1. Chapter 22 Chemistry of the Nonmetals
Lecture Presentation
Chapter 22 Chemistry of the Nonmetals

2. Descriptive Chemistry
How elements occur in nature
How elements are isolated from their sources
How elements are used
Emphasis on H, O, N, C
Look for TRENDS, rather than memorize everything (CLASSIFY)

3. Trends

4. First Group Members Often Differ
For NONMETALS, FIRST group members
are able to accommodate fewer bonded neighbors.
are more likely to form π bonds because they can get closer to other atoms.

5. Comparison of π Bonding
CO2 molecules are small molecules with π bonds.
SiO2 is an extended lattice structure with tetrahedral Si connected to four O atoms and each O atom connected to two Si atoms.

6. Importance of Oxygen and Water
O2 and H2O are abundant in our environment.
About one-third of the chemical reactions in this chapter involve either O2 or H2O.
Proton transfer: the weaker a Brønsted–Lowry acid, the stronger its conjugate base.
Combustion: burning in the presence of O2, H becomes H2O; C becomes CO2; N tends to become N2 but can form NO in special cases.

7. Hydrogen Hydrogen was discovered by Henry Cavendish (1731–1810).
It is the most abundant element in the universe.
Although 75% of the mass of the universe, it is only 0.87% of the Earth’s mass, mostly as water.

8. Isotopes of Hydrogen Protium (1H): 99.9844% of all hydrogen
Deuterium (2H): % of all hydrogen; it is NOT radioactive; it is often represented as D (e.g., D2O, deuterium oxide, is often called “heavy water”).
Tritium (3H): radioactive (half-life of 12.3 years by beta emission)
The difference in masses between 1H and 2H results in different physical properties for deuterated compounds.
2H and 3H are often used to “label” compounds to follow chemical reactions.

9. Physical Properties of Hydrogen
Hydrogen is unique.
It does not belong to any group.
It is not a member of Group IA. (IE = kJ/mol; for Li, IE = 520 kJ/mol!)
It is not a member of Group VIIA. (Although H– forms, EA = –73 kJ/mol, not nearly that of the halogens.)
It has very low melting (–259 °C) and boiling (–253 ° C) points.

10. Reactivity of Hydrogen
Hydrogen has very large bond enthalpies.
Reacts slowly, BUT when activated, H atoms react quickly and the reactions are very exothermic.
Hydrogen forms strong covalent bonds with many other elements.
It is VERY explosive in oxygen—this property led to its being used in liquid-fuel rocket engines.

11. Hydrogen Production
In the lab (small quantities), hydrogen is produced by using a more active metal and an acid.
Hydrogen is commercially produced from reaction of methane (CH4) with steam at 1100 °C or carbon and steam above 1000 °C:
CH4(g) + H2O(g)  CO(g) + 3 H2(g)
CO(g) + H2O(g)  CO2(g) + H2(g)
C(s) + H2O(g)  H2(g) + CO(g)
Its production from the electrolysis of water is not energy efficient.

12. Uses of Hydrogen
About half of the hydrogen produced is used to synthesize ammonia (NH3) in the Haber process.
Most of the remaining hydrogen is used in “cracking”—producing lower molecular weight hydrocarbons for fuel (more gasoline, diesel, etc.).
It is also used to produce methanol (CH3OH).

13. The Hydrogen Economy? 2 H2(g) + O2(g)  2 H2O(g) ΔH = –483.6 kJ
Using hydrogen as a fuel has advantages:
Its reaction with oxygen is highly exothermic.
Water is the only product.
Its low mass gives a high energy density.
Difficulties:
Producing hydrogen takes energy
Storage of hydrogen (need large volumes)
Safety (explosive with oxygen in any heat)

14. The Hydrogen Economy

15. Binary Hydrides There are three types of hydrides: Ionic Metallic
Molecular

16. Ionic Hydrides
Ionic hydrides (with H–) are formed between hydrogen and alkali metals or heavy alkaline earth metals (Ca, Sr, Ba).
They are very strong bases and reducing agents.
They react readily with water and oxygen, so they must be stored free from moisture and air.

17. Metallic Hydrides
These are formed between hydrogen and transition metals.
They often form nonstoichiometric ratios of metal to hydrogen.
TiH1.8, for example
They retain electrical conductivity and other metallic properties.

18. Molecular Hydrides
These are formed between hydrogen and nonmetals or metalloids.
They are usually gases or liquids at room temperature and normal atmospheric pressure.
Stability of the nonmetal hydrides decreases down the group.

