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Chapter 22. Chemistry of the Nonmetals
Media Resources
Figures and Tables in Transparency Pack: Section:
Figure 22.1 Trends in Elemental Properties 22.1 Periodic Trends and Chemical Reactions
Table 22.1 Properties of Xenon Compounds 22.3 Group 8A: The Noble Gases
Table 22.2 Some Properties of the Halogens 22.4 Group 7A: The Halogens
and Pb
Activities: Section:
Carbon-Silicon Orbital Overlap 22.1 Periodic Trends and Chemical Reactions
Animations: Sections:
Periodic Trends: Acid-Base Behavior of Oxides 22.5 Oxygen
Movies: Section:
3-D Models: Section:
Carbon Dioxide 22.1 Periodic Trends and Chemical Reactions
Teflon 22.4 Group 7A: The Halogens
Ozone 22.5 Oxygen
Hydrogen Peroxide 22.5 Oxygen
Water 22.5 Oxygen
Sulfate Ion 22.6 The Other Group 6A Elements: S, Se, Te,
Chemistry of the Nonmetals
315
Phosphorous Acid 22.8 The Other Group 5A Elements: P, As, Sb,
and Bi
Other Resources
Further Readings: Section:
What’s the Use? Hydrogen 22.2 Hydrogen
Electronegativities of the Noble Gases 22.3 Group 8A: The Noble Gases
Discovery and Early Uses of Iodine 22.4 Group 7A: The Halogens
Dentifrice Fluoride 22.4 Group 7A: The Halogens
Joseph Priestley, Preeminent Amateur Chemist 22.5 Oxygen
The Discovery of Oxygen and other Priestley Matters 22.5 Oxygen
The Three Forms of Molecular Oxygen 22.5 Oxygen
An Acidity Scale for Binary Oxides 22.5 Oxygen
Herman Frasch, Sulfur King 22.6 The Other Group 6A Elements: S, Se, Te,
and Pb
Boron Clusters Come of Age 22.11 Boron
Live Demonstrations: Section:
Making Hydrogen Gas from an Acid and a Base 22.2 Hydrogen
An Overhead Demonstration of Some Descriptive 22.4 Group 7A: The Halogens
Chemistry of the Halogens and Le Châtelier’s
Chapter 22
316
Chapter 22. Chemistry of the Nonmetals
Common Student Misconceptions:
• Students often find detailed discussions of descriptive chemistry to be difficult to digest.
• Students often think that all nonmetals are gaseous.
Teaching Tips:
• Videos and similar visual aids are useful in helping students learn the descriptive chemistry in this
chapter.
Lecture Outline
22.1 Periodic Trends and Chemical Reactions
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• We divide the periodic table into metals, nonmetals, and metalloids.
• Nonmetals occupy the upper right portion of the periodic table.
• H is a special case.
• The first member of a group can form bonds more readily than subsequent members.
• This is due in part to the difference in atomic size.
• Compare the elemental forms of carbon and silicon
• Carbon has 5 major crystalline allotropes: diamond, graphite, buckminsterfullerene,
graphene, and carbon nanotubes).
• Diamond is a covalent-network solid with C-C sigma bonds but not pi bonds.
• Graphite, buckminsterfullerene, graphene and carbon nanotubes have pi bonds from
sideways overlap of orbitals
Chemistry of the Nonmetals
317
Chemical Reactions
• In this chapter we focus on reactions involving O2 (oxidation or combustion) and H2O (especially
proton transfer).
FORWARD REFERENCES
• Corresponding physical properties of metals will be discussed in Chapter 23 (section 23.1).
22.2 Hydrogen
Isotopes of Hydrogen
• This phenomenon is called the kinetic isotope effect.
• Tritium (T) is radioactive with a half-life of 12.3 yr.
• Deuterium and tritium are substituted for H in compounds in order to provide a molecular marker.
Such compounds are said to be "labeled" (e.g., D2O).
Properties of Hydrogen
• Hydrogen is unique.
• Hydrogen has a 1s1 electron configuration, so it is placed above Li in the periodic table.
• However, H is significantly less reactive than the alkali metals.
• The H–H bond enthalpy is high (436 kJ/mol).
• Therefore, reactions with hydrogen are slow at room temperature.
• Often the molecules must be activated with heat, irradiation, or a catalyst.
