Answer: Starting with H2, describe how you would prepare

Without referring to Figure 22.1, state whether each of the following elements are metals, metalloids, or nonmetals: (a) Cs, (b) Ge, (c) I, (d) Kr, (e) W, (f) Ga, (g) Te, (h) Bi.

List two chemical and two physical properties that distinguish a metal from a nonmetal

Make a list of physical and chemical properties of chlorine (Cl2) and magnesium. Comment on their differences with reference to the fact that one is a metal and the other is a nonmetal

Carbon is usually classified as a nonmetal. However, the graphite used in lead pencils conducts electricity. Look at a pencil and list two nonmetallic properties of graphite

Explain why hydrogen has a unique position in the periodic table.

Describe two laboratory and two industrial preparations for hydrogen.

Hydrogen exhibits three types of bonding in its compounds. Describe each type of bonding with an example.

What are interstitial hydrides?

Give the name of (a) an ionic hydride and (b) a covalent hydride. In each case describe the preparation and give the structure of the compound.

Describe what is meant by the hydrogen economy.

Elements number 17 and 20 form compounds with hydrogen. Write the formulas for these two compounds and compare their chemical behavior in water

Give an example of hydrogen as (a) an oxidizing agent and (b) a reducing agent.

Compare the physical and chemical properties of the hydrides of each of the following elements: Na, Ca, C, N, O, Cl

Suggest a physical method that would enable you to separate hydrogen gas from neon gas.

Write a balanced equation to show the reaction between CaH2 and H2O. How many grams of CaH2 are needed to produce 26.4 L of H2 gas at 20C and 746 mmHg?

How many kilograms of water must be processed to obtain 2.0 L of D2 at 25C and 0.90 atm pressure? Assume that deuterium abundance is 0.015 percent and that recovery is 80 percent.

Predict the outcome of the following reactions: (a) CuO(s) 1 H2(g) (b) Na2O(s) 1 H2(g)

Starting with H2, describe how you would prepare (a) HCl, (b) NH3, (c) LiOH.

Give an example of a carbide and a cyanide

How are cyanide ions used in metallurgy?

Briefly discuss the preparation and properties of carbon monoxide and carbon dioxide.

What is coal?

Explain what is meant by coal gasification

Describe two chemical differences between CO and CO2.

Describe the reaction between CO2 and OH2 in terms of a Lewis acid-base reaction such as that shown on p. 705.

Draw a Lewis structure for the C2 22 ion.

Balance the following equations: (a) Be2C(s) 1 H2O(l) (b) CaC2(s) 1 H2O(l)

Unlike CaCO3, Na2CO3 does not readily yield CO2 when heated. On the other hand, NaHCO3 undergoes thermal decomposition to produce CO2 and Na2CO3. (a) Write a balanced equation for the reaction. (b) How would you test for the CO2 evolved? [Hint: Treat the gas with limewater, an aqueous solution of Ca(OH)2.]

Two solutions are labeled A and B. Solution A contains Na2CO3 and solution B contains NaHCO3. Describe how you would distinguish between the two solutions if you were provided with a MgCl2 solution. (Hint: You need to know the solubilities of MgCO3 and MgHCO3.)

Magnesium chloride is dissolved in a solution containing sodium bicarbonate. On heating, a white precipitate is formed. Explain what causes the precipitation

A few drops of concentrated ammonia solution added to a calcium bicarbonate solution cause a white precipitate to form. Write a balanced equation for the reaction.

Sodium hydroxide is hygroscopicthat is, it absorbs moisture when exposed to the atmosphere. A student placed a pellet of NaOH on a watch glass. A few days later, she noticed that the pellet was covered with a white solid. What is the identity of this solid? (Hint: Air contains CO2.)

A piece of red-hot magnesium ribbon will continue to burn in an atmosphere of CO2 even though CO2 does not support combustion. Explain.

Is carbon monoxide isoelectronic with nitrogen (N2)?

