Gián án Chapter 22 Chemistry of the 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 ioni

• nanotubes – long, thin, hollow cylinders of atoms

• carbon nanotube = sp2 C in fused hexagonal rings

Insulated Nanowire

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

compounds

 often 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

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

• silicates are covalent atomic solids of Si and O

and minor amounts of other elements

found in rocks, soils, and clays

silicates 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 hybridization

tetrahedral shape

Si-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

• a 3-dimensional covalent

network of SiO4 tetrahedrons

• generally called silica

• formula unit is SiO2

• when heated above 1500C and

cooled quickly, get amorphous

silica which we call glass

• Al substitutes for Si in some of the lattice sites

• SiO2 becomes AlO2−

• the negative charge is countered by the inclusion

of a cation

Albite = 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 SiO44−

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

Single Chain Silicates

• if the SiO44− units link as long

chains with shared O, the

structure is called a pyroxene

• formula unit SiO3

2-• chains held together by ionic

bonding to metal cations

between the chains

 diopside = CaMg(SiO3)2 where

Ca and Mg occupy lattice points

between the chains

Double Chain Silicates

• some silicates have 2

chains bonded together

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

Mica: a Phyllosilicate

Silicate Structures

• less than 0.001% in Earth’s crust, but

found concentrated in certain areas

 almost always found in compounds with O

 borax = Na 2 [B 4 O 5 (OH) 4 ]8H 2 O

 kernite = Na2[B4O5(OH)4]3H2O

 colemanite = Ca 2 B 6 O 11 5H 2 O

• used in glass manufacturing –

borosilicate glass = Pyrex

• used in control rods of nuclear reactors

Boron Trihalides

• BX3

• sp2 B

trigonal planar, 120 bond angles

forms single bonds that are shorter and stronger than

sp3 C

some 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 ring

melts at 450C

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

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

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

• 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 sp2carbons in a 2-dimensional array

sheets attracted by weak dispersion forces

• fullerenes consist of 5 and 6 member carbon

rings fused into icosahedral spheres of at least 60

Crystalline Allotropes of Carbon

Diamond Graphite

Buckminster-fullerene, C 60

Color clear-blue black black Density, g/cm 3 3.53 2.25 1.65 Hardness, Mohs Scale 10 0.5

Allotropes of Carbon - Diamond

Inert to Common Acids

Inert to Common Bases

Negative Electron Affinity

Allotropes of Carbon - Graphite

Soft and Greasy Feeling

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

 carbon black is finely divided form of carbon that is a component of soot

Allotropes of Carbon - Buckminsterfullerene

Reacts with alkali metals

Behavior more aliphatic than

aromatic

• long hollow tubes constructed of fused C6 rings

• electrical conductors

• can incorporate metals and other small

molecules and elements

used to stabilize unstable molecules

• single-walled nanotubes (SWNT) have one layer

of fused rings

• multi-walled nanotubes (MWNT) have

Nanotubes

Nanocars

element

 generally alkali or alkali earth metals

 often dicarbide ion, C 22− (aka acetylide ion)

 react with water to form acetylene, C2H2

low-electronegativity nonmetal or metalloid

 silicon carbide, SiC (aka carborundum)

Calcium Carbide

Fe3C 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 HCO 3− ions

 triple point −57C and 5.1 atm

 liquid CO 2 doesn’t exist at atmospheric pressure

 solid CO 2 = dry ice

• CO

 colorless, odorless, tasteless gas

 relatively reactive

 2 CO + O 2  2 CO 2

– burns with a blue flame

 reduces many nonmetals

– CO + Cl

 doesn’t decompose on heating

• all carbonate solutions are basic in water

 due to CO32− + H2O  OH − + HCO32−

• baking soda = NaHCO3

 decomposes on heating to Na CO , H O and CO

Elemental Phosphorus

• P

 white phosphorus

 white, soft, waxy solid that is flammable and toxic

 stored under water to prevent spontaneous combustion

 2 Ca 3 (PO 4 ) 2 (apatite) + 6 SiO 2 + 10 C  P 4 (g, wh) + 6 CaSiO 3 + 10 CO

 tetrahedron with small angles 60

 red phosphorus

 formed by heating white P to about 300C 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

