Chapter 22 Chemistry of the Nonmetals

Chapter 22Chemistry

of the Nonmetals

2008, Prentice Hall

Chemistry: A Molecular Approach, 1st Ed.Nivaldo Tro

Roy KennedyMassachusetts Bay Community College

Wellesley Hills, MA

Tro, Chemistry: A Molecular Approach 2

Nanotubes• nanotubes – long, thin, hollow cylinders of atoms• carbon nanotube = sp2 C in fused hexagonal rings

electrical conductors

• boron-nitride nanotubes = rings of alternating B and N atoms isoelectronic with Csimilar size to Caverage electronegativity of B & N about the same as Celectrical insulators


Properties of BN and C


Main Group Nonmetals


Atomic Radius and Bonding• atomic radius decreases across the period• electronegativity, ionization energy increase across the

period• nonmetals on right of p block form anions in ionic

compoundsoften reduced in chemical reactions

making them oxidizing agents

• nonmetals on left of p block can form cations and electron-deficient species in covalent bonding

• nonmetals near the center of the p block tend to use covalent bonding to complete their octets

• bonding tendency changes across the period for nonmetals from cation and covalent; to just covalent; to anion and covalent


Insulated Nanowire


Silicates

• the most abundant elements of the Earth’s crust are O and Si

• silicates are covalent atomic solids of Si and Oand minor amounts of other elementsfound in rocks, soils, and clayssilicates have variable structures – leading to the

variety of properties found in rocks, clays, and soils


Bonding in Silicates• each Si forms a single covalent bond to 4 O

sp3 hybridizationtetrahedral shapeSi-O bond length is too long to form Si=O

• to complete its octet, each O forms a single covalent bond to another Si

• the result is a covalent network solid


Quartz

• a 3-dimensional covalent network of SiO4 tetrahedrons

• generally called silica

• formula unit is SiO2

• when heated above 1500C and cooled quickly, get amorphous silica which we call glass


Aluminosilicates

• Al substitutes for Si in some of the lattice sites

• SiO2 becomes AlO2−

• the negative charge is countered by the inclusion of a cationAlbite = ¼ of Si replaced by Al; Na(AlO2)(SiO2)3

Anorthite = ½ of Si replaced by Al; Ca(AlO2)2(SiO2)2


Silicates Made of Individual Units• O of SiO4 picks up electrons from metal to form SiO4

4−

• if the SiO44− are individual units neutralized by cations,

it forms an orthosilicatewillemite = Zn2SiO4

• when two SiO4 units share an O, they form structures called pyrosilicates with the anion formula Si2O7

6−

hardystonite =Ca2ZnSi2O7


Single Chain Silicates

• if the SiO44− units link as long

chains with shared O, the structure is called a pyroxene

• formula unit SiO32-

• chains held together by ionic bonding to metal cations between the chainsdiopside = CaMg(SiO3)2 where

Ca and Mg occupy lattice points between the chains


Double Chain Silicates• some silicates have 2

chains bonded together at ½ the tetrahedra – these are called amphiboles

• often results in fibrous mineralsasbestostremolite asbestos =

Ca2(OH)2Mg5(Si4O11)2


Sheet Silicates• when 3 O of each

tetrahedron are shared, the result is a sheet structure called a phyllosilicate

• formula unit = Si2O52−

• sheets are ionically bonded to metal cations that lie between the sheets

• talc and mica


Mica: a Phyllosilicate


Silicate Structures


Boron• metalloid• at least 5 allotropes, whose structures are

icosahedrons each allotrope connects the icosahedra in

different ways

• less than 0.001% in Earth’s crust, but found concentrated in certain areas almost always found in compounds with O

borax = Na2[B4O5(OH)4]8H2O kernite = Na2[B4O5(OH)4]3H2O colemanite = Ca2B6O115H2O

• used in glass manufacturing – borosilicate glass = Pyrex

• used in control rods of nuclear reactors


Boron Trihalides

• BX3

• sp2 Btrigonal planar, 120 bond anglesforms single bonds that are shorter and stronger than

sp3 Csome overlap of empty p on B with full p on

halogen

• strong Lewis Acids


Boron-Oxygen Compounds

• form structures with trigonal BO3 units

• in B2O3, six units are linked in a flat hexagonal B6O6 ringmelts at 450C

melt dissolves many metal oxides and silicon oxides to form glasses of different compositions