• Hydrogen forms strong covalent bonds with many elements.
• When hydrogen is ignited in air, an explosion results:
2H2(g) + O2(g) → 2H2O(l) ∆H = –571.7 kJ
Production of Hydrogen
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Chapter 22
318
• For example, Zn is added to an acidic solution and hydrogen bubbles form.
• The hydrogen bubbles out of solution and is collected in a flask.
• The collection flask is usually filled with water, so the volume of hydrogen collected is the
volume of water displaced.
2NaCl(aq) + 2H2O(l) → H2(g) + Cl2(g) + 2NaOH(aq)
Uses of Hydrogen
• About two-thirds of the 2 108 kg of hydrogen produced in the United States is used for ammonia
production via the Haber process.
• Hydrogen is used to manufacture methanol:
CO(g) + 2H2(g) → CH3OH(g)
Binary Hydrogen Compounds
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• Three types of binary hydrogen compounds are formed:
• ionic hydrides (e.g., LiH)
22.3 Group 8A: The Noble Gases
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• The noble gases are all gases at room temperature.
• He is the most important noble gas.
Noble-Gas Compounds
• The noble gases are very unreactive.
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“What’s the Use? Hydrogen” from Further Readings
Chemistry of the Nonmetals
319
• All noble gases have high ionization energies.
• The heavier noble gases react more readily than the lighter ones.
22.4 Group 7A: The Halogens
• The outer electron configurations are ns2np5.
• All halogens have large electron affinities.
• They achieve a noble-gas configuration by gaining one electron.
Properties and Preparation of the Halogens
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• The properties of the halogens vary regularly with their atomic number.
• Each halogen is the most electronegative element in its row.
• Halogens exist as diatomic molecules.
• In solids and liquids, the molecules are held together by weak London-dispersion forces.
Uses of the Halogens
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• Fluorine is an important industrial chemical.
• It is used to make fluorocarbons [used as lubricants and plastics (Teflon)].
• Chlorine is used in plastics (PVC), dichloroethane, and other organic chemicals; it is also used as a
bleaching agent in the paper and textile industries.
• NaClO is the active ingredient in bleach.
• NaBr is used in photography.
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Table 22.2 from Transparency Pack
Chapter 22
320
• Iodine is a necessary nutrient.
• It is used by the body in the synthesis of thyroid hormone.
• Lack of iodine in the diet results in a thyroid condition called goiter.
The Hydrogen Halides
• All halogens form diatomic molecules with hydrogen.
Interhalogen Compounds
• Diatomic molecules containing two different halogens are called interhalogen compounds.
• Most higher interhalogen compounds have Cl, Br, or I as the central atom surrounded by 3, 5, or 7 F
Oxyacids and Oxyanions
• Acid strength increases as the oxidation state of the halogen increases.
• All are strong oxidizing agents.
• They are generally unstable and decompose readily.
FORWARD REFERENCES
• The role of chlorides in the coordination chemistry of transition metals will be tabulated in
Chapter 23 (section 23.1).
22.5 Oxygen
Properties of Oxygen
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• Oxygen has two allotropes: O2 (dioxygen) and O3 (ozone).
• O2 is a colorless, odorless gas at room temperature.
• The electron configuration is [He]2s22p4, which means the dominant oxidation state is –2.
• It can complete an octet by gaining two e– to form an oxide anion (O2–) or by sharing 2e–.
Chemistry of the Nonmetals
321
• In covalent compounds, it forms either two single bonds or a double bond.
• The O=O bond is strong (bond enthalpy 495 kJ/mol).
Production of Oxygen
• Laboratory preparation of oxygen often involves the catalytic decomposition of KClO3 in the
presence of MnO2: 2KClO3(s) → 2KCl(s) + 3O2(g)
Uses of Oxygen
• Oxygen is one of the most widely used oxidizing agents.
Ozone
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• Ozone is a pale blue poisonous gas.
• Ozone dissociates to form oxygen:
O3(g) → O2(g) + O(g) ∆H° = 105 kJ.
• In the lower atmosphere, ozone is considered an air pollutant.
Oxides
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• Oxygen is the second most electronegative element.
• Oxides are compounds with oxygen in the –2 oxidation state.
• Nonmetal oxides are covalent.
• Metal oxides are ionic.