19. Noble Gases
Noble gases are extremely stable and unreactive, as seen in the very high ionization energies.
Liquid He (boiling point 4.2 K) is used as a coolant. It is found in many natural gas wells.
Ne, Ar, and Kr are used in lighting, displays, and laser applications.
Ar is used in light bulbs and as a protective atmosphere in welding.

20. Noble Gas Compounds Discovered in 1962
Xe forms most of these compounds
Lower ionization energy and expanded valence shell lead to reactions.
KrF2 (decomposes at –10 °C) and a few other Kr compounds (unstable above –40 °C) also can be made.

21. Halogens The halogens have outer electron configurations of ns2np5.
They have large negative electron affinities and large ionization energies.
They tend to accept one electron to form anions.
Other than F, higher oxidation states (up to +7) are possible.
At is a very unstable radioactive element; little is known about its chemistry.

22. Halogens—the Elements
Most properties vary in a regular fashion from F to I.
Diatomic molecules; dispersion forces only (I2 is a solid; Br2 is a liquid; Cl2 and F2 are gases)
F2 is very reactive, difficult to work with.
Halogens are good oxidizing agents. This is a periodic property. (F2 most reactive; I2 least reactive)

23. Preparation of the Halogens
Oxidation of a halide anion occurs for any elemental halogen higher in the group.
Halides can be oxidized by electrolysis of salt solutions; bromine and iodine are usually prepared by reaction of salt solutions with chlorine.
Fluorine has an unusually high reduction potential. It can easily oxidize water. It is prepared by electrolysis of KF in anhydrous HF.

24. Uses of Halogens
Fluorine reacts to form fluorocarbon compounds used as lubricants, refrigerants, and plastics.
Teflon is a polymer of fluorocarbons.
Chlorine is the most used halogen, used for HCl, plastics, bleaches, and water purification.
Iodized salt contains KI in it, to prevent goiter, a thyroid disease.

25. Hydrogen Halides
Aqueous solutions of HCl, HBr, and HI are strong acids.
HF and HCl can be produced by reacting salts with H2SO4.
Br– and I– oxidize too easily, so one must use a weaker oxidizing acid, like H3PO4.

26. Interhalogen Compounds
Binary compounds made up of two different halogens
Most common: central Cl, Br, or I atom surrounded by many F atoms (e.g., IF3, IF5, or IF7)
Also iodine/chlorine (three or five Cl atoms) interhalogen compounds form.
These are powerful oxidizing agents.

27. Oxyacids and Oxyanions
Oxyacid strength increases with the increasing oxidation number of the central halogen.
Oxyacids are strong oxidizers.
Oxyanions are generally more stable than the corresponding acids.
Perchlorates are stable unless heated, when they are for rocket fuel.

28. Oxygen Discovered in 1774 by Joseph Priestley
Named by Lavoisier (“acid former”)
Most abundant element by mass in Earth’s crust and human body
Two allotropes: O2 (“oxygen gas”) and O3 (ozone)
Oxygen gas properties:
Colorless, odorless
Condenses at –183 °C, freezes at –218 °C

29. Bonding of Oxygen
Elemental oxygen: diatomic molecule, strong double bonds (bond enthalpy 495 kJ/mol)
Covalent molecules: single or double bonds; very strong bonds (compounds are often thermodynamically more stable than the element)
Ionic compounds: oxygen gains two electrons to fill its valence shell, producing the oxide anion (O2–).

30. Producing Oxygen
Nearly all commercial oxygen is obtained from the air.
Since oxygen gas boils at a higher temperature than liquid nitrogen, liquefied air is warmed to cause the nitrogen to boil (impurities: nitrogen and argon).
Laboratory production: heating metal chlorates or hydrogen peroxide (MnO2 catalyst)
In the atmosphere, oxygen is replenished by photosynthesis, where green plants produce sugars.

31. Uses of Oxygen Oxygen is most widely used as an oxidizing agent:
Steel industry, to remove impurities from steel
Bleach pulp and paper
Welding (with acetylene)

32. Ozone (O3) Pale blue, poisonous gas with a sharp odor
Extremely irritating to the respiratory system (ground-level pollutant)
Stronger oxidizer than O2, used to purify water (kills bacteria)
Used in organic syntheses (severs C C bonds)
Absorbs UV light in the upper atmosphere, creating reactive O atoms
Prepared by passing electricity through dry O2 (sometimes by lightning strikes)

33. Oxides –2 oxidation state most common for oxygen (by far)
Nonmetals form covalent oxides; most give oxyacids when dissolved in water (like SO2 → H2SO3).
These oxides that form acids in water are called acidic anhydrides or acidic oxides.
Ionic oxides that dissolve in water form the hydroxide ion; these are basic anhydrides or basic oxides.
Some oxides are amphoteric (act as acid or base).