Boranes closo-Boranes

• compounds of B and H

• used as reagent in hydrogenation of C=C

• closo-Boranes have formula BnHn2− and form

closed polyhedra with a BH unit at each vertex


Boranes nido-Boranes and arachno-Boranes

• nido-Boranes have formula BnHn+4 consisting of cage B missing one corner

• arachno-Boranes have formula BnHn+6 consisting of cage B missing two or three corners


Carbon• exhibits the most versatile bonding of all the

elements• diamond structure consists of tetrahedral sp3

carbons in a 3-dimensional array• graphite structures consist of trigonal planar sp2

carbons in a 2-dimensional arraysheets attracted by weak dispersion forces

• fullerenes consist of 5 and 6 member carbon rings fused into icosahedral spheres of at least 60 C

23

Crystalline Allotropes of CarbonDiamond Graphite Buckminster-

fullerene, C60

Color clear-blue black black

Density, g/cm3 3.53 2.25 1.65

Hardness, Mohs Scale 10 0.5

Electrical Conductivity, (•cm)-1 ~10-11 7.3 x 10-4 ~10-14

Thermal Conductivity, W/cm•K 23 20 ()

Melting Point, C ~3700 ~3800 800 sublimes

Heat of Formation (kcal/mol) 0.4 0.0 9.08

Refractive Index 2.42 ─ 2.2 (600 nm)

Source Kimberlite

(S. Africa)

Pegmatite

(Sri Lanka)

Shungite

(Russia)


Allotropes of Carbon - Diamond

Inert to Common AcidsInert to Common BasesNegative Electron Affinity TransparentHardestBest Thermal ConductorLeast CompressibleStiffest


Allotropes of Carbon - Graphite

Soft and Greasy FeelingSolid Lubricant Pencil “Lead”Conducts ElectricityReacts with Acids and Oxidizing Agents


Noncrystalline Forms of Carbon• coal is a mixture of hydrocarbons and carbon-rich particles

the product of carbonation of ancient plant material carbonation removes H and O from organic compounds in the form of

volatile hydrocarbons and water

• anthracite coal has highest C content• bituminous coal has high C, but high S• heating coal in the absence of air forms coke

carbon and ash

• heating wood in the absence of air forms charcoal activated carbon is charcoal used to adsorb other molecules

• soot is composed of hydrocarbons from incomplete combustion carbon black is finely divided form of carbon that is a component of

soot used as rubber strengthener


Allotropes of Carbon - Buckminsterfullerene

Sublimes between 800°CInsoluble in waterSoluble in tolueneStable in air Requires temps > 1000°C to decomposeHigh electronegativity Reacts with alkali metalsBehavior more aliphatic than aromatic


Nanotubes• long hollow tubes constructed of fused C6 rings• electrical conductors• can incorporate metals and other small

molecules and elementsused to stabilize unstable molecules

• single-walled nanotubes (SWNT) have one layer of fused rings

• multi-walled nanotubes (MWNT) have concentric layers of fused rings


Nanotubes


Nanocars


Carbides• carbides are binary compounds of C with a less electronegative

element• ionic carbides are compounds of metals with C

generally alkali or alkali earth metals often dicarbide ion, C2

2− (aka acetylide ion) react with water to form acetylene, C2H2

• covalent carbides are compounds of C with a low-electronegativity nonmetal or metalloid silicon carbide, SiC (aka carborundum)

very hard

• metallic carbides are metals in which C sits in holes in the metal lattice hardens and strengthens the metal without affecting electrical conductivity steel and tungsten carbide


Calcium Carbide


CementiteFe3C regions found in steel


Carbon Oxides• CO2

0.04% in atmosphere increased by 25% over the past century

high solubility in water due to reaction with water to form HCO3

− ions triple point −57C and 5.1 atm

liquid CO2 doesn’t exist at atmospheric pressure solid CO2 = dry ice

• CO colorless, odorless, tasteless gas relatively reactive

2 CO + O2 2 CO2

– burns with a blue flame reduces many nonmetals

– CO + Cl2 COCl2 (phosgene)– CO + S COS (fungicide)


