Key Terms, Key Equations, Summaries, and Exercises (Chapter 18)

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Key Terms, Key Equations, Summaries, and Exercises (Chapter 18)

Key Terms

acid anhydride: compound that reacts with water to form an acid or acidic solution

alkaline earth metal: any of the metals (beryllium, magnesium, calcium, strontium, barium, and radium) occupying group 2 of the periodic table; they are reactive, divalent metals that form basic oxides

allotropes: two or more forms of the same element, in the same physical state, with different chemical structures

amorphous: solid material such as a glass that does not have a regular repeating component to its three-dimensional structure; a solid but not a crystal

base anhydride: metal oxide that behaves as a base towards acids

bicarbonate anion: salt of the hydrogen carbonate ion, ${HCO}_{3}^{-}$

bismuth: heaviest member of group 15; a less reactive metal than other representative metals

borate: compound containing boron-oxygen bonds, typically with clusters or chains as a part of the chemical structure

carbonate: salt of the anion ${CO}_{3}^{2-};$ often formed by the reaction of carbon dioxide with bases

chemical reduction: method of preparing a representative metal using a reducing agent

chlor-alkali process: electrolysis process for the synthesis of chlorine and sodium hydroxide

disproportionation reaction: chemical reaction where a single reactant is simultaneously reduced and oxidized; it is both the reducing agent and the oxidizing agent

Downs cell: electrochemical cell used for the commercial preparation of metallic sodium (and chlorine) from molten sodium chloride

Frasch process: important in the mining of free sulfur from enormous underground deposits

Haber process: main industrial process used to produce ammonia from nitrogen and hydrogen; involves the use of an iron catalyst and elevated temperatures and pressures

halide: compound containing an anion of a group 17 element in the 1− oxidation state (fluoride, F⁻; chloride, Cl⁻; bromide, Br⁻; and iodide, I⁻)

Hall–Héroult cell: electrolysis apparatus used to isolate pure aluminum metal from a solution of alumina in molten cryolite

hydrogen carbonate: salt of carbonic acid, H₂CO₃ (containing the anion ${HCO}_{3}^{−})$ in which one hydrogen atom has been replaced; an acid carbonate; also known as bicarbonate ion

hydrogen halide: binary compound formed between hydrogen and the halogens: HF, HCl, HBr, and HI

hydrogen sulfate: ${HSO}_{4}^{-}$ ion

hydrogen sulfite: ${HSO}_{3}^{-}$ ion

hydrogenation: addition of hydrogen (H₂) to reduce a compound

hydroxide: compound of a metal with the hydroxide ion OH⁻ or the group −OH

interhalogen: compound formed from two or more different halogens

metal (representative): atoms of the metallic elements of groups 1, 2, 12, 13, 14, 15, and 16, which form ionic compounds by losing electrons from their outer s or p orbitals

metalloid: element that has properties that are between those of metals and nonmetals; these elements are typically semiconductors

nitrate: ${NO}_{3}^{-}$ ion; salt of nitric acid

nitrogen fixation: formation of nitrogen compounds from molecular nitrogen

Ostwald process: industrial process used to convert ammonia into nitric acid

oxide: binary compound of oxygen with another element or group, typically containing O²⁻ ions or the group –O– or =O

ozone: allotrope of oxygen; O₃

passivation: metals with a protective nonreactive film of oxide or other compound that creates a barrier for chemical reactions; physical or chemical removal of the passivating film allows the metals to demonstrate their expected chemical reactivity

peroxide: molecule containing two oxygen atoms bonded together or as the anion, $O_{2}^{2−}$

photosynthesis: process whereby light energy promotes the reaction of water and carbon dioxide to form carbohydrates and oxygen; this allows photosynthetic organisms to store energy

Pidgeon process: chemical reduction process used to produce magnesium through the thermal reaction of magnesium oxide with silicon

polymorph: variation in crystalline structure that results in different physical properties for the resulting compound

representative element: element where the s and p orbitals are filling

representative metal: metal among the representative elements

silicate: compound containing silicon-oxygen bonds, with silicate tetrahedra connected in rings, sheets, or three-dimensional networks, depending on the other elements involved in the formation of the compounds

