain. [Section 13.3] Equation Transcription: Text Transcription: B_6 B_6

Problem 8E Visualizing Concepts You take a sample of water that is at room temperature and in contact with air and put it under a vacuum. Right away, you see bubbles leave the water, but after a little while, the bubbles stop. As you keep applying the vacuum, more bubbles appear. A friend tells you that the first bubbles were water vapor, and the low pressure had reduced the boiling point of water, causing the water to boil. Another friend tells you that the first bubbles were gas molecules from the air (oxygen, nitrogen, and so forth) that were dissolved in the water. Which friend is mostly likely to be correct? What, then, is responsible for the second batch of bubbles? [Section]

The figure shows two identical volumetric flasks containing the same solution at two temperatures. 1. Does the molarity of the solution change with the change in temperature? Explain. 2. Does the molality of the solution change with the change in temperature? Explain. [Section 13.4] Equation Transcription: Text Transcription: 25 ^{o}C 55 ^{o}C

This portion of a phase diagram shows the vapor-pressure curves of a volatile solvent and of a solution of that solvent containing a nonvolatile solute. (a) Which line represents the solution? (b) What are the normal boiling points of the solvent and the solution? [Section 13.5] Equation Transcription: Text Transcription: P(atm) T(^{o}C)

Suppose you had a balloon made of some highly flexible semi-permeable membrane. The balloon is filled completely with a 0.2 M solution of some solute and is submerged in a 0.1 M solution of the same solute: Initially, the volume of solution in the balloon is 0.25 L. Assuming the volume outside the semipermeable membrane is large, as the illustration shows, what would you expect for the solution volume inside the balloon once the system has come to equilibrium through osmosis? [Section 13.5]

The molecule n-octylglucoside, shown here, is widely used in biochemical research as a nonionic detergent for “solubilizing” large hydrophobic protein molecules. What characteristics of this molecule are important for its use in this way? [Section 13.6]

Problem 13E The Solution Process (Section) Indicate whether each statement is true or false: (a) A solute will dissolve in a solvent if solute-solute interactions are stronger than solute-solvent interactions. (b) In making a solution, the enthalpy of mixing is always a positive number. (c) An increase in entropy favors mixing

Indicate whether each statement is true or false: (a) NaCl dissolves in water but not in benzene \(\left(\mathrm{C}_{6} \mathrm{H}_{6}\right)\) because benzene is denser than water. (b) NaCl dissolves in water but not in ben- zene because water has a large dipole moment and benzene has zero dipole moment. (c) NaCl dissolves in water but not in benzene because the water–ion interactions are stronger than benzene–ion interactions. Equation Transcription: Text Transcription: NaCl (C_{6}H_6) NaCl NaCl

Problem 16E The Solution Process (Section) Indicate the principal type of solute-solvent interaction in each of the following solutions and rank the solutions from weakest to strongest solute-solvent interaction: (a) KCI in water, (b) CH2CI2 in benzene (C6H6), (c) methanol (CH3OH) in water.

Problem 15E The Solution Process (Section) Indicate the type of solute-solvent interaction (Section) that should be most important in each of the following solutions: (a) CCI4 in benzene (C6H6), (b) methanol (CH30H) in water, (c) KBr in water, (d) HCI in acetonitrile (CH3CN).

Problem 17E The Solution Process (Section) An ionic compound has a very negative ?Hsoln in water. (a) Would you expect it to be very soluble or nearly insoluble in water? (b) Which term would you expect to be the largest negative number: ?Hsolvent, ?Hsolute, or ?Hmix?

Problem 18E The Solution Process (Section) When ammonium chloride dissolves in water, the solution becomes colder, (a) Is the solution process exothermic or endothermic? (b) Why does the solution form?

Problem 19E (a) In Equation, which of the enthalpy terms for dissolving an ionic solid would correspond to the lattice energy? (b) Which energy term in this equation is always exothermic? ?Hsoln = ?Hsolute + ?Hsolvent + ?Hmix

Problem 20E The Solution Process (Section) For the dissolution of LiCl in water, ?Hsoln = -37 kJ/mol. Which term would you expect to be the largest negative number: ?Hsolvent, ?Hsolute, or ?Hmix?

Two nonpolar organic liquids, hexane \(\left(\mathrm{C}_{6} \mathrm{H}_{14}\right)\) and heptane \(\left(\mathrm{C}_{7} \mathrm{H}_{16}\right)\), are mixed. (a) Do you expect \(\Delta H_{\mathrm{soln}}\) to be a large positive number, a large negative number, or close to zero? Explain. (b) Hexane and heptane are miscible with each other in all proportions. In making a solution of them, is the entropy of the system increased, decreased, or close to zero, compared to the separate pure liquids? Equation Transcription: Text Transcription: (C_{6}H_{14}) (C_{7}H_{16}) {Delta}H_{soln}

Problem 22E The Solution Process (Section) The enthalpy of solution of KBr in water is about +198 kJ/mol. Nevertheless, the solubility of KBr in water is relatively high. Why does the solution process occur even though it is endothermic?