Describe a laboratory and an industrial preparation of nitrogen gas.

What is meant by nitrogen fixation? Describe a process for fixation of nitrogen on an industrial scale.

Describe an industrial preparation of phosphorus.

Why is the P4 molecule unstable?

Nitrogen can be obtained by (a) passing ammonia over red-hot copper(II) oxide and (b) heating ammonium dichromate [one of the products is Cr(III) oxide]. Write a balanced equation for each preparation.

Write balanced equations for the preparation of sodium nitrite by (a) heating sodium nitrate and (b) heating sodium nitrate with carbon

Sodium amide (NaNH2) reacts with water to produce sodium hydroxide and ammonia. Describe this reaction as a Brnsted acid-base reaction.

Write a balanced equation for the formation of urea, (NH2)2CO, from carbon dioxide and ammonia. Should the reaction be run at a high or low pressure to maximize the yield?

Some farmers feel that lightning helps produce a better crop. What is the scientific basis for this belief?

At 620 K the vapor density of ammonium chloride relative to hydrogen (H2) under the same conditions of temperature and pressure is 14.5, although, according to its formula mass, it should have a vapor density of 26.8. How would you account for this discrepancy?

Explain, giving one example in each case, why nitrous acid can act both as a reducing agent and as an oxidizing agent.

Explain why nitric acid can be reduced but not oxidized.

Write a balanced equation for each of the following processes: (a) On heating, ammonium nitrate produces nitrous oxide. (b) On heating, potassium nitrate produces potassium nitrite and oxygen gas. (c) On heating, lead nitrate produces lead(II) oxide, nitrogen dioxide (NO2), and oxygen gas.

Explain why, under normal conditions, the reaction of zinc with nitric acid does not produce hydrogen

Potassium nitrite can be produced by heating a mixture of potassium nitrate and carbon. Write a balanced equation for this reaction. Calculate the theoretical yield of KNO2 produced by heating 57.0 g of KNO3 with an excess of carbon

Predict the geometry of nitrous oxide, N2O, by the VSEPR method and draw resonance structures for the molecule. (Hint: The atoms are arranged as NNO.)

Consider the reaction N2(g) 1 O2(g) 2NO(g) Given that the G for the reaction at 298 K is 173.4 kJ/mol, calculate (a) the standard free energy of formation of NO, (b) KP for the reaction, and (c) Kc for the reaction.

From the data in Appendix 3, calculate H for the synthesis of NO (which is the first step in the manufacture of nitric acid) at 25C: 4NH3(g) 1 5O2(g) 4NO(g) 1 6H2O(l)

Explain why two N atoms can form a double bond or a triple bond, whereas two P atoms normally can form only a single bond.

When 1.645 g of white phosphorus are dissolved in 75.5 g of CS2, the solution boils at 46.709C, whereas pure CS2 boils at 46.300C. The molal boiling-point elevation constant for CS2 is 2.34C/m. Calculate the molar mass of white phosphorus and give the molecular formula

Starting with elemental phosphorus, P4, show how you would prepare phosphoric acid.

Dinitrogen pentoxide is a product of the reaction between P4O10 and HNO3. Write a balanced equation for this reaction. Calculate the theoretical yield of N2O5 if 79.4 g of P4O10 are reacted with an excess of HNO3. (Hint: One of the products is HPO3.)

Explain why (a) NH3 is more basic than PH3, (b) NH3 has a higher boiling point than PH3, (c) PCl5 exists but NCl5 does not, (d) N2 is more inert than P4.

What is the hybridization of phosphorus in the phosphonium ion, PH4 1?

Describe one industrial and one laboratory preparation of O2.

Give an account of the various kinds of oxides that exist and illustrate each type by two examples.

Hydrogen peroxide can be prepared by treating barium peroxide with sulfuric acid. Write a balanced equation for this reaction.

Describe the Frasch process for obtaining sulfur

Describe the contact process for the production of sulfuric acid.