sulfate: ${SO}_{4}^{2−}$ ion

sulfite: ${SO}_{3}^{2−}$ ion

superoxide: oxide containing the anion $O_{2}^{-}$

Summaries

18.1 Periodicity

This section focuses on the periodicity of the representative elements. These are the elements where the electrons are entering the s and p orbitals. The representative elements occur in groups 1, 2, and 12–18. These elements are representative metals, metalloids, and nonmetals. The alkali metals (group 1) are very reactive, readily form ions with a charge of 1+ to form ionic compounds that are usually soluble in water, and react vigorously with water to form hydrogen gas and a basic solution of the metal hydroxide. The outermost electrons of the alkaline earth metals (group 2) are more difficult to remove than the outer electron of the alkali metals, leading to the group 2 metals being less reactive than those in group 1. These elements easily form compounds in which the metals exhibit an oxidation state of 2+. Zinc, cadmium, and mercury (group 12) commonly exhibit the group oxidation state of 2+ (although mercury also exhibits an oxidation state of 1+ in compounds that contain ${Hg}_{2}^{2+}) .$ Aluminum, gallium, indium, and thallium (group 13) are easier to oxidize than is hydrogen. Aluminum, gallium, and indium occur with an oxidation state 3+ (however, thallium also commonly occurs as the Tl⁺ ion). Tin and lead form stable divalent cations and covalent compounds in which the metals exhibit the 4+-oxidation state.

18.2 Occurrence and Preparation of the Representative Metals

Because of their chemical reactivity, it is necessary to produce the representative metals in their pure forms by reduction from naturally occurring compounds. Electrolysis is important in the production of sodium, potassium, and aluminum. Chemical reduction is the primary method for the isolation of magnesium, zinc, and tin. Similar procedures are important for the other representative metals.

18.3 Structure and General Properties of the Metalloids

The elements boron, silicon, germanium, arsenic, antimony, and tellurium separate the metals from the nonmetals in the periodic table. These elements, called metalloids or sometimes semimetals, exhibit properties characteristic of both metals and nonmetals. The structures of these elements are similar in many ways to those of nonmetals, but the elements are electrical semiconductors.

18.4 Structure and General Properties of the Nonmetals

Nonmetals have structures that are very different from those of the metals, primarily because they have greater electronegativity and electrons that are more tightly bound to individual atoms. Most nonmetal oxides are acid anhydrides, meaning that they react with water to form acidic solutions. Molecular structures are common for most of the nonmetals, and several have multiple allotropes with varying physical properties.

18.5 Occurrence, Preparation, and Compounds of Hydrogen

Hydrogen is the most abundant element in the universe and its chemistry is truly unique. Although it has some chemical reactivity that is similar to that of the alkali metals, hydrogen has many of the same chemical properties of a nonmetal with a relatively low electronegativity. It forms ionic hydrides with active metals, covalent compounds in which it has an oxidation state of 1− with less electronegative elements, and covalent compounds in which it has an oxidation state of 1+ with more electronegative nonmetals. It reacts explosively with oxygen, fluorine, and chlorine, less readily with bromine, and much less readily with iodine, sulfur, and nitrogen. Hydrogen reduces the oxides of metals with lower reduction potentials than chromium to form the metal and water. The hydrogen halides are all acidic when dissolved in water.

18.6 Occurrence, Preparation, and Properties of Carbonates

The usual method for the preparation of the carbonates of the alkali and alkaline earth metals is by reaction of an oxide or hydroxide with carbon dioxide. Other carbonates form by precipitation. Metal carbonates or hydrogen carbonates such as limestone (CaCO₃), the antacid Tums (CaCO₃), and baking soda (NaHCO₃) are common examples. Carbonates and hydrogen carbonates decompose in the presence of acids and most decompose on heating.

18.7 Occurrence, Preparation, and Properties of Nitrogen

Nitrogen exhibits oxidation states ranging from 3− to 5+. Because of the stability of the N≡N triple bond, it requires a great deal of energy to make compounds from molecular nitrogen. Active metals such as the alkali metals and alkaline earth metals can reduce nitrogen to form metal nitrides. Nitrogen oxides and nitrogen hydrides are also important substances.

18.8 Occurrence, Preparation, and Properties of Phosphorus

Phosphorus (group 15) commonly exhibits oxidation states of 3− with active metals and of 3+ and 5+ with more electronegative nonmetals. The halogens and oxygen will oxidize phosphorus. The oxides are phosphorus(V) oxide, P₄O₁₀, and phosphorus(III) oxide, P₄O₆. The two common methods for preparing orthophosphoric acid, H₃PO₄, are either the reaction of a phosphate with sulfuric acid or the reaction of water with phosphorus(V) oxide. Orthophosphoric acid is a triprotic acid that forms three types of salts.