Problem 23E Saturated Solutions; Factors Affecting Solubility (Sections) The solubility of Cr(N03)3*9 H20 in water is 208 g per 100 g of water at 15 °C. A solution of Cr(NO3)3*9 H2O in water at 35 °C is formed by dissolving 324 g in 100 g of water. When this solution is slowly cooled to 15 °C, no precipitate forms, (a) What term describes this solution? (b) What action might you take to initiate crystallization? Use molecular-level processes to explain how your suggested procedure works.

The solubility of \(\mathrm{MnSO}_{4} \cdot \mathrm{H}_{2} \mathrm{O}\) in water at \(20\ ^{\circ} \mathrm{C}\) is 70 g per 100 mL of water. (a) Is a 1.22 M solution of \(\mathrm{MnSO}_{4} \cdot \mathrm{H}_{2} \mathrm{O}\) in water at \(20\ ^{\circ} \mathrm{C}\) saturated, supersaturated, or unsaturated? (b) Given a solution of \(\mathrm{MnSO}_{4} \cdot \mathrm{H}_{2} \mathrm{O}\) of unknown concentration, what experiment could you perform to determine whether the new solution is saturated, supersaturated, or Unsaturated? Equation Transcription: Text Transcription: MnSO_{4}{cdot}H_{2}O 20 ^oC MnSO_{4}{cdot}H_{2}O 20 ^oC MnSO_{4}{cdot}H_{2}O

By referring to Figure 13.15, determine whether the addition of 40.0 g of each of the following ionic solids to 100 g of water at \(40^{\circ} \mathrm{C}\) will lead to a saturated solution: (a) \(\mathrm{NaNO}_{3}\), (b) KCl, (c) \(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\), (d) \(\mathrm{~Pb}\left(\mathrm{NO}_{3}\right)_{2}\). Equation Transcription: Text Transcription: 40 ^{o}C NaNO_3 KCl K_{2}Cr_{2}O_{7} Pb(NO_{3})_{2}

By referring to Figure 13.15, determine the mass of each of the following salts required to form a saturated solution in 250 g of water at \(30\ ^{\circ} \mathrm{C}\): (a) \(\mathrm{KClO}_{3}\), (b) \(\mathrm{~Pb}\left(\mathrm{NO}_{3}\right)_{2}\), (c) \(\mathrm{Ce}_{2}(\mathrm{SO} 4)_{3}\). Equation Transcription: Text Transcription: 30 ^{o}C KClO_{3} Pb(NO_{3})_{2} Ce_{2}(SO_{4})_{3}

Saturated Solutions; Factors Affecting Solubility (Sections) Consider water and glycerol, . (a) Would you expect them to be miscible in all proportions? Explain, (b) List the intermolecular attractions that occur between a water molecule and a glycerol molecule.

Saturated Solutions; Factors Affecting Solubility (Sections) Oil and water are immiscible. Which is the most likely reason? (a) Oil molecules are denser than water, (b) Oil molecules are composed mostly of carbon and hydrogen, (c) Oil molecules have higher molar masses than water, (d) Oil molecules have higher vapor pressures than water, (e) Oil molecules have higher boiling points than water.

Problem 29E Saturated Solutions; Factors Affecting Solubility (Sections) Common laboratory solvents include acetone (CH3COCH3), methanol (CH3OH), toluene (C6H5CH3), and water. Which of these is the best solvent for nonpolar solutes?

Would you expect alanine (an amino acid) to be more soluble in water or in hexane? Explain.

(a) Would you expect stearic acid, \(\mathrm{CH}_{3}\left(\mathrm{CH}_{2}\right)_{16} \mathrm{COOH}\), to be more soluble in water or in carbon tetrachloride? Explain. (b) Which would you expect to be more soluble in water, cyclohexane or dioxane? Explain. Equation Transcription: Text Transcription: CH_{3}(CH_{2})_{16}COOH H_{2}C H_{2}C CH_{2} CH_{2} CH_{2} H_{2}C H_{2}C CH_{2} CH_{2} CH_{2}

Ibuprofen, widely used as a pain reliever, has a limited solubility in water, less than 1 mg>mL. Which part of the molecule’s structure (gray, white, red) contributes to its water solubility? Which part of the molecule (gray, white, red) contributes to its water insolubility?

Saturated Solutions; Factors Affecting Solubility (Sections) Which of the following in each pair is likely to be more soluble in hexane, : (a) , (b) benzene or glycerol, , (c) octanoic acid, , or acetic acid, ? Explain your answer in each case.

Writing Condensed Structural Formulas Write the condensed structural formula for 3-ethyl-2-methylpentane How many hydrogen atoms are in 2.2-dimethylhexane? (a) 6. (b) 8. (c) 16. (d) 18. (e) 20.

Positron emission from , occurs with release of per mole of 11C. What is the mass change per mole of 11C in this nuclear reaction? The masses of are 11.009305 and 11.011434 arnu, respectively.

Drawing Isomers Draw all the structural and geometric isomers of pentene. . It has an unbranched hydrocarbon chain. Which compound does not exist? (a) 1,2.3.4.5.6.7-octaheptaene. (b) cis-2-butane. (c) trans-3-hexene. (d) 1-propene. (e)cis-4-decene.

Drawing Isomers How many straight-chain isomers are there of hexene. ?