Carbonates• solubility of CO2 in H2O due to carbonate formation

CO2 + H2O H2CO3

H2CO3 + H2O H3O+ + HCO3−

HCO3−

+ H2O H3O+ + CO32−

• washing soda = Na2CO310H2Odoesn’t decompose on heating

• all carbonate solutions are basic in waterdue to CO3

2− + H2O OH− + HCO3

2−

• baking soda = NaHCO3

decomposes on heating to Na2CO3, H2O and CO2


Elemental Nitrogen

• N2

78% of atmospherepurified by distillation of liquid air, or

filtering air through zeolitesvery stable, very unreactive

NN


Elemental Phosphorus• P

white phosphorus white, soft, waxy solid that is flammable and toxic stored under water to prevent spontaneous combustion 2 Ca3(PO4)2 (apatite) + 6 SiO2 + 10 C P4(g, wh) + 6 CaSiO3 + 10 CO tetrahedron with small angles 60

red phosphorus formed by heating white P to about 300C in absence of air amorphous mostly linked tetrahedra not as reactive or toxic as white P used in match heads

black phosphorus formed by heating white P under pressure most thermodynamically stable form, therefore least reactive layered structure similar to graphite


Phosphorus

White PhosphorusRed Phosphorus


Hydrides of Nitrogen• ammonia, NH3

pungent gas basic NH3 + H2O NH4

+ + OH−

reacts with acids to make NH4+ salts

– used as chemical fertilizers made by fixing N from N2 using the Haber-Bosch process

• hydrazine, N2H4 colorless liquid basic N2H4 + H2O N2H5

+ + OH−

powerful reducing agent

• hydrogen azide, HN3 acidic HN3 + H2O H3O+ + N3

−

thermodynamically unstable and decomposes explosively to its elements


Hydrazine


Oxides of Nitrogen• formed by reaction of N2 or NOx with O2

• all unstable and will eventually decompose into N2 and O2

• NO = nitrogen monoxide = nitric oxide important in living systems free radical

• NO2 = nitrogen dioxide 2 NO2 N2O4

red-brown gas free radical

• N2O = dinitrogen monoxide = nitrous oxide laughing gas made by heating ammonium nitrate NH4NO3 N2O + H2O oxidizing agent Mg + N2O N2 + MgO decomposes on heating 2 N2O 2 N2 + O2

pressurize food dispensers



Nitric Acid• HNO3 = nitric acid

produced by the Ostwald Process

4 NH3(g) + 5 O2(g) 4 NO(g) + 6 H2O(g)

2 NO(g) + O2(g) 2 NO2(g)

3 NO2(g) + H2O(l) 2 HNO3(l) + NO(g)

strong acidstrong oxidizing agentconcentrated = 70% by mass = 16 M

some HNO3 in bottle reacts with H2O to form NO2

main use to produce fertilizers and explosives

NH3(g) + HNO3(aq) NH4NO3(aq)


Nitrates and Nitrites• NO3

− = nitrateANFO = ammonium nitrate fuel oil

used as explosive in Oklahoma Cityammonium nitrate can decompose explosively

and other nitrates

2 NH4NO3 2 N2 + O2 + 4 H2Ometal nitrates used to give colors to fireworksvery soluble in wateroxidizing agent

• NO2− = nitrite

NaNO2 used as food preservative in processed meatskills botulism bacteriakeeps meat from browning when exposed to aircan form nitrosamines which may increase risk of colon cancer??


Phosphine• PH3

colorless, poisonous gas that smells like rotting fish formed by reacting metal phosphides with water

Ca3P2(s) + 6 H2O(l) 2 PH3(g) + 3 Ca(OH)2(aq)

also by reaction of wh P with H2O in basic solution

2 P4(s) + 9 H2O(l) + 3 OH−(aq) 5 PH3(g) + 3 H2PO4

−(aq)

decomposes on heating to elements

4 PH3(g) P4(s) + 6 H2(g)

reacts with acids to form PH4+ ion

does not form basic solutions


Phosphorus Halides• P4 can react directly with halogens to form PX3 and

PX5 compounds

• PX3 can react with water to form H3PO3

PX5 can react with water to form H3PO4

PCl3(l) + 3 H2O(l) H3PO3(aq) + 3 HCl(aq)

• PCl3 reacts with O2 to form POCl3(l) phosphorus oxychlorideother oxyhalides made by substitution on POCl3