34. Peroxides and Superoxides
Peroxides have O—O bonds, giving an oxidation state of –1 for O22–. They are formed with the cations Na, Ca, Sr, and Ba.
Hydrogen peroxide (H2O2) is very reactive; it disproportionates—reacts with itself as oxidizing and reducing agent. Uses: antiseptic, bleach fabric.
The superoxide ion is O2–. It is formed with the most active metals as cations (K, Rb, and Cs).
Superoxides react with water to produce O2. Use: in a breathing apparatus.

35. Other Group 6A Elements
S, Se, and Te have oxidation states of –2 and positive oxidation states up to +6.
They can use d-orbitals to expand beyond the octet.
Po has no stable isotopes.

36. Selenium and Tellurium
These elements are anions in minerals with Cu and Pb.
They are naturally found as helical chains of atoms.
Selenium is not electrically conductive in the dark, but it is quite so in light. For that reason, it is used in light meters, photoelectric cells, and photocopiers.

37. Sulfur in Nature
Elemental sulfur is a yellow solid; it consists of rings of eight S atoms.
Sulfur can also be found as sulfide and sulfate minerals.
One problem with the appearance of sulfur in nature: it is an impurity in coal and petroleum, creating sulfur oxides (acid rain) when it is burned.

38. Uses of Sulfur Most common uses:
Production of sulfuric acid (an extremely important industrial compound because it is a strong acid, an oxidizing agent, and a dehydrating agent)
Vulcanization of rubber

39. Sulfides and Disulfide
Sulfide is sulfur in the –2 oxidation state.
Hydrogen sulfide and organic sulfides have an unpleasant odor (rotten egg smell); dimethyl sulfide is added to natural gas as a safety agent—you smell it if the gas isn’t burning.
Many minerals, such as galena (PbS) and cinnabar (HgS), are ionic sulfides.
Pyrites contain the disulfide ion, S22–, found in minerals like iron pyrite (fool’s gold).

40. Sulfur Oxides, Oxyacids, and Oxyanions
Sulfuric acid (most important!) has already been mentioned.
Bisulfates are used in toilet bowl cleaners and to adjust pH in swimming pools and hot tubs.
SO2 is a poison, particularly to lower organisms.
Sulfites and bisulfites are added to foods and wines to kill bacteria but can cause severe asthmatic reactions.

41. Thio In naming, thio means sulfur in place of oxygen.
Thiosulfate ion is S2O32– (one S atom replaces one of the O atoms in sulfate).

42. Nitrogen 78% of Earth’s atmosphere
Sources: nitrate salts (saltpeter in India; Chile saltpeter in Chile); mostly by distilling liquid air
Colorless, odorless, tasteless
m.p. = –210 °C; b.p. = –196 °C
Not very reactive as element due to N N; inert gas in industrial usage
Most important use: fertilizers
Oxidation states from –3 to +5

43. Nitrogen is Converted to Ammonia
Nitrogen is converted to ammonia using the Haber process. The primary use of ammonia is to produce fertilizers.
Ammonia reacts with hypochlorite to produce hydrazine. It is a strong reducing agent. Methylhydrazine is a rocket fuel.
Ammonia is converted to NO by the Ostwald process: its reaction with oxygen at high temperature on a Pt catalyst.

44. Nitrogen Oxides
Nitrous oxide (N2O, laughing gas) was the first general anesthetic.
It is also used in aerosol products like whipped cream.
Nitric oxide (NO) is a slightly toxic, colorless gas. It is a neurotransmitter involved in vasodilation in humans.
NO reacts with O2 in the air to produce nitrogen dioxide (NO2), which is a pollutant found in smog.

45. Oxyacids of Nitrogen Nitric acid (HNO3) A strong oxidizing acid
Used in fertilizer and explosive production (TNT, nitroglycerine, nitrocellulose)
Very important industrial compound
Nitrous acid (HNO2)
Less stable, weak acid
Tends to disproportionate to NO and HNO3

46. Other Group 5A Elements
Nonmetals (N, P); metalloids (As, Sb); metal (Bi)
Size and metallic character increase down Group 5A.
N2: diatomic molecules; all others: only single bonds between atoms

47. Phosphorus Main source: phosphate minerals
Two allotropes of phosphorus:
White phosphorus (P4), which is formed by reduction of Ca3(PO4)2 with C in the presence of SiO2 in large furnaces; it bursts into flames if exposed to O2 in the air.
Red phosphorus, which is produced by heating white phosphorus to 400 °C in the absence of air, is more stable in air.