The partial pressure of \(\mathrm{O}_{2}\) in air at sea level is 0.21 atm. Using the data in Table 13.1, together with Henry’s law, calculate the molar concentration of \(\mathrm{O}_{2}\) in the surface water of a mountain lake saturated with air at \(20\ ^{\circ} \mathrm{C}\) and an atmospheric pressure of 650 torr. Equation Transcription: Text Transcription: O_2 O_2 20 ^{o}C

Draw the condensed Structural formula for 4-methyl-2-pentyne.

What product is formed from the hydrogenation of 2-methylpropene? (a) propane, (b) butane, (c) 2-methylbutane. (d) 2-methylpropane, (e) 2-methylpropyne

Problem 41E Concentrations of Solutions (Section) A solution is made containing 14.6 g of CH30H in 184 g of H2O. Calculate (a) the mole fraction of CH3OH, (b) the mass percent of CH3OH, (c) the molality of CH3OH.

Concentrations of Solutions (Section) A solution is made containing 20.8 g of phenol in 425 g of ethanol . Calculate (a) the mole fraction of phenol, (b) the mass percent of phenol, (c) the molality of phenol.

Problem 43E Concentrations of Solutions (Section) Calculate the molarity of the following aqueous solutions: (a) 0.540 g of Mg(N03)2 in 250.0 ml of solution, (b) 22.4 g of LiCIO4.3 H2O in 125 mL of solution, (c) 25.0 ml of 3.50 M HNO3 diluted to 0.250 L.

Problem 44E Concentrations of Solutions (Section) What is the molarity of each of the following solutions: (a) 15.0 g of AI2(SO4)3 in 0.250 ml solution, (b) 5.25 g of Mn(NO3)2*2 H20 in 175 ml of solution (c) 35.0 ml of 9.00 MH2SO4 diluted to 0.500 L?

Problem 45E Concentrations of Solutions (Section) Calculate the molality of each of the following solutions: (a) 8.66 g of benzene (C6H6) dissolved in 23.6 g of carbon tetrachloride (CCI4), (b) 4.80 g of NaCI dissolved in 0.350 L of water.

Problem 46E Concentrations of Solutions (Section) (a) What is the molality of a solution formed by dissolving 1.12 mol of KCI in 16.0 mol of water? (b) How many grams of sulfur (S8) must be dissolved in 100.0 g of naphthalene (C10H8) to make a 0.12 m solution?

Concentrations of Solutions (Section) A sulfuric acid solution containing 571.6 g of per liter of solution has a density of . Calculate (a) the mass percentage, (b) the mole fraction, (c) the molality, (d) the molarity of in this solution.

Concentrations of Solutions (Section) Ascorbic acid (vitamin C, ) is a water-soluble vitamin. A solution containing 80.5 g of ascorbic acid dissolved in 210 g of water has a density of 1.22 g/mL at 55 °C. Calculate (a) the mass percentage, (b) the mole fraction, (c) the molality, (d) the molarity of ascorbic acid in this solution.

Problem 49E Concentrations of Solutions (Section) The density of acetonitrile (CH3CN) is 0.786 g/mL and the density of methanol (CH3OH) is 0.791 g/mL. A solution is made by dissolving 22.5 mL of CH3OH in 98.7 mL of CH3CN. (a) What is the mole fraction of methanol in the solution? (b) What is the molality of the solution? (c) Assuming that the volumes are additive, what is the molarity of CH3OH in the solution?

Problem 50E Concentrations of Solutions (Section) The density of toluene (C7H8) is 0.867 g/mL, and the density of thiophene (C4H4S) is 1.065 g/ml. A solution is made by dissolving 8.10 g of thiophene in 250.0 ml of toluene, (a) Calculate the mole fraction of thiophene in the solution, (b) Calculate the molality of thiophene in the solution, (c) Assuming that the volumes of the solute and solvent are additive, what is the molarity of thiophene in the solution?

Concentrations of Solutions (Section) Calculate the number of moles of solute present in each of the following aqueous solutions: (a) 600 ml of 0.250 M , (b) 86.4 g of 0.180 m , (c) 124.0 g of a solution that is 6.45% glucose by mass.

Problem 52E Concentrations of Solutions (Section) Calculate the number of moles of solute present in each of the following solutions: (a) 255 mL of 1.50 MHNO3(ag), (b) 50.0 mg of an aqueous solution that is 1.50 m NaCI, (c) 75.0 g of an aqueous solution that is 1.50% sucrose (C12H22O11) by mass.

Problem 53E Concentrations of Solutions (Section) Describe how you would prepare each of the following aqueous solutions, starting with solid KBr: (a) 0.75 L of 1.5 x 10-2 M KBr, (b) 125 g of 0.180 m KBr, (c) 1.85 L of a solution that is 12.0% KBr by mass (the density of the solution is 1.10 g/mL), (d) a 0.150 M solution of KBr that contains just enough KBr to precipitate 16.0 g of AgBr from a solution containing 0.480 mol of AgNO3.

Problem 54E Concentrations of Solutions (Section) Describe how you would prepare each of the following aqueous solutions: (a) 1.50 L of 0.110 M(NH4)2SO4 solution, starting with solid (NH4)2SO4; (b) 225 g of a solution that is 0.65 m in Na2CO3, starting with the solid solute; (c) 1.20 L of a solution that is 15.0% Pb(NO3)2 by mass (the density of the solution is 1.16 g/mL), starting with solid solute; (d) a 0.50 M solution of HCI that would just neutralize 5.5 g of Ba(OH)2 starting with 6.0 M HCI.