• Oxides that react with water to form hydroxides are called basic anhydrides, or basic oxides.
• Example: BaO in water produces Ba(OH)2.
BaO(s) + H2O(l) → Ba(OH)2(aq)
• Oxides that exhibit both acidic and basic properties are said to be amphoteric (e.g., Cr2O3).
“Water” 3-D Model from Instructor’s Resource CD/DVD
Chapter 22
322
Peroxides and Superoxides
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• Peroxides have an O–O bond and O in the –1 oxidation state.
• An example is hydrogen peroxide (H2O2).
• Superoxides have an O–O bond and O in an oxidation state of –½ .
• The superoxide ion is O2–.
• Superoxides usually form with very active metals (KO2, RbO2 and CsO2).
• Uses of superoxides:
FORWARD REFERENCES
• Several oxide minerals of transition metals will be tabulated in Chapter 23 (section 23.1).
22.6 The Other Group 6A Elements: S, Se, Te, and Po
General Characteristics of Group 6A Elements
• The outermost electron configuration is ns2np4.
Occurrences and Production of S, Se, and Te
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• Sulfur, selenium, and tellurium can all be mined from the earth.
Properties and Uses of Sulfur, Selenium, and Tellurium
• Sulfur is yellow, tasteless, and almost odorless.
• Sulfur is insoluble in water.
Chemistry of the Nonmetals
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• Its electrical conductivity is poor in the dark and increases greatly when exposed to light.
Sulfides
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• S is in the –2 oxidation state in sulfides.
• Many metals are found in the form of sulfides in ores.
Oxides, Oxyacids, and Oxyanions of Sulfur
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• SO2 is produced when sulfur is combusted in air.
• SO2 in water produces sulfurous acid, H2SO3, a weak diprotic acid.
• SO2 is toxic to fungi and is used to sterilize dried fruit and wine.
• Na2SO3 and NaHSO3 are used as preservatives.
8SO32– (aq) + S8(s) → 8S2O32– (aq)
FORWARD REFERENCES
• Several sulfide minerals of transition metals will be tabulated in Chapter 23 (section 23.1).
22.7 Nitrogen
Properties of Nitrogen
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• It is a colorless, odorless, tasteless gas composed of N2 molecules.
• It is unreactive because of the strong triple bond.
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“Herman Frasch, Sulfur King” from Further Readings
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“Sulfate Ion” 3-D Model from Instructor’s Resource CD/DVD
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Table 22.6 from Transparency Pack
Chapter 22
324
• The most common oxidation states are +5, 0, and –3 (nitrogen has an [He]2s22p3 electron
configuration).
Production and Uses of Nitrogen
• N2 is produced by fractional distillation of air.
• Nitrogen is used as an inert gas to exclude oxygen from packaged foods and in the manufacture of
Hydrogen Compounds of Nitrogen
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• Ammonia is one of the most important compounds of nitrogen.
• Ammonia is a colorless toxic gas with an irritating aroma.
• In the laboratory, ammonia is produced by the reaction between NaOH and an ammonium salt:
NH4Cl(aq) + NaOH(aq) → NH3(g) + H2O(l) + NaCl(aq)
• Commercially, ammonia is prepared by the Haber process.
N2(g) + 3H2(g) → 2NH3(g)
Oxides and Oxyacids of Nitrogen
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• There are three common oxides of nitrogen:
• N2O (nitrous oxide).
• This is also known as laughing gas and is used as an anesthetic.
• NO (nitric oxide).
• This is a toxic, colorless gas; it is an important neurotransmitter.
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“Hydrazine” 3-D Model from Instructor’s Resource CD/DVD
32
“Some History of Nitrates” from Further Readings
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“Nitrogen Dioxide and Dinitrogen Tetroxide” Movie from Instructor’s Resource CD/DVD
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“Dinitrogen Trioxide” 3-D Model from Instructor’s Resource CD/DVD
Chemistry of the Nonmetals
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• oxidation of NH3 by oxygen to form NO (usually using a Pt catalyst).
4NH3(g) +5O2(g) → 4NO(g) + 6H2O(g)
• oxidation of NO by oxygen to form NO2 (unreacted NO is recycled).
• NO2 dissolution in water to form nitric acid.
FORWARD REFERENCES
• Ammonia complexes of cobalt will be discussed in Chapter 23 (section 23.2).