18.9 Occurrence, Preparation, and Compounds of Oxygen

Oxygen is one of the most reactive elements. This reactivity, coupled with its abundance, makes the chemistry of oxygen very rich and well understood.

Compounds of the representative metals with oxygen exist in three categories (1) oxides, (2) peroxides and superoxides, and (3) hydroxides. Heating the corresponding hydroxides, nitrates, or carbonates is the most common method for producing oxides. Heating the metal or metal oxide in oxygen may lead to the formation of peroxides and superoxides. The soluble oxides dissolve in water to form solutions of hydroxides. Most metals oxides are base anhydrides and react with acids. The hydroxides of the representative metals react with acids in acid-base reactions to form salts and water. The hydroxides have many commercial uses.

All nonmetals except fluorine form multiple oxides. Nearly all of the nonmetal oxides are acid anhydrides. The acidity of oxyacids requires that the hydrogen atoms bond to the oxygen atoms in the molecule rather than to the other nonmetal atom. Generally, the strength of the oxyacid increases with the number of oxygen atoms bonded to the nonmetal atom and not to a hydrogen.

18.10 Occurrence, Preparation, and Properties of Sulfur

Sulfur (group 16) reacts with almost all metals and readily forms the sulfide ion, S²⁻, in which it has as oxidation state of 2−. Sulfur reacts with most nonmetals.

18.11 Occurrence, Preparation, and Properties of Halogens

The halogens form halides with less electronegative elements. Halides of the metals vary from ionic to covalent; halides of nonmetals are covalent. Interhalogens form by the combination of two or more different halogens.

All of the representative metals react directly with elemental halogens or with solutions of the hydrohalic acids (HF, HCl, HBr, and HI) to produce representative metal halides. Other laboratory preparations involve the addition of aqueous hydrohalic acids to compounds that contain such basic anions, such as hydroxides, oxides, or carbonates.

18.12Occurrence, Preparation, and Properties of the Noble Gases

The most significant property of the noble gases (group 18) is their inactivity. They occur in low concentrations in the atmosphere. They find uses as inert atmospheres, neon signs, and as coolants. The three heaviest noble gases react with fluorine to form fluorides. The xenon fluorides are the best characterized as the starting materials for a few other noble gas compounds.

Exercises

18.1 Periodicity

1.

How do alkali metals differ from alkaline earth metals in atomic structure and general properties?

2.

Why does the reactivity of the alkali metals decrease from cesium to lithium?

3.

Predict the formulas for the nine compounds that may form when each species in column 1 of Exercise 18.3 reacts with each species in column 2.

1	2
Na	I
Sr	Se
Al	O

4.

Predict the best choice in each of the following. You may wish to review the chapter on electronic structure for relevant examples.

(a) the most metallic of the elements Al, Be, and Ba

(b) the most covalent of the compounds NaCl, CaCl₂, and BeCl₂

(c) the lowest first ionization energy among the elements Rb, K, and Li

(d) the smallest among Al, Al⁺, and Al³⁺

(e) the largest among Cs⁺, Ba²⁺, and Xe

5.

Sodium chloride and strontium chloride are both white solids. How could you distinguish one from the other?

6.

The reaction of quicklime, CaO, with water produces slaked lime, Ca(OH)₂, which is widely used in the construction industry to make mortar and plaster. The reaction of quicklime and water is highly exothermic:
$CaO (s) + H_{2} O (l) ⟶ Ca {(OH)}_{2} (s) Δ H = −350 kJ {mol}^{−1}$

(a) What is the enthalpy of reaction per gram of quicklime that reacts?

(b) How much heat, in kilojoules, is associated with the production of 1 ton of slaked lime?

7.

Write a balanced equation for the reaction of elemental strontium with each of the following:

(a) oxygen

(b) hydrogen bromide

(c) hydrogen

(d) phosphorus

(e) water

8.

How many moles of ionic species are present in 1.0 L of a solution marked 1.0 M mercury(I) nitrate?

9.

What is the mass of fish, in kilograms, that one would have to consume to obtain a fatal dose of mercury, if the fish contains 30 parts per million of mercury by weight? (Assume that all the mercury from the fish ends up as mercury(II) chloride in the body and that a fatal dose is 0.20 g of HgCl₂.) How many pounds of fish is this?

10.

The elements sodium, aluminum, and chlorine are in the same period.

(a) Which has the greatest electronegativity?

(b) Which of the atoms is smallest?

(c) Write the Lewis structure for the simplest covalent compound that can form between aluminum and chlorine.