Problem 55E Commercial aqueous nitric acid has a density of 1.42 g/mL and is 16 M. Calculate the percent HNaO3 by mass in the solution.

Commercial concentrated aqueous ammonia is 28% by mass and has a density of 0.90 g/mL. What is the molarity of this solution?

Problem 57E Concentrations of Solutions (Section) Brass is a substitutional alloy consisting of a solution of copper and zinc. A particular sample of red brass consisting of 80.0% Cu and 20.0% Zn by mass has a density of 8750 kg/m3. (a) What is the molality of Zn in the solid solution? (b) What is the molarity of Zn in the solution?

Caffeine \(\left(\mathrm{C}_{8} \mathrm{H}_{10} \mathrm{~N}_{4} \mathrm{O}_{2}\right)\) is a stimulant found in coffee and tea. If a solution of caffeine in the solvent chloroform \(\left(\mathrm{CHCl}_{3}\right)\) has a concentration of 0.0500 m, calculate (a) the percentage of caffeine by mass, (b) the mole fraction of caffeine in the solution. Equation Transcription: Text Transcription: (C_{8}H_{10}N_{4}O_{2}) (CHCl_{3})

Problem 60E Concentrations of Solutions (Section) Breathing air that contains 4.0% by volume CO2 over time causes rapid breathing, throbbing headache, and nausea, among other symptoms. What is the concentration of CO2 in such air in terms of (a) mol percentage, (b) molarity, assuming 1 atm pressure and a body temperature of 37 °C?

Problem 59E Concentrations of Solutions (Section) During a person's typical breathing cycle, the CO2 concentration in the expired air rises to a peak of 4.6% by volume, (a) Calculate the partial pressure of the CO2 in the expired air at its peak, assuming 1 atm pressure and a body temperature of 37 °C. (b) What is the molarity of the CO2 in the expired air at its peak, assuming a body temperature of 37 °C?

Problem 61E Colligative Properties (Section) You make a solution of a nonvolatile solute with a liquid solvent. Indicate whether each of the following statements is true or false, (a) The freezing point of the solution is higher than that of the pure solvent, (b) The freezing point of the solution is lower than that of the pure solvent, (c) The boiling point of the solution is higher than that of the pure solvent, (d) The boiling point of the solution is lower than that of the pure solvent.

Problem 62E You make a solution of a nonvolatile solute with a liquid solvent. Indicate if each of the following statements is true or false, (a) The freezing point of the solution is unchanged by addition of the solvent, (b) The solid that forms as the solution freezes is nearly pure solute, (c) The freezing point of the solution is independent of the concentration of the solute, (d) The boiling point of the solution increases in proportion to the concentration of the solute, (e) At any temperature, the vapor pressure of the solvent over the solution is lower than what it would be for the pure solvent.

Problem 63E Colligative Properties (Section) Consider two solutions, one formed by adding 10 g of glucose (C6H12O6) to 1 L of water and the other formed by adding 10 g of sucrose (C12H22O11) to 1 L of water. Calculate the vapour pressure for each solution at 20 °C; the vapor pressure of pure water at this temperature is 17.5 torr.

Problem 64E Colligative Properties (Section) (a) What is an ideal solution? (b) The vapor pressure of pure water at 60 °C is 149 torr. The vapor pressure of water over a solution at 60 °C containing equal numbers of moles of water and ethylene glycol (a nonvolatile solute) is 67 torr. Is the solution ideal according to Raoult's law? Explain.

Problem 65E Colligative Properties (Section) (a) Calculate the vapor pressure of water above a solution prepared by adding 22.5 g of lactose (C12H22011) to 200.0 g of water at 338 K. (Vapor-pressure data for water are given in Appendix B.) (b) Calculate the mass of propylene glycol (C3H802) that must be added to 0.340 kg of water to reduce the vapor pressure by 2.88 torr at 40 °C.

Problem 66E Colligative Properties (Section) (a) Calculate the vapor pressure of water above a solution prepared by dissolving 28.5 g of glycerin (C3H8O3) in 125 g of water at 343 K. (The vapor pressure of water is given in Appendix B.) (b) Calculate the mass of ethylene glycol (C2H6O2) that must be added to 1.00 kg of ethanol (C2H5OH) to reduce its vapor pressure by 10.0 torr at 35 °C. The vapour pressure of pure ethanol at 35 °C is 1.00 * 102 torr.

Problem 67E Colligative Properties (Section) At 63.5 °C, the vapor pressure of H20 is 175 torr, and that of ethanol (C2H5OH) is 400 torr. A solution is made by mixing equal masses of H2O and C2H5OH. (a) What is the mole fraction of ethanol in the solution? (b) Assuming ideal-solution behavior, what is the vapor pressure of the solution at 63.5 °C? (c) What is the mole fraction of ethanol in the vapor above the solution?

Problem 68E Colligative Properties (Section) At 20 °C, the vapor pressure of benzene (C6H6) is 75 torr, and that of toluene (C7H8) is 22 torr. Assume that benzene and toluene form an ideal solution, (a) What is the composition in mole fraction of a solution that has a vapor pressure of 35 torr at 20 °C? (b) What is the mole fraction of benzene in the vapor above the solution described in part (a)?

