Question: Solve the combined gas law expression for

> 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

> Compare and contrast the gas, liquid, and solid states with regard to the nature of the interactions among the particles.

> 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.

> 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

> Express each of the following in units of atm: a. 128 cm Hg b. 255 torr c. 1405 mm Hg d. 303 kPa

> Describe the molecular/atomic basis of gas pressure.

> What is the trend for atom size from top to bottom down a group?

> Why is diamond used as an industrial cutting tool?

> Describe one property that is characteristic of: a. molecular solids b. metallic solids

> Distinguish between amorphous and crystalline solids.

> Which of the following pairs of atoms and ions are isoelectronic? a. Cl-, Ar b. Na+, Ne c. Mg2+, Na+ d. Li+, Ne e. O2-, F- f. N3-, Cl-

> Predict the compound expected to have the greatest surface tension in the liquid state ннн H ннн H H H Н—С—С—С—н н—С —С—С—Н н—с—с—С—Н H нн H он он propane isopropyl alcohol propylene glycol

> Which of these molecules would you expect to have the highest boiling point? Why? H H. Н—С — Н H-C-CI H-C-0--H H H methane chloromethane methanol

> Which of these molecules exhibit dipole-dipole forces? Why? H H. Н—С — Н H-C-CI H-C-0--H H H methane chloromethane methanol

> Distinguish between the terms evaporation and boiling.

> How can the periodic table be used to determine the number of valence electrons in a representative element atom?

> Convert 3.0 m to: a. yd b. in c. ft d. cm e. mm

> Identify the element for each of the orbital diagrams (once corrected) in Question 2.91. Question 2.91: а. 1s 2s 2p b. 1s 2s 2p C. 1s 2s 2p

> Ozone, O3, has two resonance forms. Draw each form.

> Using the periodic table, write the electron configuration of each of the following atoms: a. Ca b. Fe c. Cl

> Ethylamine is an example of an important class of organic compounds. The molecular formula of ethylamine is CH3CH2NH2. Draw its Lewis structure.

> What is the role of coenzyme A in the reaction catalyzed by pyruvate dehydrogenase?

> Formaldehyde, H2CO, in water solution has been used as a preservative for biological specimens. Draw the Lewis structure of formaldehyde.

> How is a 2s orbital different from a 1s orbital?

> Refer to the periodic table, and find the following information: a. the symbol of the noble gas in period 3 b. the element in Group IVA (or 14) with the smallest mass c. the only metalloid in Group IIIA (or 13) d. the element whose atoms contain 18 proto

> True or false? Molecules that have only polar bonds will always be polar. Explain your reasoning.

> What is the bond angle of a tetrahedral molecule?

> For each of the elements Ca, K, Cu, Zn, Br, and Kr, provide the following information: a. Which are metals? b. Which are representative metals? c. Which are inert or noble gases?

> Which group of the periodic table is known as the noble gases? List their names.

> Rank the following in order of increasing bond length: single bond, double bond, triple bond

> Use LeChatelier’s principle to predict the effects, if any, of each of the following changes on the equilibrium system, described below, in a closed container. C(s) + 2H2 (g) ↽−−−−⇀ CH4 (g) + 18 kcal a. C is added. b. H2 is added. c. CH4 is removed. d.

> Calculate the pH of a solution that is: a. 1.0 × 10-1 M in HCl b. 1.0 × 10-5 M in HNO3

> Express each of the following in units of psi: a. 12.5 cm Hg b. 46.0 torr c. 254 mm Hg d. 0.48 atm

> A 0.0500-mol sample of a nutrient substance is burned in a bomb calorimeter containing 2.00 × 102 g H2O. If the formula weight of this nutrient substance is 114 g/mol, what is the fuel value (in nutritional Cal) if the temperature of the water increased

> A 0.325-mol sample of ammonium nitrate was dissolved in water producing a 4.00 × 102 g solution. The temperature decreased from 25.08C to 15.70C. If the specific heat of the resulting solution is 1.00 cal/g · 0C, calculate the quantity of energy absorbed

> What is plane-polarized light?