How is hydrogen sulfide generated in the laboratory?

Draw molecular orbital energy level diagrams for O2, O2 2, and O2 22

One of the steps involved in the depletion of ozone in the stratosphere by nitric oxide may be represented as NO(g) 1 O3(g) NO2(g) 1 O2(g) From the data in Appendix 3 calculate G, KP, and Kc for the reaction at 25C.

Hydrogen peroxide is unstable and decomposes readily: 2H2O2(aq) 2H2O(l) 1 O2(g) This reaction is accelerated by light, heat, or a catalyst. (a) Explain why hydrogen peroxide sold in drugstores comes in dark bottles. (b) The concentrations of aqueous hydrogen peroxide solutions are normally expressed as percent by mass. In the decomposition of hydrogen peroxide, how many liters of oxygen gas can be produced at STP from 15.0 g of a 7.50 percent hydrogen peroxide solution?

What are the oxidation numbers of O and F in HFO?

Oxygen forms double bonds in O2, but sulfur forms single bonds in S8. Explain.

In 2008, about 48 million tons of sulfuric acid were produced in the United States. Calculate the amount of sulfur (in grams and moles) used to produce that amount of sulfuric acid.

Sulfuric acid is a dehydrating agent. Write balanced equations for the reactions between sulfuric acid and the following substances: (a) HCOOH, (b) H3PO4, (c) HNO3, (d) HClO3. (Hint: Sulfuric acid is not decomposed by the dehydrating action.)

Calculate the amount of CaCO3 (in grams) that would be required to react with 50.6 g of SO2 emitted by a power plant.

SF6 exists but OF6 does not. Explain

Explain why SCl6, SBr6, and SI6 cannot be prepared.

Compare the physical and chemical properties of H2O and H2S.

The bad smell of water containing hydrogen sulfide can be removed by the action of chlorine. The reaction is H2S(aq) 1 Cl2(aq) 2HCl(aq) 1 S(s) If the hydrogen sulfide content of contaminated water is 22 ppm by mass, calculate the amount of Cl2 (in grams) required to remove all the H2S from 2.0 3 102 gallons of water. (1 gallon 5 3.785 L.)

Describe two reactions in which sulfuric acid acts as an oxidizing agent

Concentrated sulfuric acid reacts with sodium iodide to produce molecular iodine, hydrogen sulfide, and sodium hydrogen sulfate. Write a balanced equation for the reaction.

Describe an industrial method for preparing each of the halogens.

Name the major uses of the halogens.

Metal chlorides can be prepared in a number of ways: (a) direct combination of metal and molecular chlorine, (b) reaction between metal and hydrochloric acid, (c) acid-base neutralization, (d) metal carbonate treated with hydrochloric acid, (e) precipitation reaction. Give an example for each type of preparation

Sulfuric acid is a weaker acid than hydrochloric acid. Yet hydrogen chloride is evolved when concentrated sulfuric acid is added to sodium chloride. Explain

Show that chlorine, bromine, and iodine are very much alike by giving an account of their behavior (a) with hydrogen, (b) in producing silver salts, (c) as oxidizing agents, and (d) with sodium hydroxide. (e) In what respects is fluorine not a typical halogen element?

A 375-gallon tank is filled with water containing 167 g of bromine in the form of Br2 ions. How many liters of Cl2 gas at 1.00 atm and 20C will be required to oxidize all the bromide to molecular bromine?

Draw structures for (a) (HF)2 and (b) HF2 2

Hydrogen fluoride can be prepared by the action of sulfuric acid on sodium fluoride. Explain why hydrogen bromide cannot be prepared by the action of the same acid on sodium bromide

Aqueous copper(II) sulfate solution is blue. When aqueous potassium fluoride is added to the CuSO4 solution, a green precipitate is formed. If aqueous potassium chloride is added instead, a bright-green solution is formed. Explain what happens in each case.