22.8 The Other Group 5A Elements: P, As, Sb, and Bi
General Characteristics of the Group 5A Elements
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• The outermost shell electron configuration is ns2np3.
Occurrence, Isolation, and Properties of Phosphorus
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• Occurs mainly in phosphorus minerals [e.g., phosphate rock, Ca3(PO4)2].
• Elemental P4 is produced by reduction:
2 Ca3(PO4)2(s) + 6SiO2(s) + 10C(s) → P4(g) + 6CaSiO3(l) + 10CO(g).
• There are two allotropes of phosphorus: red and white.
Phosphorus Halides
• Phosphorus forms a variety of compounds with halogens with the tri- and pentahalides being the most
important.
• The most important is PCl3 which is used in soap, detergent, plastic, and insecticide production.
• Preparation of phosphorus halides involves direct oxidation of elemental phosphorous with elemental
halogen.
• For example: 2P(s) + 3Cl2(g) → 2PCl3(l)
• In the presence of excess chlorine:
Chapter 22
326
PCl3(l) + Cl2(g) PCl5(s)
• In the presence of water hydrolysis occurs readily:
PBr3(g) + 3H2O(l) → H3PO3(aq) + 3HBr(aq)
PCl5(l) + 4H2O(l) → H3PO4(aq) + 5HCl(aq)
Oxy Compounds of Phosphorus
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• Oxygen-containing phosphorus compounds are extremely important.
• Phosphorus(III) oxide, P4O6 is made by reacting white phosphorus with a limited supply of
oxygen.
• Phosphorus(V) oxide, P4O10 is made by reacting phosphorus with excess oxygen.
• Phosphorus(III) oxide, P4O6 produces phosphorous acid, H3PO3 in water.
• The oxides of phosphorus are acidic.
• H3PO3 is a weak diprotic acid (the H attached to P is not acidic).
FORWARD REFERENCES
• Phosphate groups in RNA and DNA will be discussed in Chapter 24 (section 24.10).
22.9 Carbon
Elemental Forms of Carbon
• Carbon constitutes about 0.027% of the Earth’s crust.
• Carbon is the main constituent of living matter.
• The study of carbon compounds (organic compounds) is called organic chemistry.
• There are five allotropic forms of carbon: two of these are graphite and diamond:
Oxides of Carbon
• Carbon forms two principal oxides: CO and CO2.
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“Tetraphosphorus Hexoxide” 3-D Model from Instructor’s Resource CD/DVD
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“Tetraphosphorus Decoxide” 3-D Model from Instructor’s Resource CD/DVD
Chemistry of the Nonmetals
327
• CO (carbon monoxide) is formed when carbon or hydrocarbons are burned in a limited supply of
oxygen. 2C(s) + O2(g) → CO(g)
• CO is very toxic (binds irreversibly to hemoglobin, interfering with oxygen transport).
• It is odorless, colorless, and tasteless.
Carbonic Acid and Carbonates
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• When CO2 dissolves in water (moderately soluble) a diprotic acid, carbonic acid, forms:
CO2(aq) + H2O(l) H2CO3(aq)
• Carbonic acid is responsible for giving carbonated beverages a sharp acidic taste.
• Two salts of carbonic acid may be obtained by neutralization.
• Partial neutralization of H2CO3 gives hydrogen carbonates (bicarbonates): HCO3–.
Carbides
• Carbides are binary compounds of C and metals, metalloids, and certain nonmetals.
• There are three types of carbides:
• Ionic carbides (formed by active metals)
• Most contain the acetylide ion, C22–.
• An example is CaC2.
Chapter 22
328
• CaC2 is used in the formation of acetylene:
CaC2(s) + 2H2O(l) → Ca(OH)2(aq) + C2H2(g)
Other Inorganic Compounds of Carbon
• Two interesting inorganic compounds of carbon are HCN and CS2.
• HCN (hydrogen cyanide) is an extremely toxic gas.
• HCN is produced by reacting a salt, (e.g., NaCN) with acid.
FORWARD REFERENCES
• Chemistry of carbon will be discussed in detail throughout Chapter 24.
22.10 The Other Group 4A Elements: Si, Ge, Sn, and Pb
General Characteristics of Group 4A Elements
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• The outermost electron configuration is ns2np2.
• The electronegativities are low.
