Question: The traditional method of analysis for the

> What major assumption (that was analogous to what had already been demonstrated for electromagnetic radiation) did de Broglie and Schrödinger make about the motion of tiny particles?

> How does the Bohr theory account for the observed phenomenon of the emission of discrete wavelengths of light by excited atoms?

> What does it mean to say that the hydrogen atom has discrete energy levels? How is this fact reflected in the radiation that excited hydrogen atoms emit?

> Describe briefly why the study of electromagnetic radiation has been important to our understanding of the arrangement of electrons in atoms.

> The “Chemistry in Focus” segment Atmospheric Effects discusses the greenhouse effect. How do the greenhouse gases CO2, H2O, and CH4 have an effect on the temperature of the atmosphere?

> Calculate ∆H for the reaction N2H4(l) + O2(g) N2(g) + 2H2O(l) given the following data: Equation ΔH (kI) 2NH3(g) + 3N,0(g) 4N2(g) + 3H,0(/) - 1010 N,0(g) + 3H2(g) → N2Hạ(/) + H20(1) 2NH3(g) + }02(g)→ N2Ha(/

> A swimming pool, 10.0 m by 4.0 m, is filled with water to a depth of 3.0 m at a temperature of 20.2 °C. How much energy is required to raise the temperature of the water to 24.6 °C?

> Which of the following reactions is/are endothermic? a. CO2(s) CO2(g) b. NH3(g) NH3(l) c. 2H2(g) + O2(g) 2H2O(g) d. H2O(l) H2O(s) e. Cl2(g) 2Cl(g)

> Use the average atomic masses given inside the front cover of this book to calculate the number of moles of the element present in each of the following samples. a. 4.95 g of neon b. 72.5 g of nickel c. 115 mg of silver d. 6.22 µg of uranium (µ is a

> It has been determined that the body can generate 5500 kJ of energy during one hour of strenuous exercise. Perspiration is the body’s mechanism for eliminating this heat. How many grams and how many liters of water would have to be evaporated through per

> Consider the following equations: 3A + 6B 3D ∆H = -403 kJ/mol E + 2F A ∆H = -105.2 kJ/mol C E + 3D ∆H = +64.8 kJ/mol Suppose the first equation is reversed and multiplied by 1/6 , the second and third equati

> The overall reaction in commercial heat packs can be represented as 4Fe(s) + 3O2(g) 2Fe2O3(s) ∆H = -1652 kJ a. How much heat is released when 4.00 moles of iron is reacted with excess O2? b. How much heat is released when 1.00 mole

> For each of the substances listed in Table 10.1, calculate the quantity of heat required to heat 150. g of the substance by 11.2 °C. Table 10.1 The Specific Heat Capacities of Some Common Substances Specific Heat Capacity (J/g °C) Substan

> A 25.0-g sample of pure iron at 85 °C is dropped into 75 g of water at 20. °C. What is the final temperature of the water–iron mixture?

> A 50.0-g sample of water at 100. °C is poured into a 50.0-g sample of water at 25 °C. What will be the final temperature of the water?

> What is the “greenhouse effect”? Why is a certain level of greenhouse gases beneficial, but too high a level dangerous to life on earth? What is the most common greenhouse gas?

> What does petroleum consist of? What are some “fractions” into which petroleum is refined? How are these fractions related to the sizes of the molecules involved?

> Given the following data: C2H2(g) + 5/2O2(g) 2CO2(g) + H2O(l) ∆H = -1300. kJ C(s) + O2(g) CO2(g) ∆H = -394 kJ H2(g) + 1/2O2(g) H2O(l) ∆H = -286 kJ Calculate ∆H for the reaction 2C(s)

> Given the following data: S(s) + 3/2 O2(g) SO3(g) ∆H = -395.2 kJ 2SO2(g) + O2(g) 2SO3(g) ∆H = -198.2 kJ Calculate ∆H for the reaction S(s) + O2(g) SO2(g).

> In Fig. 10.1, what kind of energy does ball A possess initially when at rest at the top of the hill? What kind of energies are involved as ball A moves down the hill? What kind of energy does ball A possess when it reaches the bottom of the hill and stop

> Given the following data: C(s) + O2(g) CO2(g) ∆H = -393 kJ 2CO(g) + O2(g) 2CO2(g) ∆H = -566 kJ Calculate ∆H for the reaction 2C(s) + O2(g) CO(g).