(d) Will the oxide of each element be acidic, basic, or amphoteric?

11.

Does metallic tin react with HCl?

12.

What is tin pest, also known as tin disease?

13.

Compare the nature of the bonds in PbCl₂ to that of the bonds in PbCl₄.

14.

Is the reaction of rubidium with water more or less vigorous than that of sodium? How does the rate of reaction of magnesium compare?

18.2 Occurrence and Preparation of the Representative Metals

15.

Write an equation for the reduction of cesium chloride by elemental calcium at high temperature.

16.

Why is it necessary to keep the chlorine and sodium, resulting from the electrolysis of sodium chloride, separate during the production of sodium metal?

17.

Give balanced equations for the overall reaction in the electrolysis of molten lithium chloride and for the reactions occurring at the electrodes. You may wish to review the chapter on electrochemistry for relevant examples.

18.

The electrolysis of molten sodium chloride or of aqueous sodium chloride produces chlorine.

Calculate the mass of chlorine produced from 3.00 kg sodium chloride in each case. You may wish to review the chapter on electrochemistry for relevant examples.

19.

What mass, in grams, of hydrogen gas forms during the complete reaction of 10.01 g of calcium with water?

20.

How many grams of oxygen gas are necessary to react completely with 3.01 $\times$ 10²¹ atoms of magnesium to yield magnesium oxide?

21.

Magnesium is an active metal; it burns in the form of powder, ribbons, and filaments to provide flashes of brilliant light. Why is it possible to use magnesium in construction?

22.

Why is it possible for an active metal like aluminum to be useful as a structural metal?

23.

Describe the production of metallic aluminum by electrolytic reduction.

24.

What is the common ore of tin and how is tin separated from it?

25.

A chemist dissolves a 1.497-g sample of a type of metal (an alloy of Sn, Pb, Sb, and Cu) in nitric acid, and metastannic acid, H₂SnO₃, is precipitated. She heats the precipitate to drive off the water, which leaves 0.4909 g of tin(IV) oxide. What was the percentage of tin in the original sample?

26.

Consider the production of 100 kg of sodium metal using a current of 50,000 A, assuming a 100% yield.

(a) How long will it take to produce the 100 kg of sodium metal?

(b) What volume of chlorine at 25 °C and 1.00 atm forms?

27.

What mass of magnesium forms when 100,000 A is passed through a MgCl₂ melt for 1.00 h if the yield of magnesium is 85% of the theoretical yield?

18.3 Structure and General Properties of the Metalloids

28.

Give the hybridization of the metalloid and the molecular geometry for each of the following compounds or ions. You may wish to review the chapters on chemical bonding and advanced covalent bonding for relevant examples.

(a) GeH₄

(b) SbF₃

(c) Te(OH)₆

(d) H₂Te

(e) GeF₂

(f) TeCl₄

(g) ${SiF}_{6}^{2−}$

(h) SbCl₅

(i) TeF₆

29.

Write a Lewis structure for each of the following molecules or ions. You may wish to review the chapter on chemical bonding.

(a) H₃BPH₃

(b) ${BF}_{4}^{-}$

(c) BBr₃

(d) B(CH₃)₃

(e) B(OH)₃

30.

Describe the hybridization of boron and the molecular structure about the boron in each of the following:

(a) H₃BPH₃

(b) ${BF}_{4}^{-}$

(c) BBr₃

(d) B(CH₃)₃

(e) B(OH)₃

31.

Using only the periodic table, write the complete electron configuration for silicon, including any empty orbitals in the valence shell. You may wish to review the chapter on electronic structure.

32.

Write a Lewis structure for each of the following molecules and ions:

(a) (CH₃)₃SiH

(b) ${SiO}_{4}^{4−}$

(c) Si₂H₆

(d) Si(OH)₄

(e) ${SiF}_{6}^{2−}$

33.

Describe the hybridization of silicon and the molecular structure of the following molecules and ions:

(a) (CH₃)₃SiH

(b) ${SiO}_{4}^{4−}$

(c) Si₂H₆

(d) Si(OH)₄

(e) ${SiF}_{6}^{2−}$

34.

Describe the hybridization and the bonding of a silicon atom in elemental silicon.

35.

Classify each of the following molecules as polar or nonpolar. You may wish to review the chapter on chemical bonding.

(a) SiH₄

(b) Si₂H₆

(c) SiCl₃H

(d) SiF₄

(e) SiCl₂F₂

36.

