Question: Compare and contrast the gas, liquid, and

> Predict whether a reaction with a negative H and a negative S will be spontaneous, nonspontaneous, or temperature dependent. Explain your reasoning.

> Energy is required to break chemical bonds during the course of a reaction. When is energy released?

> The Henry’s law constant, k, for N2 in aqueous solution is 6.1 × 10-4 mol/ (L · atm) at 250C. When a diver is at a depth of 240 m, the pressure is approximately 25 atm. Calculate the equilibrium solubility of N2 at this depth (250C) in units of mol/L.

> Explain what is meant by the term entropy.

> Write the expression for free energy. When G is a negative value, what does it indicate about the spontaneity of the reaction?

> Provide an explanation for the fact that most decomposition reactions are endothermic but most combination reactions are exothermic.

> Describe what is meant by an endothermic reaction.

> What type of bond is the peptide bond? Explain why the peptide bond is rigid.

> Are the following processes exothermic or endothermic? a. S(s) + O2(g) −−−−→ SO2(g), H = -71 kcal b. N2(g) + 2O2(g) + 16.2 kcal −−−−→ 2NO2(g)

> A certain change in reaction conditions for an equilibrium process was found to increase the rate of the forward reaction much more than that of the reverse reaction. Did the amount of product increase, decrease, or remain the same? Explain your reasonin

> Bacterial growth decreases markedly in a refrigerator. Why?

> Describe conditions that can lead to dangerously low concentrations of potassium in the blood.

> Suggest a change in experimental conditions that would increase the yield of H2 in the reaction in Question 7.104. Question 7.104: C(s) + H2O(g)↽−−−−⇀CO(g) + H2 (g)

> a. Write the equilibrium constant expression for the reaction: C(s) + H2O(g)↽−−−−⇀CO(g) + H2 (g) b. Calculate the equilibrium constant if [H2O] = 0.40 M [CO] = 0.40 M [H2] = 0.20 M

> Carbonated beverages quickly go flat (lose CO2) when heated. Explain, using LeChatelier’s principle.

> Distinguish between the terms net energy and activation energy.

> How can the octet rule be used to determine the number of electrons gained or lost by an atom as it becomes an ion?

> Name the two most important cations in biological fluids.

> Explain how the pH of blood would change under each of the conditions described in Questions 8.15 and 8.16. Questions 8.15: Explain how the molar concentration of H2CO3 in the blood would change if the partial pressure of CO2 in the lungs were to incre

> label solution as isotonic, hypotonic, or hypertonic in comparison to 0.9% (m/V) NaCl (0.15 M NaCl). 3% (m/V) NaCl

> label solution as isotonic, hypotonic, or hypertonic in comparison to 0.9% (m/V) NaCl (0.15 M NaCl). 0.35 M glucose

> A gas mixture has three components, N2, F2, and He. The partial pressure of N2 is 0.35 atm and F2 is 0.45 atm. If the total pressure is 1.20 atm, what is the partial pressure of helium?

> Determine the osmolarity of 2.5 × 10-4 M C6H12O6 (nonelectrolyte).

> Two solutions, A and B, are separated by a semipermeable membrane. For each case, predict whether there will be a net flow of water in one direction and, if so, which direction. A is 0.10 M NaCl and B is 0.20 M glucose.

> Two solutions, A and B, are separated by a semipermeable membrane. For each case, predict whether there will be a net flow of water in one direction and, if so, which direction. A is 0.10 M glucose and B is 0.10 M KCl.

> Calculate the number of g of silver nitrate required to prepare 2.00 L of 0.500 M AgNO3.

> A sample of O2 gas occupies 257 mL at 208C and 1.20 atm. What is the volume of this O2 gas sample at STP?

> Using salt to try to melt ice on a day when the temperature is -200C will be unsuccessful. Why?

> What is the major importance of Raoult’s law?

> Describe the transport of lipids digested in the lumen of the intestines to the cells of the body.