48. Oxy Compounds of Phosphorus
Phosphorus(III) oxide (P4O6) and phosphorus(V) oxide (P4O10) are anhydride forms of phosphorous (H3PO3) and phosphoric (H3PO4) acids.
Phosphoric acid and its salts are used in fertilizers and in detergents as water softeners (binding metal cations).
Polyphosphates form by dehydration between phosphate groups.
Phosphates are important in biological systems (ATP/ADP conversion for energy).

49. Carbon
Several allotropes: graphite, diamond, fullerenes, carbon nanotubes, graphene
Occurs in nature as carbonates, and in coal, petroleum, and natural gas
Some uses: pigment in black ink (carbon black); remove odors from air and drinking water (activated charcoal); strong composite materials (carbon fibers); cutting tools (diamond)

50. Graphite vs. Diamond Graphite Diamond Metallic luster
Conducts electricity
Sheets held together by dispersion forces
D = 2.25 g/cm3
Soft, black, slippery solid
Crystalline
Electrical insulator
sp3-hybridized three-dimensional network of C atoms
D = 3.51 g/cm3
Clear, hard solid

51. Oxides of Carbon Carbon monoxide (CO) Carbon dioxide (CO2)
This is an odorless, colorless, tasteless gas.
It acts as a Lewis base with transition metals.
CO binds preferentially to iron in hemoglobin, inhibiting O2 transport.
It is used as a fuel and as a reducing agent in metallurgy.
Carbon dioxide (CO2)
It is a colorless, odorless gas.
Although a minor component of the atmosphere, it has a major role in the greenhouse effect.
Uses include production of carbonated beverages; it is also used as a refrigerant (as dry ice).

52. Carbonic Acid and Carbonates
Dissolved CO2 in water is in equilibrium with carbonic acid, H2CO3.
Carbonates are found as minerals such as calcite, CaCO3, the primary constituent of limestone, marble, and shells of marine animals.
Washing soda and baking soda are forms of sodium carbonate and sodium bicarbonate.
Lime, CaO, is formed from CaCO3. It is used to make mortar.

53. Carbides
Binary compounds of carbon and a metal, metalloid, or certain nonmetals
Calcium carbide is a source for acetylene (used in welding).
Many transition metals form interstitial carbides, making very hard and heat-resistant materials, such as tungsten carbide.
Covalent carbides are made with Si and B; SiC is almost as hard as diamond and used in cutting tools.

54. Group 4A
Clear trend of increasing metallic character down group (C is nonmetal; Si, Ge are metalloids; Sn, Pb are metals)
+2 oxidation state common for Ge, Sn, Pb
Other than C, able to form more than four bonds
Carbon only group member that commonly forms bonds to itself

55. Silicon The second most abundant element in Earth’s crust
Occurs as SiO2 and silicate minerals
Obtained by heating SiO2 and C
A semiconductor used in making transistors and solar cells
Purified by a process known as zone refining

56. Silicates
Silicates have a central silicon atom that is surrounded by four oxygens.
These units can share oxygen atoms to connect into sheets or strands.
Talc and asbestos are two examples of molecules containing these structures.

57. Glass Silicones Formed by rapidly cooling molten quartz
Soda-lime glass (in windows and bottles): made from sand, lime (CaO), and soda ash (Na2O)
Cobalt glass: add CoO to soda-lime glass
Harder glass if K2O replaces Na2O
Lead crystal: replace CaO with PbO
Borosilicates (Pyrex, Kimax) withstand higher temperatures.
Silicones
Contain O—Si—O chains
Chain length and cross-linkages applied to make materials for lubricants and sealants

58. Boron Boron is the only nonmetal in group 3A.
Compounds of boron and hydrogen are called boranes, the simplest one being diborane, B2H6.
Diborane demonstrates an unusual type of bonding in which two boron atoms share one hydrogen atom (a bridging hydrogen).

59. Boron with H and O
Borane anions, such as borohydride, BH4–, are good reducing agents and sources of hydride ion.
Diborane is extremely reactive, spontaneously flammable in air (exothermic!), creating B2O3, boric oxide.
Boric oxide is the anhydride of boric acid, which is actually amphoteric: H3BO3 or B(OH)3.
Borax, the salt of tetraboric acid, H2B4O7, is used in cleaning products.