Silicon reacts with sulfur at elevated temperatures. If 0.0923 g of silicon reacts with sulfur to give 0.3030 g of silicon sulfide, determine the empirical formula of silicon sulfide.

37.

Name each of the following compounds:

(a) TeO₂

(b) Sb₂S₃

(c) GeF₄

(d) SiH₄

(e) GeH₄

38.

Write a balanced equation for the reaction of elemental boron with each of the following (most of these reactions require high temperature):

(a) F₂

(b) O₂

(c) S

(d) Se

(e) Br₂

39.

Why is boron limited to a maximum coordination number of four in its compounds?

40.

Write a formula for each of the following compounds:

(a) silicon dioxide

(b) silicon tetraiodide

(c) silane

(d) silicon carbide

(e) magnesium silicide

41.

From the data given in Appendix I , determine the standard enthalpy change and the standard free energy change for each of the following reactions:

(a) ${BF}_{3} (g) + {3H}_{2} O (l) ⟶ B {(OH)}_{3} (s) + 3HF (g)$

(b) ${BCl}_{3} (g) + {3H}_{2} O (l) ⟶ B {(OH)}_{3} (s) + 3HCl (g)$

(c) $B_{2} H_{6} (g) + {6H}_{2} O (l) ⟶ 2B {(OH)}_{3} (s) + {6H}_{2} (g)$

42.

A hydride of silicon prepared by the reaction of Mg₂Si with acid exerted a pressure of 306 torr at 26 °C in a bulb with a volume of 57.0 mL. If the mass of the hydride was 0.0861 g, what is its molecular mass? What is the molecular formula for the hydride?

43.

Suppose you discovered a diamond completely encased in a silicate rock. How would you chemically free the diamond without harming it?

18.4 Structure and General Properties of the Nonmetals

44.

Carbon forms a number of allotropes, two of which are graphite and diamond. Silicon has a diamond structure. Why is there no allotrope of silicon with a graphite structure?

45.

Nitrogen in the atmosphere exists as very stable diatomic molecules. Why does phosphorus form less stable P₄ molecules instead of P₂ molecules?

46.

Write balanced chemical equations for the reaction of the following acid anhydrides with water:

(a) SO₃

(b) N₂O₃

(c) Cl₂O₇

(d) P₄O₁₀

(e) NO₂

47.

Determine the oxidation number of each element in each of the following compounds:

(a) HCN

(b) OF₂

(c) AsCl₃

48.

Determine the oxidation state of sulfur in each of the following:

(a) SO₃

(b) SO₂

(c) ${SO}_{3}^{2−}$

49.

Arrange the following in order of increasing electronegativity: F; Cl; O; and S.

50.

Why does white phosphorus consist of tetrahedral P₄ molecules while nitrogen consists of diatomic N₂ molecules?

18.5 Occurrence, Preparation, and Compounds of Hydrogen

51.

Why does hydrogen not exhibit an oxidation state of 1− when bonded to nonmetals?

52.

The reaction of calcium hydride, CaH₂, with water can be characterized as a Lewis acid-base reaction:
${CaH}_{2} (s) + {2H}_{2} O (l) ⟶ Ca {(OH)}_{2} (a q) + {2H}_{2} (g)$

Identify the Lewis acid and the Lewis base among the reactants. The reaction is also an oxidation-reduction reaction. Identify the oxidizing agent, the reducing agent, and the changes in oxidation number that occur in the reaction.

53.

In drawing Lewis structures, we learn that a hydrogen atom forms only one bond in a covalent compound. Why?

54.

What mass of CaH₂ is necessary to react with water to provide enough hydrogen gas to fill a balloon at 20 °C and 0.8 atm pressure with a volume of 4.5 L? The balanced equation is:
${CaH}_{2} (s) + {2H}_{2} O (l) ⟶ Ca {(OH)}_{2} (a q) + {2H}_{2} (g)$

55.

What mass of hydrogen gas results from the reaction of 8.5 g of KH with water?
$KH + H_{2} O ⟶ KOH + H_{2}$

18.6 Occurrence, Preparation, and Properties of Carbonates

56.

Carbon forms the ${CO}_{3}^{2−}$ ion, yet silicon does not form an analogous ${SiO}_{3}^{2−}$ ion. Why?

57.

Complete and balance the following chemical equations:

(a) hardening of plaster containing slaked lime
$Ca {(OH)}_{2} + {CO}_{2} ⟶$

(b) removal of sulfur dioxide from the flue gas of power plants
$CaO + {SO}_{2} ⟶$

(c) the reaction of baking powder that produces carbon dioxide gas and causes bread to rise
${NaHCO}_{3} + {NaH}_{2} {PO}_{4} ⟶$

58.