Occurrence and Preparation of Silicon
• Si is the second most abundant element in the Earth’s crust.
• Elemental Si is prepared by reducing SiO2:
SiO2(l) + 2C(s) → Si(l) + 2CO(g)
• Silicon has many important uses in the electronics industry.
• Wafers of Si are cut from cylindrical Si crystals.
Table 22.8 from Transparency Pack
Chemistry of the Nonmetals
329
Silicates
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• More than 90% of the Earth’s crust is composed of compounds of Si and O.
• The most common oxidation state of Si is +4.
• Silicates are compounds in which Si has four O atoms surrounding it in a tetrahedral arrangement.
• A single-strand silicate chain can form with a Si2O64– repeating unit.
• An example is enstatite (MgSiO3).
• Consider a structure with two vertices linked to three other tetrahedra.
• A two-dimensional sheet results.
• The mineral talc [talcum powder, Mg3(Si2O5)2(OH)2] results.
• Many minerals are based on silicates.
Glass
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• Glasses result when silicates are heated (Si–O bonds are broken) and then rapidly cooled.
• The Si–O bonds are re-formed before the atoms are able to organize into an ordered arrangement.
• The amorphous solid is called quartz glass or silica glass.
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Figure 22.34 from Transparency Pack
52
“Silicate Tetrahedron” 3-D Model from Instructor’s Resource CD/DVD
Chapter 22
330
Silicones
• Silicones consist of O–Si–O chains with Si–R (R is an organic group such as CH3) bonds filling the
Si valency.
22.11 Boron
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• Boranes are compounds of boron and hydrogen.
• BH3 is the simplest borane.
• It reacts with itself to form diborane, B2H6.
• Hydrogen appears to form two bonds.
• These are called bridging hydrogens.
Chemistry of the Nonmetals
331
Further Readings:
1. Alton Banks, “What’s the Use? Hydrogen,” J. Chem. Educ., Vol. 66, 1989, 801. Each of the “What’s
the Use?” articles, written by Alton Banks, focuses on the uses of a specific element. See volumes 66
(1989), 67 (1990), 68 (1991) and 69 (1992) for other elements!
6. Kathryn R. Williams, “The Discovery of Oxygen and other Priestley Matters,” J. Chem. Educ., Vol. 80,
2003, 1129–1131.
7. Michael Laing, “The Three Forms of Molecular Oxygen,” J. Chem. Educ., Vol. 66, 1989, 453–454.
8. Derek W. Smith, “An Acidity Scale for Binary Oxides,” J. Chem. Educ., Vol. 64, 1987, 480–481.
13. Natalie I. Foster and Ned D. Heindel, “The Discovery of Nitroglycerine: Its Preparation and
Therapeutic Utility,” J. Chem. Educ., Vol 58, 1981, 364–365.
14. Annette Lykknes and Lise Kvittinger, “Arsenic: Not So Evil After All?” J. Chem. Educ., Vol. 80,
2003, 497–500.
15. Terence P. Lee, “Keeping the Fire Cold,” Chemistry in Britain, January 1996, 41–45. An article on
the importance of phosphorous.
Chapter 22
332
20. Russell N. Grimes, “Boron Clusters Come of Age,” J. Chem. Educ., Vol. 81, 2004, 658–672.
Live Demonstrations:
1. Lee R. Summerlin, Christie L. Borgford, and Julie B. Ealy, “Making Hydrogen Gas from an Acid and a
Base,” Chemical Demonstrations, A Sourcebook for Teachers, Volume 2 (Washington: American
Chemical Society, 1987), pp. 33–34. Gas-generating reactions are carried out in two flasks fitted with
balloons; the balloons inflate as they collect the hydrogen gas generated.
4. Thomas G. Richmond and Paul F. Kraus, “Demonstrating a Lack of Reactivity Using a Teflon-Coated
Pan,” J. Chem. Educ., Vol. 72, 1995, 731. A short demonstration of the wonders of Teflon.
5. Bassam Z. Shakhashiri, “Combining Volume of Oxygen with Sulfur,” Chemical Demonstrations: A
Handbook for Teachers of Chemistry, Volume 2 (Madison: The University of Wisconsin Press, 1985), pp.
190–192. A demonstration of the production of SO2, rather than SO3, upon combustion of sulfur in O2.