Problem 69E Colligative Properties (Section) (a) Does a 0.10 m aqueous solution of NaCI have a higher boiling point, a lower boiling point, or the same boiling point as a 0.10 m aqueous solution of C6H12O6? (b) The experimental boiling point of the NaCI solution is lower than that calculated assuming that NaCI is completely dissociated in solution. Why is this the case?

Problem 70E Colligative Properties (Section) Arrange the following aqueous solutions, each 10% by mass in solute, in order of increasing boiling point: glucose (C6H12O6), sucrose (C12H22O11), sodium nitrate (NaNO3).

Problem 71E Colligative Properties (Section) List the following aqueous solutions in order of increasing boiling point: 0.120 m glucose, 0.050 m LiBr, 0.050 m Zn(NO3)2.

Problem 72E Colligative Properties (Section) List the following aqueous solutions in order of decreasing freezing point: 0.040 m glycerin (C3H8O3), 0.020 m KBr, 0.030 m phenol (C6H5OH).

Using data from Table 13.3, calculate the freezing and boiling points of each of the following solutions: (a) 0.22 m glycerol \(\mathrm {\left(C_{3} H_{8} O_{3}\right)}\) in ethanol, (b) 0.240 mol of naphthalene \(\mathrm {\left(C_{10} H_{8}\right)}\) in 2.45 mol of chloroform, (c) 1.50 g NaCl in 0.250 kg of water, (d) 2.04 g KBr and 4.82 g glucose \(\mathrm {\left(C_{6} H_{12} O_{6}\right)}\) in 188 g of water. Equation Transcription: Text Transcription: (C_{3}H_{8}O_{3}) (C_{10}H_{8}) (C_{6}H_{12}O_{6})

Using data from Table 13.3, calculate the freezing and boiling points of each of the following solutions: (a) 0.25 m glucose in ethanol; (b) 20.0 g of decane, \(\mathrm{C}_{10} \mathrm{H}_{22}\), in 50.0 g \(\mathrm{CHCl}_{3}\);(c) 3.50 g NaOH in 175 g of water, (d) 0.45 mol ethylene glycol and 0.15 mol KBr in 150 g \(\mathrm{H}_{2} \mathrm{O}\). Equation Transcription: Text Transcription: C_{10}H_{22} CHCl_{3} NaOH KBr H_{2}O

Problem 75E Colligative Properties (Section) How many grams of ethylene glycol (C2H602) must be added to 1.00 kg of water to produce a solution that freezes at -5.00 °C?

Problem 76E Colligative Properties (Section) What is the freezing point of an aqueous solution that boils at 105.0 °C?

Problem 78E Colligative Properties (Section) Seawater contains 3.4 g of salts for every liter of solution. Assuming that the solute consists entirely of NaCI (in fact, over 90% of the salt is indeed NaCI), calculate the osmotic pressure of seawater at 20 °C.

Problem 77E Colligative Properties (Section) What is the osmotic pressure formed by dissolving 44.2 mg of aspirin (C9H8O4) in 0.358 L of water at 25 °C?

Adrenaline is the hormone that triggers the release of extra glucose molecules in times of stress or emergency. A solution of 0.64 g of adrenaline in 36.0 g of \(\mathrm{CCl}_{4}\) elevates the boiling point by \(0.49\ { }^{\circ} \mathrm{C}\). Calculate the approximate molar mass of adrenaline from this data. Equation Transcription: Text Transcription: CCl4 0.49 ^{o}C

Problem 80E Lauryl alcohol is obtained from coconut oil and is used to make detergents. A solution of 5.00 g of lauryl alcohol in 0.100 kg of benzene freezes at 4.1 °C. What is the molar mass of lauryl alcohol from this data?

Problem 81E Lysozyme is an enzyme that breaks bacterial cell walls. A solution containing 0.150 g of this enzyme in 210 mL of solution has an osmotic pressure of 0.953 torr at 25 °C. What is the molar mass of lysozyme?

Problem 82E A dilute aqueous solution of an organic compound soluble in water is formed by dissolving 2.35 g of the compound in water to form 0.250 L of solution. The resulting solution has an osmotic pressure of 0.605 atm at 25 °C. Assuming that the organic compound is a nonelectrolyte, what is its molar mass?

Problem 83E The osmotic pressure of a 0.010 M aqueous solution of CaCI2 is found to be 0.674 atm at 25 °C. (a) Calculate the van't Hoff factor, /, for the solution, (b) How would you expect the value of / to change as the solution becomes more concentrated? Explain.

Based on the data given in Table 13.4, which solution would give the larger freezing-point lowering, a 0.030 m solution of NaCl or a 0.020 m solution of \(\mathrm{K}_{2} \mathrm{SO}_{4}\)? How do you explain the departure from ideal behavior and the differences observed between the two salts? Equation Transcription: Text Transcription: NaCl K_{2}SO_4

Problem 85E Colloids (Section) (a) Do colloids made only of gases exist? Why or why not? (b) In the 1850's, Michael Faraday prepared ruby-red colloids of gold nanoparticles in water that are still stable today. These brightly colored colloids look like solutions. What experiment(s) could you do to determine whether a given colored preparation is a solution or colloid?