> Given the following hypothetical data: X(g) + Y(g) XY(g) for which ∆H = a kJ X(g) + Z(g) XZ(g) for which ∆H = b kJ Calculate ∆H for the reaction Y(g) + XZ(g) XY(g) + Z(g)

> When ethanol (grain alcohol, C2H5OH) is burned in oxygen, approximately 1360 kJ of heat energy is released per mole of ethanol. C2H5OH(l) + 3O2(g) 2CO2(g) + 3H2O(g) a. What quantity of heat is released for each gram of ethanol

> For the reaction S(s) + O2(g) SO2(g), ∆H = -296 kJ per mole of SO2 formed. a. Calculate the quantity of heat released when 1.00 g of sulfur is burned in oxygen. b. Calculate the quantity of heat released when 0.501 mole of sulfur is b

> The enthalpy change for the reaction of hydrogen gas with fluorine gas to produce hydrogen fluoride is 2542 kJ for the equation as written: H2(g) + F2(g) 2HF(g) ∆H = -542 kJ a. What is the enthalpy change per mole of hydrogen fluorid

> The “Chemistry in Focus” segment Nature Has Hot Plants discusses thermogenic, or heat-producing, plants. For some plants, enough heat is generated to increase the temperature of the blossom by 15 °C. About how much heat is required to increase the temper

> Convert the following numbers of calories or kilocalories into joules or kilojoules. a. 7845 cal b. 4.55 * 104 cal c. 62.142 kcal d. 43,024 cal

> How is the calorie defined? How does a Calorie differ from a calorie? How is the joule related to the calorie?

> How are the temperature of an object and the thermal energy of an object related?

> The production capacity for acrylonitrile (C3H3N) in the United States is over 2 billion pounds per year. Acrylonitrile, the building block for polyacrylonitrile fibers and a variety of plastics, is produced from gaseous propylene, ammonia, and oxygen:

> What is meant by a state function? Give an example.

> Sulfur dioxide gas reacts with sodium hydroxide to form sodium sulfite and water. The unbalanced chemical equation for this reaction is as follows: SO2(g) + NaOH(s) Na2SO3(s) + H2O(l) Assuming you react 38.3 g of sulfur dioxide with

> Ammonia gas reacts with sodium metal to form sodium amide (NaNH2) and hydrogen gas. The unbalanced chemical equation for this reaction is as follows: NH3(g) + Na(s) NaNH2(s) + H2(g) Assuming that you start with 32.8 g of ammonia

> Over the years, the thermite reaction has been used for welding railroad rails, in incendiary bombs, and to ignite solid fuel rocket motors. The reaction is Fe2O3(s) + 2Al(s) / 2Fe(l) + Al2O3(s) a. What mass of iron(III) oxide must be used to produce 25

> Consider the following reaction: 4NH3(g) + 5O2(g) ( 4NO(g) + 6H2O(g) a. If a container were to have only 10 molecules of O2(g) and 10 molecules of NH3(g), how many total molecules (reactant and product) would be present in the container after the above

> Hydrazine, N2H4, emits a large quantity of energy when it reacts with oxygen, which has led to hydrazine’s use as a fuel for rockets: N2H4(l) + O2(g) / N2(g) + 2H2O(g) How many moles of each of the gaseous products are produced when 20.0 g of pure hydra

> For each of the following unbalanced chemical equations, suppose 25.0 g of each reactant is taken. Show by calculation which reactant is limiting. Calculate the theoretical yield in grams of the product in boldface. a. C2H5OH(l) + O2(g)

> For each of the following unbalanced chemical equations, suppose exactly 5.0 g of each reactant is taken. Determine which reactant is limiting, and calculate what mass of each product is expected, assuming that the limiting reactant is completely consume

> The gaseous hydrocarbon acetylene, C2H2, is used in welders’ torches because of the large amount of heat released when acetylene burns with oxygen. 2C2H2(g) + 5O2(g) / 4CO2(g) + 2H2O(g) How many grams of oxygen gas are needed for the complete combustion

> When small quantities of elemental hydrogen gas are needed for laboratory work, the hydrogen is often generated by chemical reaction of a metal with acid. For example, zinc reacts with hydrochloric acid, releasing gaseous elemental hydrogen: Zn(s) + 2HC

> When elemental copper is placed in a solution of silver nitrate, the following oxidation–reduction reaction takes place, forming elemental silver: Cu(s) + 2AgNO3(aq) / Cu(NO3)2(aq) + 2Ag(s) What mass of copper is required to remove all the silver from a

> For each of the following unbalanced equations, indicate how many moles of the first product are produced if 0.625 mole of the second product forms. State clearly the mole ratio used for each conversion. a. KO2(s) + H2O(l) O2(g

> Many metals occur naturally as sulfide compounds; examples include ZnS and CoS. Air pollution often accompanies the processing of these ores, because toxic sulfur dioxide is released as the ore is converted from the sulfide to the oxide by roasting (smel

> One step in the commercial production of sulfuric acid, H2SO4, involves the conversion of sulfur dioxide, SO2, into sulfur trioxide, SO3. 2SO2(g) + O2(g) / 2SO3(g) If 150 kg of SO2 reacts completely, what mass of SO3 should result?