• phosphous halide and oxyhalides are key starting materials in the production of many P compounds fertilizers, pesticides, oil-additives, fire-retardants,

surfactants


Phosphorus Oxides

• P4 reacts with O2 to make P4O6(s) or P4O10(s)

get P4O10 with excess O2


Phosphoric Acid and Phosphates• H3PO4 = phosphoric acid

white solid that melts at 42Cconcentrated = 85% by mass = 14.7 Mproduced by reacting P4O10 with water or the reaction

of Ca3(PO4)2 with sulfuric acid

P4O10(s) + 6 H2O(l) 4 H3PO4(aq)

Ca3(PO4)2(s) + 3 H2SO4(l) 3 CaSO4(s) + 2 H3PO4(qa)

used in rust removal, fertilizers, detergent additives and food preservativesodium pyrophosphate = Na4P2O7

sodium tripolyphosphate = Na5P3O10


Use of Phosphates in Food


Oxygen• 2s22p4

6 valence electrons

• stronger oxidizing agent than other 6A elementsused by living system to acquire energy

• second highest electronegativity (3.5)

• very high abundance in crust, and highest abundance of any element on Earth

• found in most common compounds


Elemental Oxygen• O2

nonpolar, colorless, odorless gas freezing point −183C at which it becomes a pale blue liquid slightly soluble in water

0.04 g/L mainly produced by fractional distillation of air

also by the electrolysis of water can be synthesized by heating metal oxides, chlorates, or nitrates

HgO(s) Hg(l) + O2(g) 2 NaNO3(s) 2 NaNO2(s) + O2(g) 2 KClO3(s) 2 KCl(s) + 3 O2(g)

used in high temperature combustion blast furnace, oxyacetylene torch

used to create artificial atmospheres divers, high-altitude flight

medical treatment lung disease, hyperbaric O2 to treat skin wounds


Oxides• reacts with most other elements to form oxides

both metals and nonmetals

• oxides containing O2− with −2 oxidation state most stable for small ions with high charge

• oxides containing O2− with −½ oxidation state

most stable for large ions with smaller charge


Ozone• O3

toxic, pungent, blue, diamagnetic gasdenser than O2

freezing point −112C, where it becomes a blue liquidsynthesized naturally from O2 through the activation by

ultraviolet lightmainly in the stratosphereprotecting the living Earth from harmful UV rays

spontaneously decomposes into O2

commercial use as a strong oxidizing agent and disinfectant formed in the troposphere by interaction of UV light and auto

exhaustoxidation damages skin, lungs, eyes, and cracks plastics and rubbers


Sulfur• large atom and weaker oxidizer than oxygen• often shows +2, +4, or +6 oxidation numbers in its

compounds, as well as −2• composes 0.06% of Earth’s crust• elemental sulfur found in a few natural deposits

some on the surface

• below ground recovered by the Frasch Processsuperheated water pumped down into deposit, melting the

sulfur and forcing it up the recovery pipe with the water

• also obtained from byproducts of several industrial processes


Natural Sulfur Deposit


Frasch Process


Allotropes of Sulfur• several crystalline forms

• the most common naturally occurring allotrope has S8 rings most others also ring structures of various sizes

• when heated to 112C, S8 melts to a yellow liquid with low viscosity

• when heated above 150C, rings start breaking and a dark brown viscous liquid forms darkest at 180C above 180C the liquid becomes less viscous

• if the hot liquid is quenched in cold water, a plastic amorphous solid forms that becomes brittle and hard on cooling


sulfur at ~150C sulfur at ~180C


Amorphous Sulfur


Other Sources of Sulfur• H2S(g) from oil and natural gas deposits

toxic gas (death > 100 ppm), smells like rotten eggs bond angle only 92.5 nonpolar S-H bond weaker and longer than O-H bond oxidized to elemental S through the Claus Process

2 H2S(g) + 2 O2(g) 2 SO2(g) + 2 H2O(g) 4 H2S(g) + 2 SO2(g) 6 S(s) + 4 H2O(g)

• FeS2 (iron pyrite) roasted in absence of air forming FeS(s) and S2(g)

• metal sulfides roasted in air to make SO2(g), which is later reduced react with acids to make H2S most insoluble in water used as bactericide and stop dandruff in shampoo