Problem 86E Colloids (Section) Choose the best answer: A colloidal dispersion of one liquid in another is called (a) a gel, (b) an emulsion, (c) a foam, (d) an aerosol.

Problem 87E An ‘emulsifying agent' is a compound that helps stabilize a hydrophobic colloid in a hydrophilic solvent (or a hydrophilic colloid in a hydrophobic solvent). Which of the following choices is the best emulsifying agent? (a) CH3COOH, (b) CH3CH2CH2COOH, (c) CH3(CH2)11COOH, (d) CH3(CH2)11COONa.

Problem 88E Aerosols are important components of the atmosphere. Does the presence of aerosols in the atmosphere increase or decrease the amount of sunlight that arrives at the Earth's surface, compared to an “aerosol-free' atmosphere? Explain your reasoning.

Problem 89E Proteins can be precipitated out of aqueous solution by the addition of an electrolyte; this process is called “salting out'the protein, (a) Do you think that all proteins would be precipitated out to the same extent by the same concentration of the same electrolyte? (b) If a protein has been salted out, are the protein-protein interactions stronger or weaker than they were before the electrolyte was added? (c) A friend of yours who is taking a biochemistry class says that salting out works because the waters of hydration that surround the protein prefer to surround the electrolyte as the electrolyte is added; therefore, the protein's hydration shell is stripped away, leading to protein precipitation. Another friend of yours in the same biochemistry class says that salting out works because the incoming ions adsorb tightly to the protein, making ion pairs on the protein surface, which end up giving the protein a zero net charge in water and therefore leading to precipitation. Discuss these two hypotheses. What kind of measurements would you need to make to distinguish between these two hypotheses?

Problem 90E Explain how (a) a soap such as sodium stearate stabilizes a colloidal dispersion of oil droplets in water; (b) milk curdles upon addition of an acid.

Problem 92AE A saturated solution of sucrose (C12H22011) is made by dissolving excess table sugar in a flask of water. There are 50 g of undissolved sucrose crystals at the bottom of the flask in contact with the saturated solution. The flask is stoppered and set aside. A year later a single large crystal of mass 50 g is at the bottom of the flask. Explain how this experiment provides evidence for a dynamic equilibrium between the saturated solution and the undissolved solute.

Butylated hydroxytoluene (BHT) has the following molecular structure: It is widely used as a preservative in a variety of foods, including dried cereals. Based on its structure, would you expect BHT to be more soluble in water or in hexane \(\left(\mathrm{C}_{6} \mathrm{H}_{14}\right)\)? Explain. Equation Transcription: Text Transcription: (BHT) CH_3 CH_3 CH_3 H_{3}C-C C-CH_3 CH_3 OH CH_3 BHT (C_{6}H_{14})

Problem 93AE Most fish need at least 4 ppm dissolved 02 in water for survival, (a) What is this concentration in mol/L? (b) What partial pressure of 02 above water is needed to obtain 4 ppm 02 in water at 10 °C? (The Henry's law constant for 02 at this temperature is 1.71 * 10-3 mol/L-atm.)

Problem 94AE The presence of the radioactive gas radon (Rn) in well water presents a possible health hazard in parts of the United States, (a) Assuming that the solubility of radon in water with 1 atm pressure of the gas over the water at 30 °C is 7.27 * 10-3 M, what is the Henry's law constant for radon in water at this temperature? (b) A sample consisting of various gases contains 3.5 * 10-6 mole fraction of radon. This gas at a total pressure of 32 atm is shaken with water at 30 °C. Calculate the molar concentration of radon in the water.

Problem 95AE Glucose makes up about 0.10% by mass of human blood. Calculate this concentration in (a) ppm, (b) molality, (c) What further information would you need to determine the molarity of the solution?

Problem 96AE The concentration of gold in seawater has been reported to be between 5 ppt (parts per trillion) and 50 ppt. Assuming that seawater contains 13 ppt of gold, calculate the number of grams of gold contained in 1.0 * 103 gal of seawater.

Problem 97AE The maximum allowable concentration of lead in drinking water is 9.0 ppb. (a) Calculate the molarity of lead in a 9.0-ppb solution, (b) How many grams of lead are in a swimming pool containing 9.0 ppb lead in 60 m3 of water?

Problem 98AE Acetonitrile (CH3CN) is a polar organic solvent that dissolves a wide range of solutes, including many salts. The density of a 1.80 M LiBr solution in acetonitrile is 0.826 g/cm3. Calculate the concentration of the solution in (a) molality, (b) mole fraction of LiBr, (c) mass percentage of CH3CN.

Problem 99AE A ‘canned heat' product used to warm buffet dishes consists of a homogeneous mixture of ethanol (C2H5OH) and paraffin, which has an average formula of C24H5O. What mass of C2H5OH should be added to 620 kg of the paraffin to produce 8 torr of ethanol vapor pressure at 35 °C? The vapor pressure of pure ethanol at 35 °C is 100 torr.

Problem 100AE A solution contains 0.115 mol H2O and an unknown number of moles of sodium chloride. The vapor pressure of the solution at 30 °C is 25.7 torr. The vapor pressure of pure water at this temperature is 31.8 torr. Calculate the number of grams of sodium chloride in the solution. (Hint Remember that sodium chloride is a strong electrolyte.)