> For each of the following incomplete and unbalanced equations, indicate how many moles of the second reactant would be required to react completely with 0.145 mole of the first reactant. a. BaCl2(aq) + H2SO4(aq) b. AgNO3(aq) + NaCl(aq) c. Pb(NO3)2(aq

> Using the average atomic masses given inside the front cover of the text, calculate the mass in grams of each of the following samples. a. 5.0 moles of nitric acid b. 0.000305 mole of mercury c. 2.31 * 10-5 mole of potassium chromate d. 10.5 moles of

> For each of the following balanced equations, indicate how many moles of the product could be produced by complete reaction of 1.00 g of the reactant indicated in boldface. Indicate clearly the mole ratio used for the conversion. a. NH3(g) + HCl(g) / NH

> For each of the following balanced reactions, calculate how many moles of each product would be produced by complete conversion of 0.50 mole of the reactant indicated in boldface. Indicate clearly the mole ratio used for the conversion. a. 2H2O2(l) / 2H

> For each of the following reactions, give the balanced equation for the reaction and state the meaning of the equation in terms of numbers of individual molecules and in terms of moles of molecules. a. UO2(s) + 4HF(aq) / UF4(aq) + 2H2O(l) b. 2NaC2H3O2(a

> Barium chloride solutions are used in chemical analysis for the quantitative precipitation of sulfate ion from solution. Ba2+(aq) + SO42–(aq) / BaSO4(s) Suppose a solution is known to contain on the order of 150 mg of sulfate ion. What mass of barium c

> When elemental copper is strongly heated with sulfur, a mixture of CuS and Cu2S is produced, with CuS predominating. Cu(s) + S(s) / CuS(s) 2Cu(s) + S(s) / Cu2S(s) What is the theoretical yield of CuS when 31.8 g of Cu(s) is heated with 50.0 g of

> For each of the following unbalanced equations, indicate how many moles of the second reactant would be required to react exactly with 0.275 mole of the first reactant. State clearly the mole ratio used for the conversion. a. Cl2(g) + KI(aq)

> When the sugar glucose, C6H12O6, is burned in air, carbon dioxide and water vapor are produced. Write the balanced chemical equation for this process, and calculate the theoretical yield of carbon dioxide when 1.00 g of glucose is burned completely.

> One process for the commercial production of baking soda (sodium hydrogen carbonate) involves the following reaction, in which the carbon dioxide is used in its solid form (“dry ice”) both to serve as a source of reactant and to cool the reaction system

> Natural waters often contain relatively high levels of calcium ion, Ca21, and hydrogen carbonate ion (bicarbonate), HCO3-, from the leaching of minerals into the water. When such water is used commercially or in the home, heating of the water leads to th

> Solid copper can be produced by passing gaseous ammonia over solid copper(II) oxide at high temperatures. The other products of the reaction are nitrogen gas and water vapor. The balanced equation for this reaction is: 2NH3(g) + 3CuO(s)

> Although they were formerly called the inert gases, at least the heavier elements of Group 8 do form relatively stable compounds. For example, xenon combines directly with elemental fluorine at elevated temperatures in the presence of a nickel catalyst.

> Alkali metal hydroxides are sometimes used to “scrub” excess carbon dioxide from the air in closed spaces (such as submarines and spacecraft). For example, lithium hydroxide reacts with carbon dioxide according to the unbalanced chemical equation LiOH(s

> The compound sodium thiosulfate pentahydrate, Na2S2O3.5H2O, is important commercially to the photography business as “hypo,” because it has the ability to dissolve unreacted silver salts from photographic film during development. Sodium thiosulfate penta

> An air bag is deployed by utilizing the following reaction (the nitrogen gas produced inflates the air bag): 2NaN3(s) 2Na(s) + 3N2(g) If 10.5 g of NaN3 is decomposed, what theoretical mass of sodium should be produced? If only 2.84 g of

> Your text talks about several sorts of “yield” when experiments are performed in the laboratory. Students often confuse these terms. Define, compare, and contrast what are meant by theoretical yield, actual yield, and percent yield.