> A 300.0-mL portion of H2O is added to 300.0 mL of 0.250 M H2SO4. What is the new molarity?

> After the soda in Question 6.5 is opened, the “fizz” shows a loss of CO2. If the partial pressure of CO2 in the atmosphere is 5.0 × 10-4 atm, calculate the equilibrium CO2 concentration in mol/L in the open bottle of soda. Question 6.5: An unopened bot

> Calculate the volume of a 1.00 × 10-2 M KOH solution containing 3.00 × 10-1 mol of solute.

> Will the volume of gas increase, decrease, or remain the same if the temperature is decreased and the pressure is increased? Explain.

> Write the dilution expression and define each term.

> Calculate the molarity of a solution that contains 1.75 mol of KNO3 dissolved in 3.00 L.

> Calculate the molarity of 2.75 L of solution containing 1.35 × 10-2 mol HCl.

> How many g of solute are needed to prepare each of the following solutions? a. 2.50 × 102 g of 5.00% (m/m) NH4Cl (ammonium chloride) b. 2.50 × 102 g of 3.50% (m/m) Na2CO3

> How many mL of 4.0 mass/volume % Mg (NO3)2 solution would contain 1.2 g of magnesium nitrate?

> Would the soft drink in Question 6.3 go “flat” faster if the bottle warmed to room temperature? Why? Question 6.3: Explain why, over time, a bottle of soft drink goes “flat” after it is opened.

> Use the kinetic molecular theory to explain why aerosol cans carry instructions warning against heating or disposing of the container in a fire.

> Describe the evidence that suggests that mitochondria evolved from free-living bacteria.

> A solution was prepared by dissolving 12.4 g of NaNO3 in sufficient water to produce 95.0 mL of solution. What is the mass/volume % of this solution?

> Calculate the composition of each of the following solutions in mass/mass %: a. 1.00 g KCl in 1.00 × 102 g solution b. 50.0 g KCl in 5.00 × 102 mL solution (d = 1.14 g/mL)

> Calculate the composition of each of the following solutions in mass/volume %: a. 20.0 g benzene dissolved in 1.00 × 102 mL solution b. 20.0 g acetic acid dissolved in 2.50 L solution

> Calculate the composition of each of the following solutions in mass/volume %: a. 0.700 g KCl in 1.00 mL solution b. 95.2 g MgCl2 in 0.250 L solution

> Fish kills (the sudden death of thousands of fish) often occur during periods of prolonged elevated temperatures. Pollution is often, but not always, the cause. Suggest another reason, based on solubility trends.

> Is CH3OH more likely to form a solution in water or benzene (C6H6)? Explain your reasoning.

> An isotope of technetium, mixed with sulfur and colloidally dispersed in water, is frequently used in diagnosing various medical conditions because it is readily taken up by various tissues prior to excretion. Explain why this important mixture is not a

> What role does electronegativity play in determining the bonding between atoms in a compound?

> Describe how you would distinguish experimentally between a true solution and a colloidal dispersion.

> In what form are the vitamins riboflavin, thiamine, niacin, and pantothenic acid needed by the pyruvate dehydrogenase complex?

> Which of the following solute(s) would form an electrolytic solution in water? Explain your reasoning. a. HCl b. Na2SO4 c. Ethanol (CH3CH2OH)

> Two liters of liquid A are mixed with two liters of liquid B. The resulting volume is only 3.95 L. Explain what happened on the molecular level.

> Draw the appropriate Lewis symbol for each of the following atoms: a. Be b. B c. F d. S

> Explain the large difference in boiling point (b.p.) for the isomer’s butanol and diethyl ether. нн нн нн ||| | Н-с—с- с —с—о — н нн Н-с—с- о —с—с — Н H H нн нн H H diethyl ether b.p. %3D 34.5°С butanol b.p. = 117°C

> Describe the clinical effects of depressed concentrations of potassium ions in the blood.