Problem 101AE Two beakers are placed in a sealed box at 25 °C. One beaker contains 30.0 ml of a 0.050 M aqueous solution of a non-volatile nonelectrolyte. The other beaker contains 30.0 ml of a 0.035 M aqueous solution of NaCI. The water vapor from the two solutions reaches equilibrium, (a) In which beaker does the solution level rise, and in which one does it fall? (b) What are the volumes in the two beakers when equilibrium is attained, assuming ideal behavior?

Problem 102AE A car owner who knows no chemistry has to put antifreeze in his car's radiator. The instructions recommend a mixture of 30% ethylene glycol and 70% water. Thinking he will improve his protection he uses pure ethylene glycol, which is a liquid at room temperature. He is saddened to find that the solution does not provide as much protection as he hoped. The pure ethylene glycol freezes solid in his radiator on a very cold day, while his neighbor, who did use the 30/70 mixture, has no problem. Suggest an explanation.

Problem 104AE Carbon disulfide (CS2) boils at 46.30 °C and has a density of 1.261 g/ml. (a) When 0.250 mol of a nondissociating solute is dissolved in 400.0 ml_ of CS2, the solution boils at 47.46 °C. What is the molal boiling-point-elevation constant for CS2? (b) When 5.39 g of a nondissociating unknown is dissolved in 50.0 ml of CS2, the solution boils at 47.08 °C. What is the molecular weight of the unknown?

Calculate the freezing point of a 0.100 m aqueous solution of \(\mathrm{K}_{2} \mathrm{SO}_{4}\), (a) ignoring interionic attractions, and (b) taking interionic attractions into consideration by using the van’t Hoff factor (Table 13.4). Equation Transcription: Text Transcription: K_{2}SO_4

Problem 105AE A lithium salt used in lubricating grease has the formula LiCnH2n + 102. The salt is soluble in water to the extent of0.036 g per 100 g of water at 25 °C. The osmotic pressure of this solution is found to be 57.1 torr. Assuming that molality and molarity in such a dilute solution are the same and that the lithium salt is completely dissociated in the solution, determine an appropriate value of n in the formula for the salt

Fluorocarbons (compounds that contain both carbon and fluorine) were, until recently, used as refrigerants. The compounds listed in the following table are all gases at \(25\ ^{\circ} \mathrm{C}\), and their solubilities in water at \(25\ ^{\circ} \mathrm{C}\) and 1 atm fluorocarbon pressure are given as mass percentages. (a) For each fluorocarbon, calculate the molality of a saturated solution. (b) Explain why the molarity of each of the solutions should be very close numerically to the molality. (c) Based on their molecular structures, account for the differences in solubility of the four fluorocarbons. (d) Calculate the Henry’s law constant at \(25\ ^{\circ} \mathrm{C}\) for \(\mathrm{CHCIF}_{2}\), and compare its magnitude to that for \(\mathrm{~N}_{2}\left(6.8 \times 10^{-4} \mathrm{~mol} / \text {L-atm }\right)\). Suggest a reason for the difference in magnitude. Fluorocarbon Solubility (mass %) \(\mathrm{CF}_{4}\) 0.0015 \(\mathrm{CClF}_{3}\) 0.009 \(\mathrm{CCl}_{2} \mathrm{~F}_{2}\) 0.028 \(\mathrm{CHClF}_{2}\) 0.30 Equation Transcription: Text Transcription: 25 ^{o}C 25 ^{o}C 25 ^{o}C CHCIF_2 N_{2}(6.8x10^-4 mol/L-atm) CF_4 CCIF_3 CCl_{2}F_2 CHClF_2

Consider the following values for enthalpy of vaporization (kJ/mol) of several organic substances: (a) Account for the variations in heats of vaporization for these substances, considering their relative intermolecular forces. (b) How would you expect the solubilities of these sub- stances to vary in hexane as solvent? In ethanol? Use inter- molecular forces, including hydrogen-bonding interactions where applicable, to explain your responses. Equation Transcription: Text Transcription: (kJ/mol) CH_{3}C-H H_{2}C-CH_{2} CH_{3}CCH_{3} CH_{2} H_{2}C-CH_{2}

Problem 109AE A textbook on chemical thermodynamics states, "The heat of solution represents the difference between the lattice energy of the crystalline solid and the solvation energy of the gaseous ions.' (a) Draw a simple energy diagram to illustrate this statement, (b) A salt such ? as NaBr is insoluble in most polar nonaqueous solvents such as acetonitrile (CH3CN) or nitromethane (CH3NO2), but salts of large cations, such as tetramethylammonium bromide [(CH3)4NBr], are generally more soluble. Use the thermochemical cycle you drewin part (a) and the factors that determine the lattice energy (Section) to explain this fact.

Problem 110AE (a) A sample of hydrogen gas is generated in a closed container by reacting 2.050 g of zinc metal with 15.0 ml of 1.00 M sulfuric acid. Write the balanced equation for the reaction, and calculate the number of moles of hydrogen formed, assuming that the reaction is complete. (b) The volume over the solution in the container is 122 ml. Calculate the partial pressure of the hydrogen gas in this volume at 25 °C, ignoring any solubility of the gas in the solution, (c) The Henry's law constant for hydrogen in water at 25 °C is 7.8 * 10-4 mol/L-atm. Estimate the number of moles of hydrogen gas that remain dissolved in the solution. What fraction of the gas molecules in the system is dissolved in the solution? Was it reasonable to ignore any dissolved hydrogen in part (b)?