> Silicon carbide, SiC, is one of the hardest materials known. Surpassed in hardness only by diamond, it is sometimes known commercially as carborundum. Silicon carbide is used primarily as an abrasive for sandpaper and is manufactured by heating common sa

> For each of the following balanced chemical equations, calculate how many moles and how many grams of each product would be produced by the complete conversion of 0.50 mole of the reactant indicated in boldface. State clearly the mole ratio used for each

> Hydrogen peroxide is used as a cleaning agent in the treatment of cuts and abrasions for several reasons. It is an oxidizing agent that can directly kill many microorganisms; it decomposes upon contact with blood, releasing elemental oxygen gas (which in

> Copper(II) sulfate has been used extensively as a fungicide (kills fungus) and herbicide (kills plants). Copper(II) sulfate can be prepared in the laboratory by reaction of copper(II) oxide with sulfuric acid. The unbalanced equation is CuO(s) + H2SO4(aq

> Lead(II) carbonate, also called “white lead,” was formerly used as a pigment in white paints. However, because of its toxicity, lead can no longer be used in paints intended for residential homes. Lead(II) carbonate is prepared industrially by reaction o

> For each of the following unbalanced chemical equations, suppose that exactly 15.0 g of each reactant are taken. Using Before–Change–After (BCA) tables, determine which reactant is limiting, and calculate what mass of each product is expected. (Assume th

> For each of the following unbalanced chemical equations, suppose 1.00 g of each reactant is taken. Show by calculation which reactant is limiting. Calculate the mass of each product that is expected. a. UO2(s) + HF(aq) UF4(aq) +

> A compound with molar mass 180.1 g/mol has the following composition by mass: C ………...40.0% H ………... 6.70% O ………...53.3% Determine the empirical and molecular formulas of the compound.

> For each of the following unbalanced chemical equations, suppose that exactly 1.00 g of each reactant is taken. Determine which reactant is limiting, and calculate what mass of the product in boldface is expected (assuming that the limiting reactant is c

> For each of the following unbalanced chemical equations, suppose 10.0 g of each reactant is taken. Show by calculation which reactant is the limiting reagent. Calculate the mass of each product that is expected. a. C3H8(g) + O2(g)

> Consider samples of phosphine (PH3), water (H2O), hydrogen sulfide (H2S), and hydrogen fluoride (HF), each with a mass of 119 g. Rank the compounds from the least to the greatest number of hydrogen atoms contained in the samples.

> For each of the following unbalanced chemical equations, suppose that exactly 5.00 g of each reactant is taken. Determine which reactant is limiting, and calculate what mass of each product is expected (assuming that the limiting reactant is completely c

> For each of the following balanced chemical equations, calculate how many grams of the product(s) would be produced by complete reaction of 0.125 mole of the first reactant. a. AgNO3(aq) + LiOH(aq) AgOH(s) + LiNO3(aq) b. Al2(SO4)3(aq)

> For each of the following unbalanced reactions, suppose exactly 5.00 moles of each reactant are taken. Determine which reactant is limiting, and also determine what mass of the excess reagent will remain after the limiting reactant is consumed. For each

> Vitamin B12, cyanocobalamin, is essential for human nutrition. Its molecular formula is C63H88CoN14O14P. A lack of this vitamin in the diet can lead to anemia. Cyanocobalamin is the form of the vitamin found in vitamin supplements. a. What is the molar

> Determine the molar mass for the following compounds to four significant figures. Compound Molar Mass (g/mol) water iron(III) chloride potassium bromide ammonium nitrate sodium hydroxide

> When barium metal is heated in chlorine gas, a binary compound forms that consists of 65.95% Ba and 34.05% Cl by mass. Calculate the empirical formula of the compound.

> When 1.00 g of metallic chromium is heated with elemental chlorine gas, 3.045 g of a chromium chloride salt results. Calculate the empirical formula of the compound.

> Explain how one determines which reactant in a process is the limiting reactant. Does this depend only on the masses of the reactant present? Give an example of how to determine the limiting reactant by using a Before–Change–After (BCA) table with a bala

> When 2.004 g of calcium is heated in pure nitrogen gas, the sample gains 0.4670 g of nitrogen. Calculate the empirical formula of the calcium nitride formed.