Heating a sample of Na₂CO₃⋅xH₂O weighing 4.640 g until the removal of the water of hydration leaves 1.720 g of anhydrous Na₂CO₃. What is the formula of the hydrated compound?

18.7 Occurrence, Preparation, and Properties of Nitrogen

59.

Write the Lewis structures for each of the following:

(a) NH²⁻

(b) N₂F₄

(c) ${NH}_{2}^{-}$

(d) NF₃

(e) $N_{3}^{-}$

60.

For each of the following, indicate the hybridization of the nitrogen atom (for $N_{3}^{−},$ the central nitrogen).

(a) N₂F₄

(b) ${NH}_{2}^{-}$

(c) NF₃

(d) $N_{3}^{-}$

61.

Explain how ammonia can function both as a Brønsted base and as a Lewis base.

62.

Determine the oxidation state of nitrogen in each of the following. You may wish to review the chapter on chemical bonding for relevant examples.

(a) NCl₃

(b) ClNO

(c) N₂O₅

(d) N₂O₃

(e) ${NO}_{2}^{-}$

(f) N₂O₄

(g) N₂O

(h) ${NO}_{3}^{-}$

(i) HNO₂

(j) HNO₃

63.

For each of the following, draw the Lewis structure, predict the ONO bond angle, and give the hybridization of the nitrogen. You may wish to review the chapters on chemical bonding and advanced theories of covalent bonding for relevant examples.

(a) NO₂

(b) ${NO}_{2}^{-}$

(c) ${NO}_{2}^{+}$

64.

How many grams of gaseous ammonia will the reaction of 3.0 g hydrogen gas and 3.0 g of nitrogen gas produce?

65.

Although PF₅ and AsF₅ are stable, nitrogen does not form NF₅ molecules. Explain this difference among members of the same group.

66.

The equivalence point for the titration of a 25.00-mL sample of CsOH solution with 0.1062 M HNO₃ is at 35.27 mL. What is the concentration of the CsOH solution?

18.8 Occurrence, Preparation, and Properties of Phosphorus

67.

Write the Lewis structure for each of the following. You may wish to review the chapter on chemical bonding and molecular geometry.

(a) PH₃

(b) ${PH}_{4}^{+}$

(c) P₂H₄

(d) ${PO}_{4}^{3−}$

(e) PF₅

68.

Describe the molecular structure of each of the following molecules or ions listed. You may wish to review the chapter on chemical bonding and molecular geometry.

(a) PH₃

(b) ${PH}_{4}^{+}$

(c) P₂H₄

(d) ${PO}_{4}^{3−}$

69.

Complete and balance each of the following chemical equations. (In some cases, there may be more than one correct answer.)

(a) $P_{4} + Al ⟶$

(b) $P_{4} + Na ⟶$

(c) $P_{4} + F_{2} ⟶$

(d) $P_{4} + {Cl}_{2} ⟶$

(e) $P_{4} + O_{2} ⟶$

(f) $P_{4} O_{6} + O_{2} ⟶$

70.

Describe the hybridization of phosphorus in each of the following compounds: P₄O₁₀, P₄O₆, PH₄I (an ionic compound), PBr₃, H₃PO₄, H₃PO₃, PH₃, and P₂H₄. You may wish to review the chapter on advanced theories of covalent bonding.

71.

What volume of 0.200 M NaOH is necessary to neutralize the solution produced by dissolving 2.00 g of PCl₃ is an excess of water? Note that when H₃PO₃ is titrated under these conditions, only one proton of the acid molecule reacts.

72.

How much POCl₃ can form from 25.0 g of PCl₅ and the appropriate amount of H₂O?

73.

How many tons of Ca₃(PO₄)₂ are necessary to prepare 5.0 tons of phosphorus if the yield is 90%?

74.

Write equations showing the stepwise ionization of phosphorous acid.

75.

Draw the Lewis structures and describe the geometry for the following:

(a) ${PF}_{4}^{+}$

(b) PF₅

(c) ${PF}_{6}^{-}$

(d) POF₃

76.

Why does phosphorous acid form only two series of salts, even though the molecule contains three hydrogen atoms?

77.

Assign an oxidation state to phosphorus in each of the following:

(a) NaH₂PO₃

(b) PF₅

(c) P₄O₆

(d) K₃PO₄

(e) Na₃P

(f) Na₄P₂O₇

78.