White Phosphorus Red Phosphorus

Hydrides of Nitrogen

• ammonia, NH 3

 pungent gas

 basic NH3 + H2O  NH4+ + OH −

 reacts with acids to make NH 4+ salts

– used as chemical fertilizers

 made by fixing N from N2 using the Haber-Bosch process

Hydrazine

Oxides of Nitrogen

• formed by reaction of N2 or NOx with O2

• all unstable and will eventually decompose into N 2 and O 2

• NO = nitrogen monoxide = nitric oxide

 important in living systems

 made by heating ammonium nitrate NH4NO3  N2O + H2O

 oxidizing agent Mg + N2O  N2 + MgO

 decomposes on heating 2 N O  2 N + O

Nitric Acid

• HNO3 = nitric acid

 some HNO 3 in bottle reacts with H 2 O to form NO 2

 main use to produce fertilizers and explosives

Nitrates and Nitrites

• NO3− = nitrate

 ANFO = ammonium nitrate fuel oil

 used as explosive in Oklahoma City

 ammonium nitrate can decompose explosively

 and other nitrates

2 NH4NO3  2 N2 + O2 + 4 H2O

 metal nitrates used to give colors to fireworks

 very soluble in water

 oxidizing agent

• NO2− = nitrite

 NaNO 2 used as food preservative in processed meats

 kills botulism bacteria

• PH3

 colorless, poisonous gas that smells like rotting fish

 formed by reacting metal phosphides with water

 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 oxychloride

 other oxyhalides made by substitution on POCl3

• phosphous halide and oxyhalides are key starting

materials in the production of many P compounds

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 42C

concentrated = 85% by mass = 14.7 M

produced 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 preservative

Use of Phosphates in Food

6 valence electrons

• stronger oxidizing agent than other 6A elements

used 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 −183C 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

• 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

• O3

 toxic, pungent, blue, diamagnetic gas

 denser than O2

 freezing point −112C, where it becomes a blue liquid

 synthesized naturally from O2 through the activation by

ultraviolet light

 mainly in the stratosphere

 protecting the living Earth from harmful UV rays

 commercial use as a strong oxidizing agent and disinfectant

 formed in the troposphere by interaction of UV light and auto exhaust

• 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 Process

 superheated 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

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 112C, S 8 melts to a yellow liquid with low

viscosity

• when heated above 150C, rings start breaking and a dark brown viscous liquid forms

 darkest at 180C

 above 180C 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 ~150C sulfur at ~180C

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 H 2 S(g) + 2 O2(g)  2 SO2(g) + 2 H 2 O(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

Metal Sulfides

Sulfur Dioxide

• SO2

 colorless, dense, acrid gas that is toxic

 produced naturally by volcanic action and as a byproduct of industrial processes

 including electrical generation by burning oil and coal, as well as metal extraction

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 10.4C, boiling point 337C

 oily, dense liquid at room temperature

• reacts vigorously and exothermically with water

 “you always oughter(sic) add acid to water”

Dehydration of Sucrose

Production of H2SO4

• step 1: combustion of elemental S

 complete using V2O5 catalyst

• most reactive nonmetal group

• never found in elemental form in nature

• come from dissolved salts in seawater

 except fluorine, which comes from minerals fluorospar

(CaF 2 ) and fluoroapatite [Ca 10 F 2 (PO 4 ) 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

• F2 is a yellow-green toxic gas

• F 2 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

 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) + H 2 O(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 AB 3 ; only AB 5 has B = F, IF 7 only known n = 7

 only ClF3 used industrially

 to produce UF 6 in nuclear fuel enrichment

• most halogen oxides are unstable

 tend to be explosive

 OF 2 only compound with O = +2 oxidation state

 ClO2(g) is strong oxidizer used to bleach flour and wood pulp

 explosive – so diluted with CO 2 and N 2