> In the “Chemistry in Focus” segment Cars of the Future, the claim is made that the combustion of gasoline for some cars causes about 1 lb of CO2 to be produced for each mile traveled. Estimate the gas mileage of a car that produces about 1 lb of CO2 per

> A 1.2569-g sample of a new compound has been analyzed and found to contain the following masses of elements: carbon, 0.7238 g; hydrogen, 0.07088 g; nitrogen, 0.1407 g; oxygen, 0.3214 g. Calculate the empirical formula of the compound.

> Ammonium nitrate has been used as a high explosive because it is unstable and decomposes into several gaseous substances. The rapid expansion of the gaseous substances produces the explosive force. NH4NO3(s) N2(g) + O2(g) + H2O(g) C

> For each of the following unbalanced chemical equations, calculate how many moles of each product would be produced by the complete conversion of 0.125 mole of the reactant indicated in boldface. State clearly the mole ratio used for the conversion. a.

> Calculate the percent by mass of the element mentioned first in the formulas for each of the following compounds. a. sodium azide, NaN3 b. copper(II) sulfate, CuSO4 c. gold(III) chloride, AuCl3 d. silver nitrate, AgNO3 e. rubidium sulfate, Rb2SO4 f

> Although we tend to make less use of mercury these days because of the environmental problems created by its improper disposal, mercury is still an important metal because of its unusual property of existing as a liquid at room temperature. One process b

> Calculate the number of moles of hydrogen atoms present in each of the following samples. a. 2.71 g of ammonia b. 0.824 mole of water c. 6.25 mg of sulfuric acid d. 451 g of ammonium carbonate

> Calculate the number of molecules present in each of the following samples. a. 3.45 g of C6H12O6 b. 3.45 moles of C6H12O6 c. 25.0 g of ICl5 d. 1.00 g of B2H6 e. 1.05 mmol of Al(NO3)3

> Calculate the mass in grams of each of the following samples. a. 3.09 moles of ammonium carbonate b. 4.01 * 10-6 moles of sodium hydrogen carbonate c. 88.02 moles of carbon dioxide d. 1.29 mmol of silver nitrate e. 0.0024 mole of chromium(III) chlor

> If you have equal-mole samples of each of the following compounds, which compound contains the greatest number of oxygen atoms? a. magnesium nitrate b. dinitrogen pentoxide c. iron(III) phosphate d. barium oxide e. potassium acetate

> Calculate the number of moles of the indicated substance present in each of the following samples. a. 1.28 g of iron(II) sulfate b. 5.14 mg of mercury(II) iodide c. 9.21 µg of tin(IV) oxide d. 1.26 lb of cobalt(II) chloride e. 4.25 g of copper(II) ni

> Calculate the number of moles of the indicated substance present in each of the following samples. a. 21.2 g of ammonium sulfide b. 44.3 g of calcium nitrate c. 4.35 g of dichlorine monoxide d. 1.0 lb of ferric chloride e. 1.0 kg of ferric chloride

> Calculate the molar mass for each of the following substances. a. adipic acid, C6H10O4 b. caffeine, C8H10N4O2 c. eicosane, C20H42 d. cyclohexanol, C6H11OH e. vinyl acetate, C4H6O2 f. dextrose, C6H12O6

> Calculate the molar mass for each of the following substances. a. ferrous sulfate b. mercuric iodide c. stannic oxide d. cobaltous chloride e. cupric nitrate

> For each of the following balanced chemical equations, calculate how many moles of product(s) would be produced if 0.500 mole of the first reactant were to react completely. a. CO2(g) + 4H2(g) CH4(g) + 2H2O(l) b. BaCl2(aq) + 2AgNO3(

> Using the average atomic masses given inside the front cover of this text, calculate the number of atoms present in each of the following samples. a. 2.89 g of gold b. 0.000259 mole of platinum c. 0.000259 g of platinum d. 2.0 lb of magnesium e. 1.9

> Using the average atomic masses given inside the front cover of this text, calculate the mass in grams of each of the following samples. a. 5.0 moles of potassium b. 0.000305 mole of mercury c. 2.31 * 10-5 moles of manganese d. 10.5 moles of phosphor

> Using the average atomic masses given inside the front cover of this text, calculate how many moles of each element the following masses represent. a. 1.5 mg of chromium b. 2.0 * 10-3 g of strontium c. 4.84 * 104 g of boron d. 3.6 * 10-6 µg of califo

> When elemental carbon is burned in the open atmosphere, with plenty of oxygen gas present, the product is carbon dioxide. C(s) + O2(g) CO2(g) However, when the amount of oxygen present during the burning of the carbon is restricted, carbon