Problem 107AE At ordinary body temperature (37 °C), the solubility of N2 in water at ordinary atmospheric pressure (1.0 atm) is 0.015 g/L. Air is approximately 78 mol % N2. (a) Calculate the number of moles of N2 dissolved per liter of blood, assuming blood is a simple aqueous solution, (b) At a depth of 100 ft in water, the external pressure is 4.0 atm. What is the solubility of N2 from air in blood at this pressure? (c) If a scuba diver suddenly surfaces from this depth, how many milliliters of N2 gas, in the form of tiny bubbles, are released into the bloodstream from each liter of blood?

Problem 112IE A small cube of lithium (density = 0.535 g/cm3) measuring 1.0 mm on each edge is added to 0.500 L of water. The following reaction occurs: 2 Li(s) + 2 H20(l). 2 LiOH(aq) + H2(g) What is the freezing point of the resultant solution, assuming that the reaction goes to completion?

The following table presents the solubilities of several gases in water at \(25^{\circ} \mathrm{C}\) under a total pressure of gas and water vapor of 1 atm. (a) What volume of \(\mathrm{CH}_{4}(\mathrm{~g})\) under standard conditions of temperature and pressure is contained in 4.0 L of a saturated solution at \(25^{\circ} \mathrm{C}\)? (b) Explain the variation in solubility among the hydrocarbons listed (the first three compounds), based on their molecular structures and intermolecular forces. (c) Compare the solubilities of \(\mathrm{O}_{2}, \mathrm{~N}_{2}\), and NO, and account for the variations based on molecular structures and intermolecular forces. (d) Account for the much larger values observed for \(\mathrm{H}_{2} \mathrm{S}\) and \(\mathrm{SO}_{2}\) as compared with the other gases listed. (e) Find several pairs of substances with the same or nearly the same molecular masses (for example, \(\mathrm{C}_{2} \mathrm{H}_{4}\) and \(\mathrm{N}_{2}\)), and use intermolecular interactions to explain the differences in their solubilities. Gas Solubility (mM) \(\mathrm{CH}_{4}\) (methane) 1.3 \(\mathrm{C}_{2} \mathrm{H}_{6}\) (ethane) 1.8 \(\mathrm{C}_{2} \mathrm{H}_{4}\) (ethylene) 4.7 \(\mathrm{N}_{2}\) 0.6 \(\mathrm{O}_{2}\) 1.2 NO 1.9 \(\mathrm{H}_{2} \mathrm{S}\) 99 \(\mathrm{SO}_{2}\) 1476 Equation Transcription: Text Transcription: 25 ^{o}C CH_{4}(g) 25 ^{o}C O_{2},N_{2} H_{2}S SO_{2} C_{2}H_4 N_{2} CH_4 C_{2}H_6 C_{2}H_4 N_{2} O_2 NO H_{2}S SO_{2}

At \(35\ ^{\circ} \mathrm{C}\) the vapor pressure of acetone, \(\left(\mathrm{CH}_{3}\right)_{2} \mathrm{CO}\), is 360 torr, and that of chloroform, \(\mathrm{CHCl}_{3}\), is 300 torr. Acetone and chloroform can form very weak hydrogen bonds between one another; the chlorines on the carbon give the carbon a sufficient partial positive charge to enable this behavior: A solution composed of an equal number of moles of acetone and chloroform has a vapor pressure of 250 torr at \(35\ ^{\circ} \mathrm{C}\). (a) What would be the vapor pressure of the solution if it exhibited ideal behavior? (b) Use the existence of hydrogen bonds between acetone and chloroform molecules to explain the deviation from ideal behavior. (c) Based on the behavior of the solution, predict whether the mixing of acetone and chloroform is an exothermic \(\left(\Delta H_{\mathrm{soln}}<0\right)\) or endothermic \(\left(\Delta H_{\mathrm{soln}}>0\right)\) process. (d) Would you expect the same vapor-pressure behavior for acetone and chloromethane \(\left(\mathrm{CH}_{3} \mathrm{Cl}\right)\)? Explain. Equation Transcription: Text Transcription: 35 ^{o}C (CH_{3})_{2}CO CHCl_{3} Cl Cl-C-HO=C Cl CH_{3} CH_{3} 35 ^{o}C (H_{soln}<0) (H_{soln}>0) (CH_{3}Cl)

Compounds like sodium stearate, called “surfactants” in general, can form structures known as micelles in water, once the solution concentration reaches the value known as the critical micelle concentration (cmc). Micelles contain dozens to hundreds of molecules. The cmc depends on the substance, the solvent, and the temperature. At and above the cmc, the properties of the solution vary Drastically. (a) The turbidity (the amount of light scattering) of solutions increases dramatically at the cmc. Suggest an explanation. (b) The ionic conductivity of the solution dramatically changes at the cmc. Suggest an explanation. (c) Chemists have developed fluorescent dyes that glow brightly only when the dye molecules are in a hydrophobic environment. Predict how the intensity of such fluorescence would relate to the concentration of sodium stearate as the sodium stearate concentration approaches and then increases past the cmc.