Phosphoric acid, one of the acids used in some cola drinks, is produced by the reaction of phosphorus(V) oxide, an acidic oxide, with water. Phosphorus(V) oxide is prepared by the combustion of phosphorus.

(a) Write the empirical formula of phosphorus(V) oxide.

(b) What is the molecular formula of phosphorus(V) oxide if the molar mass is about 280.

(c) Write balanced equations for the production of phosphorus(V) oxide and phosphoric acid.

(d) Determine the mass of phosphorus required to make 1.00 $\times$ 10⁴ kg of phosphoric acid, assuming a yield of 98.85%.

18.9 Occurrence, Preparation, and Compounds of Oxygen

79.

Predict the product of burning francium in air.

80.

Using equations, describe the reaction of water with potassium and with potassium oxide.

81.

Write balanced chemical equations for the following reactions:

(a) zinc metal heated in a stream of oxygen gas

(b) zinc carbonate heated until loss of mass stops

(c) zinc carbonate added to a solution of acetic acid, CH₃CO₂H

(d) zinc added to a solution of hydrobromic acid

82.

Write balanced chemical equations for the following reactions:

(a) cadmium burned in air

(b) elemental cadmium added to a solution of hydrochloric acid

(c) cadmium hydroxide added to a solution of acetic acid, CH₃CO₂H

83.

Illustrate the amphoteric nature of aluminum hydroxide by citing suitable equations.

84.

Write balanced chemical equations for the following reactions:

(a) metallic aluminum burned in air

(b) elemental aluminum heated in an atmosphere of chlorine

(c) aluminum heated in hydrogen bromide gas

(d) aluminum hydroxide added to a solution of nitric acid

85.

Write balanced chemical equations for the following reactions:

(a) sodium oxide added to water

(b) cesium carbonate added to an excess of an aqueous solution of HF

(c) aluminum oxide added to an aqueous solution of HClO₄

(d) a solution of sodium carbonate added to solution of barium nitrate

(e) titanium metal produced from the reaction of titanium tetrachloride with elemental sodium

86.

What volume of 0.250 M H₂SO₄ solution is required to neutralize a solution that contains 5.00 g of CaCO₃?

87.

Which is the stronger acid, HClO₄ or HBrO₄? Why?

88.

Write a balanced chemical equation for the reaction of an excess of oxygen with each of the following. Remember that oxygen is a strong oxidizing agent and tends to oxidize an element to its maximum oxidation state.

(a) Mg

(b) Rb

(c) Ga

(d) C₂H₂

(e) CO

89.

Which is the stronger acid, H₂SO₄ or H₂SeO₄? Why? You may wish to review the chapter on acid-base equilibria.

18.10 Occurrence, Preparation, and Properties of Sulfur

90.

Explain why hydrogen sulfide is a gas at room temperature, whereas water, which has a lower molecular mass, is a liquid.

91.

Give the hybridization and oxidation state for sulfur in SO₂, in SO₃, and in H₂SO₄.

92.

Which is the stronger acid, NaHSO₃ or NaHSO₄?

93.

Determine the oxidation state of sulfur in SF₆, SO₂F₂, and KHS.

94.

Which is a stronger acid, sulfurous acid or sulfuric acid? Why?

95.

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

96.

Give the Lewis structure of each of the following:

(a) SF₄

(b) K₂SO₄

(c) SO₂Cl₂

(d) H₂SO₃

(e) SO₃

97.

Write two balanced chemical equations in which sulfuric acid acts as an oxidizing agent.

98.

Explain why sulfuric acid, H₂SO₄, which is a covalent molecule, dissolves in water and produces a solution that contains ions.

99.

How many grams of Epsom salts (MgSO₄⋅7H₂O) will form from 5.0 kg of magnesium?

18.11 Occurrence, Preparation, and Properties of Halogens

100.

What does it mean to say that mercury(II) halides are weak electrolytes?

101.

Why is SnCl₄ not classified as a salt?

102.

The following reactions are all similar to those of the industrial chemicals. Complete and balance the equations for these reactions:

(a) reaction of a weak base and a strong acid
${NH}_{3} + {HClO}_{4} ⟶$

(b) preparation of a soluble silver salt for silver plating
${Ag}_{2} {CO}_{3} + {HNO}_{3} ⟶$

(c) preparation of strontium hydroxide by electrolysis of a solution of strontium chloride
${SrCl}_{2} (a q) + H_{2} O (l) \overset{electrolysis}{\to}$

103.