> What type of solute dissolves readily in benzene (C6H6)?

> Sketch the interaction of a water molecule with a chloride ion.

> Why does water’s abnormally high boiling point help to make it a desirable solvent?

> Comparing pure water and a 0.10 m glucose solution, which has the higher boiling point?

> What is the relationship between the strength of the attractive forces in a liquid and its vapor pressure?

> Explain how the molar concentration of H2CO3 in the blood would change if the partial pressure of CO2 in the lungs were to decrease.

> Compare the strength of intermolecular forces in liquids with those in solids.

> Would CO behave more like an ideal gas at 2 atm or 20 atm? Explain your reasoning.

> H2O and CH4 are gases at 1500C. Which exhibits more ideal behavior? Why?

> Fill in the blank with the missing abbreviation and name the prefix. a. 106 m = 1 _____m b. 10-3 L = 1 _____L c. 10-9 g = 1 _____g

> State Dalton’s law in equation form.

> Calculate the volume of 6.00 mol O2 gas at 30 cm Hg and 720F.

> How many g of O2 gas occupy 10.0 L at STP?

> What molecular properties favor high surface tension?

> Calculate the volume of 44.0 g of carbon dioxide at STP.

> Calculate the pressure, in atmosphere, of 7.0 mol of carbon monoxide stored in a 30.0-L container at 650C.

> Describe the structure of the eukaryotic chromosome.

> What are the units and numerical value of standard pressure?

> Would C6H12O6 in H2O form an electrolytic solution? Why?

> State Avogadro’s law in equation form.

> Would CCl4 be expected to be a solid at room temperature? Why?

> Hydrogen sulfide (H2S) is a gas at 08C. When its temperature is decreased, does it behave more or less ideally? Explain your answer.

> A sealed balloon filled with helium gas occupies 2.50 L at 250C and 1.00 atm. When released, it rises to an altitude where the temperature is 200C and the pressure is only 0.800 atm. Calculate the new volume of the balloon.

> Solve the combined gas law expression for the final temperature.

> Name each of the following compounds: a. N2O4 b. CCl4 c. N2O5

> The balloon described in Question 5.49 was then placed in a refrigerator at 398F. Calculate its new volume. Question 5.49: A balloon containing a sample of helium gas is warmed in an oven. If the balloon measures 1.25 L at room temperature (200C), what

> Determine the change in volume that takes place when a 2.00-L sample of N2(g) is heated from 250 K to 500 K.

> Describe how you would distinguish experimentally between a colloidal dispersion and a suspension.

> Write the correct formula for each of the following: a. manganese (II) oxide b. manganese (III) oxide

> The temperature on a summer day may be 908F. Convert this value to Kelvin units.

> State Charles’s law in equation form.

> A balloon filled with helium gas at 1.00 atm occupies 15.6 L. What volume would the balloon occupy in the upper atmosphere at a pressure of 0.150 atm?

> At what temperature will 2.00 mol of He fill a 2.00-L container at standard pressure?

> Calculate the Boyle’s law constant at a pressure of 2 atm.

> What is the volume of the gas at a pressure of 5 atm?

> By what factor will the volume of the gas in Question 5.33 change? Question 5.33: The pressure on a fixed mass of a gas is tripled at constant temperature. Will the volume increase, decrease, or remain the same?

> State Boyle’s law in equation form.

> Name each of the following ions: a. Cu2+ b. Fe2+ c. Fe3+

> Explain the relationship you described in the answer to Question 17.24 for the effect of the number of carbon-carbon double bonds in fatty acids on their melting points. Question 17.24: As the number of carbon-carbon double bonds in fatty acids increas

> Do gases exhibit more ideal behavior at low or high temperatures? Why?

> Why do gases with lower molar masses diffuse more rapidly than gases with higher molar masses?

> Why are gas densities much lower than those of liquids or solids?

> Name each of the following compounds: a. Li2CO3 b. FeBr2 c. CuSO4