What volume of bromine (Br2) vapor measured at 100C and 700 mmHg pressure would be obtained if 2.00 L of dry chlorine (Cl2), measured at 15C and 760 mmHg, were absorbed by a potassium bromide solution?

Use the VSEPR method to predict the geometries of the following species: (a) I3 2, (b) SiCl4, (c) PF5, (d) SF4.

Iodine pentoxide, I2O5, is sometimes used to remove carbon monoxide from the air by forming carbon dioxide and iodine. Write a balanced equation for this reaction and identify species that are oxidized and reduced.

Write a balanced equation for each of the following reactions: (a) Heating phosphorous acid yields phosphoric acid and phosphine (PH3). (b) Lithium carbide reacts with hydrochloric acid to give lithium chloride and methane. (c) Bubbling HI gas through an aqueous solution of HNO2 yields molecular iodine and nitric oxide. (d) Hydrogen sulfide is oxidized by chlorine to give HCl and SCl2

(a) Which of the following compounds has the great est ionic character? PCl5, SiCl4, CCl4, BCl3 (b) Which of the following ions has the smallest ionic radius? F2, C42, N32, O22 (c) Which of the following atoms has the highest ionization energy? F, Cl, Br, I (d) Which of the following oxides is most acidic? H2O, SiO2, CO2

Both N2O and O2 support combustion. Suggest one physical and one chemical test to distinguish between the two gases.

What is the change in oxidation number for the following reaction? 3O2 2O3

Describe the bonding in the C2 22 ion in terms of the molecular orbital theory.

Starting with deuterium oxide (D2O), describe how you would prepare (a) NaOD, (b) DCl, (c) ND3, (d) C2D2, (e) CD4, (f) D2SO4.

Solid PCl5 exists as [PCl 4 1][PCl6 2]. Draw Lewis structures for these ions. Describe the hybridization state of the P atoms.

Consider the Frasch process. (a) How is it possible to heat water well above 100C without turning it into steam? (b) Why is water sent down the outermost pipe? (c) Why would excavating a mine and digging for sulfur be a dangerous procedure for obtaining the element?

Predict the physical and chemical properties of astatine, a radioactive element and the last member of Group 7A.

Lubricants used in watches usually consist of longchain hydrocarbons. Oxidation by air forms solid polymers that eventually destroy the effectiveness of the lubricants. It is believed that one of the initial steps in the oxidation is removal of a hydrogen atom (hydrogen abstraction). By replacing the hydrogen atoms at reactive sites with deuterium atoms, it is possible to substantially slow down the overall oxidation rate. Why? (Hint: Consider the kinetic isotope effect.)

How are lightbulbs frosted? (Hint: Consider the action of hydrofluoric acid on glass, which is made of silicon dioxide.)

Life evolves to adapt to its environment. In this respect, explain why life most frequently needs oxygen for survival, rather than the more abundant nitrogen

As mentioned in Chapter 3, ammonium nitrate is the most important nitrogen-containing fertilizer in the world. Given only air and water as starting materials and any equipment and catalyst at your disposal, describe how you would prepare ammonium nitrate. State conditions under which you can increase the yield in each step.

As we saw in Section 21.2, the reduction of iron oxides is accomplished by using carbon monoxide as a reducing agent. Starting with coke in a blast furnace, the following equilibrium plays a key role in the extraction of iron: C(s) 1 CO2(g) 2CO(g) Use the data in Appendix 3 to calculate the equilibrium constant at 25C and 1000C. Assume H and S to be independent of temperature.

Assuming ideal behavior, calculate the density of gaseous HF at its normal boiling point (19.5C). The experimentally measured density under the same conditions is 3.10 g/L. Account for the discrepancy between your calculated value and the experimental result

A 10.0-g sample of white phosphorus was burned in an excess of oxygen. The product was dissolved in enough water to make 500 mL of solution. Calculate the pH of the solution at 25C.