© 2015 pearson education, inc. chapter 22 chemistry of the nonmetals lecture presentation

© 2015 Pearson Education, Inc.

Chapter 22

Chemistry of the Nonmetals

Lecture Presentation



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)



Trends



First Group Members Often Differ

For NONMETALS, FIRST group memberso are able to

accommodate fewer bonded neighbors.

o are more likely to form π bonds because they can get closer to other atoms.



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.



Importance of Oxygen and WaterO2 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.



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.



Isotopes of Hydrogen• Protium (1H): 99.9844% of all hydrogen• Deuterium (2H): 0.0156% 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.



Physical Properties of Hydrogen• Hydrogen is unique.

– It does not belong to any group.– It is not a member of Group IA. (IE =

1312 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.



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.



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.



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).



The Hydrogen Economy?

• 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)

2 H2(g) + O2(g) 2 H2O(g) ΔH = –483.6 kJ



The Hydrogen Economy



Binary Hydrides

There are three types of hydrides:– Ionic– Metallic– Molecular



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.



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.



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.



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.



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.



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.



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)



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.



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.



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.



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.



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.



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



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–).



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.



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)



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)



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).



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.



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.



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.



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.



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



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, S2

2–, found in minerals like iron pyrite (fool’s gold).



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.



Thio

• In naming, thio means sulfur in place of oxygen.• Thiosulfate ion is S2O3

2– (one S atom replaces one of the O atoms in sulfate).



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



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.



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.



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



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



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.



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).



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)



Graphite vs. Diamond

Graphite• Metallic luster• Conducts electricity• Sheets held together

by dispersion forces• D = 2.25 g/cm3

• Soft, black, slippery solid

Diamond• Crystalline• Electrical insulator• sp3-hybridized three-

dimensional network of C atoms

• D = 3.51 g/cm3

• Clear, hard solid



Oxides of Carbon• Carbon monoxide (CO)

– 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).



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.



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.



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



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



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.



Glass• 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



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).



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.

© 2015 pearson education, inc. chapter 22 chemistry of the nonmetals lecture presentation

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