Metal Sulfides


Sulfur Dioxide• SO2

colorless, dense, acrid gas that is toxicproduced naturally by volcanic action and as a byproduct of

industrial processes including electrical generation by burning oil and coal, as well as

metal extractionacidic

SO2(g) + H2O(l) H2SO3(aq) forms acid rain in the air

2 SO2(g) + O2(g) + 2 H2O(l) 2 H2SO4(aq) removed from stack by scrubbing with limestone

CaCO3(s) CaO(s) + O2(g) 2 CaO(g) + 2 SO2(g) + O2(g) 2 CaSO4(g)

used to treat fruits and vegetables as a preservative


Sulfuric Acid

• most produced chemical in the world• strong acid, good oxidizing agent, dehydrating agent• used in production of fertilizers, dyes, petrochemicals,

paints, plastics, explosives, batteries, steel, and detergents

• melting point 10.4C, boiling point 337Coily, dense liquid at room temperature

• reacts vigorously and exothermically with water“you always oughter(sic) add acid to water”


Dehydration of Sucrose

C12H22O11(s) + H2SO4(l) 12 C(s) + 11 H2O(g) + H2SO4(aq)


Production of H2SO4

• contact process• step 1: combustion of elemental S

complete using V2O5 catalyst

S(g) + O2(g) SO2(g)

2 SO2(g) + O2(g) 2 SO3(g)

• step 2: absorbing the SO2 into conc. H2SO4 to form oleum, H2S2O7

SO3(g) + H2SO4(l) H2S2O7(l)

• step 3: dissolve the oleum in waterH2S2O7(l) + H2O(l) 2 H2SO4(aq)


Halogens• most reactive nonmetal group, never found in

elemental form in nature• come from dissolved salts in seawater

except fluorine, which comes from minerals fluorospar (CaF2) and fluoroapatite [Ca10F2(PO4)6]

• atomic radius increases down the column• most electronegative element in its period, decreasing

down the column• fluorine only has oxidation states of -1 or 0, others

have oxidation states ranging from -1 to +7


Properties of the Halogens


Fluorine• F2 is a yellow-green toxic gas• F2 is the most reactive nonmetal and forms binary compounds

with every element except He, Ne, and Ar including XeF2, XeF6, XeOF4, KrF2

so reactive it reacts with other elements of low reactivity resulting in flames

even reacts with the very unreactive asbestos and glass stored in Fe, Cu, or Ni containers because the metal fluoride that forms coats

the surface protecting the rest of the metal

• F2 bond weakest of the X2 bonds, allowing reactions to be more exothermic

• small ion size of F− leads to large lattice energies in ionic compounds

• produced by the electrolysis of HF


Hydrofluoric Acid• HF

produced by the reaction of fluorospar with H2SO4

CaF2(s) + H2SO4(l) CaSO4(s) + 2 HF(g) crystalline HF is zig-zag chains HF is weak acid, Ka = 6.8 x 10-4 at 25C F− can combine with HF to form complex ion HF2

−

with bridging H strong oxidizing agent

strong enough to react with glass, so generally stored in plastic used to etch glass

SiO2(g) + 4 HF(aq) SiF4(g) + H2O(l) very toxic because it penetrates tissues and reacts with internal organs

and bones


Halogen Compounds• form ionic compounds with metals and molecular compounds

having covalent bonds with nonmetals• halogens can also form compounds with other halogens – called

interhalides for interhalides, the larger has lower electronegativity – so it is central in

the molecule; with a number of more electronegative halides attached general formula ABn where n can be 1, 3, 5, or 7

most common AB or AB3; only AB5 has B = F, IF7 only known n = 7

only ClF3 used industrially to produce UF6 in nuclear fuel enrichment

• most halogen oxides are unstable tend to be explosive OF2 only compound with O = +2 oxidation state ClO2(g) is strong oxidizer used to bleach flour and wood pulp

explosive – so diluted with CO2 and N2

produced by oxidation of NaClO2 with Cl2 or the reduction of NaClO3 with HCl2 NaClO2 + Cl2 2 NaCl + 2 ClO2

2 NaClO3 + 4 HCl 2 ClO2 + 2 H2O + 2 NaCl

Documents

Chapter 22 Chemistry of the Nonmetals