Which is the stronger acid, HClO₃ or HBrO₃? Why?

104.

What is the hybridization of iodine in IF₃ and IF₅?

105.

Predict the molecular geometries and draw Lewis structures for each of the following. You may wish to review the chapter on chemical bonding and molecular geometry.

(a) IF₅

(b) $I_{3}^{-}$

(c) PCl₅

(d) SeF₄

(e) ClF₃

106.

Which halogen has the highest ionization energy? Is this what you would predict based on what you have learned about periodic properties?

107.

Name each of the following compounds:

(a) BrF₃

(b) NaBrO₃

(c) PBr₅

(d) NaClO₄

(e) KClO

108.

Explain why, at room temperature, fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid.

109.

What is the oxidation state of the halogen in each of the following?

(a) H₅IO₆

(b) ${IO}_{4}^{-}$

(c) ClO₂

(d) ICl₃

(e) F₂

110.

Physiological saline concentration—that is, the sodium chloride concentration in our bodies—is approximately 0.16 M. A saline solution for contact lenses is prepared to match the physiological concentration. If you purchase 25 mL of contact lens saline solution, how many grams of sodium chloride have you bought?

18.12 Occurrence, Preparation, and Properties of the Noble Gases

111.

Give the hybridization of xenon in each of the following. You may wish to review the chapter on the advanced theories of covalent bonding.

(a) XeF₂

(b) XeF₄

(c) XeO₃

(d) XeO₄

(e) XeOF₄

112.

What is the molecular structure of each of the following molecules? You may wish to review the chapter on chemical bonding and molecular geometry.

(a) XeF₂

(b) XeF₄

(c) XeO₃

(d) XeO₄

(e) XeOF₄

113.

Indicate whether each of the following molecules is polar or nonpolar. You may wish to review the chapter on chemical bonding and molecular geometry.

(a) XeF₂

(b) XeF₄

(c) XeO₃

(d) XeO₄

(e) XeOF₄

114.

What is the oxidation state of the noble gas in each of the following? You may wish to review the chapter on chemical bonding and molecular geometry.

(a) XeO₂F₂

(b) KrF₂

(c) ${XeF}_{3}^{+}$

(d) ${XeO}_{6}^{4−}$

(e) XeO₃

115.

A mixture of xenon and fluorine was heated. A sample of the white solid that formed reacted with hydrogen to yield 81 mL of xenon (at STP) and hydrogen fluoride, which was collected in water, giving a solution of hydrofluoric acid. The hydrofluoric acid solution was titrated, and 68.43 mL of 0.3172 M sodium hydroxide was required to reach the equivalence point. Determine the empirical formula for the white solid and write balanced chemical equations for the reactions involving xenon.

116.

Basic solutions of Na₄XeO₆ are powerful oxidants. What mass of Mn(NO₃)₂•6H₂O reacts with 125.0 mL of a 0.1717 M basic solution of Na₄XeO₆ that contains an excess of sodium hydroxide if the products include Xe and solution of sodium permanganate?

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Key Terms

Summaries

18.1 Periodicity

18.2 Occurrence and Preparation of the Representative Metals

18.3 Structure and General Properties of the Metalloids

18.4 Structure and General Properties of the Nonmetals

18.5 Occurrence, Preparation, and Compounds of Hydrogen

18.6 Occurrence, Preparation, and Properties of Carbonates

18.7 Occurrence, Preparation, and Properties of Nitrogen

18.8 Occurrence, Preparation, and Properties of Phosphorus

18.9 Occurrence, Preparation, and Compounds of Oxygen

18.10 Occurrence, Preparation, and Properties of Sulfur

18.11 Occurrence, Preparation, and Properties of Halogens

18.12Occurrence, Preparation, and Properties of the Noble Gases

Exercises

18.1 Periodicity

18.2 Occurrence and Preparation of the Representative Metals

18.3 Structure and General Properties of the Metalloids

18.4 Structure and General Properties of the Nonmetals

18.5 Occurrence, Preparation, and Compounds of Hydrogen

18.6 Occurrence, Preparation, and Properties of Carbonates

18.7 Occurrence, Preparation, and Properties of Nitrogen

18.8 Occurrence, Preparation, and Properties of Phosphorus

18.9 Occurrence, Preparation, and Compounds of Oxygen

18.10 Occurrence, Preparation, and Properties of Sulfur

18.11 Occurrence, Preparation, and Properties of Halogens

18.12 Occurrence, Preparation, and Properties of the Noble Gases

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