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Chemistry: A Molecular Approach, 1st Ed.
Nivaldo Tro
Chapter 22
Chemistry
of the
Nonmetals
Roy Kennedy
Massachusetts Bay Community College
Wellesley Hills, MA
2008, Prentice Hall
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 C
 similar size to C
 average electronegativity of B & N about the same as C
 electrical insulators
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Properties of BN and C
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Main Group Nonmetals
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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
anion and covalent
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Insulated Nanowire
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Silicates
• 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
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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
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Quartz
• 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
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Aluminosilicates
• 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
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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,
it forms an orthosilicate
 willemite = Zn2SiO4
• when two SiO4 units share an O, they form structures
called pyrosilicates with the anion formula Si2O76−
 hardystonite =Ca2ZnSi2O7
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Single Chain Silicates
• if the SiO44− units link as long
•
•
chains with shared O, the
structure is called a pyroxene
formula unit SiO32chains 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
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Double Chain Silicates
• some silicates have 2
chains bonded together
at ½ the tetrahedra –
these are called
amphiboles
• often results in fibrous
minerals
asbestos
tremolite asbestos =
Ca2(OH)2Mg5(Si4O11)2
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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
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Mica: a Phyllosilicate
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Silicate Structures
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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 = Ca2B6O115H2O
• used in glass manufacturing –
borosilicate glass = Pyrex
• used in control rods of nuclear reactors
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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
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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
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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
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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
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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 array
sheets attracted by weak dispersion forces
• fullerenes consist of 5 and 6 member carbon
rings fused into icosahedral spheres of at least
60 C
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Crystalline Allotropes of Carbon
Diamond
Graphite
Buckminsterfullerene, C60
Color
clear-blue
black
black
Density, g/cm3
3.53
2.25
1.65
Hardness, Mohs Scale
10
0.5
Electrical Conductivity, (m•cm)-1
~10-11
7.3 x 10-4
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)
~10-14
23
Allotropes of Carbon - Diamond
Inert to Common Acids
Inert to Common Bases
Negative Electron Affinity
Transparent
Hardest
Best Thermal Conductor
Least Compressible
Stiffest
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Allotropes of Carbon - Graphite
Soft and Greasy Feeling
Solid Lubricant
Pencil “Lead”
Conducts Electricity
Reacts with Acids and
Oxidizing Agents
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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
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Allotropes of Carbon Buckminsterfullerene
Sublimes between 800°C
Insoluble in water
Soluble in toluene
Stable in air
Requires temps > 1000°C to
decompose
High electronegativity
Reacts with alkali metals
Behavior more aliphatic than
aromatic
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Nanotubes
• 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
concentric layers of fused rings
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Nanotubes
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Nanocars
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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, C22− (aka acetylide ion)
 react with water to form acetylene, C2H2
• covalent carbides are compounds of C with a lowelectronegativity 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
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Calcium Carbide
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Cementite
Fe3C regions found in steel
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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 −57C and 5.1 atm
• CO
 liquid CO2 doesn’t exist at atmospheric pressure
 solid CO2 = dry ice
 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)
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Carbonates
• solubility of CO2 in H2O due to carbonate formation
 CO2 + H2O  H2CO3
 H2CO3 + H2O  H3O+ + HCO3−
 HCO3− + H2O  H3O+ + CO32−
• washing soda = Na2CO310H2O
 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 Na2CO3, H2O and CO2
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Elemental Nitrogen
• N2
78% of atmosphere
purified by distillation of liquid air, or
filtering air through zeolites
very stable, very unreactive
NN
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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 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
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Phosphorus
White
Red Phosphorus
Phosphorus
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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
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Hydrazine
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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
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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 acid
 strong oxidizing agent
 concentrated = 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)
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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
 NaNO2 used as food preservative in processed meats
 kills botulism bacteria
 keeps meat from browning when exposed to air
 can form nitrosamines which may increase risk of colon cancer??
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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
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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
 fertilizers, pesticides, oil-additives, fire-retardants,
surfactants
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Phosphorus Oxides
• P4 reacts with O2 to make P4O6(s) or P4O10(s)
get P4O10 with excess O2
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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
sodium pyrophosphate = Na4P2O7
sodium tripolyphosphate = Na5P3O10
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Use of Phosphates in Food
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Oxygen
• 2s22p4
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
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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
 medical treatment
 lung disease, hyperbaric O2 to treat skin wounds
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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
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Ozone
• 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
 spontaneously decomposes into O2
 commercial use as a strong oxidizing agent and disinfectant
 formed in the troposphere by interaction of UV light and auto
exhaust
 oxidation damages skin, lungs, eyes, and cracks plastics and rubbers
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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 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
processes
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Natural Sulfur Deposit
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Frasch Process
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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, S8 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
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sulfur at ~150C
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sulfur at ~180C
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Amorphous Sulfur
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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
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Metal Sulfides
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• SO2
Sulfur Dioxide
 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
 acidic
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
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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”
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Dehydration of Sucrose
C12H22O11(s) + H2SO4(l)  12 C(s) + 11 H2O(g) + H2SO4(aq)
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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 water
H2S2O7(l) + H2O(l)  2 H2SO4(aq)
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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
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Properties of the Halogens
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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
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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 25C
 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
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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 HCl
2 NaClO2 + Cl2  2 NaCl + 2 ClO2
2 NaClO3 + 4 HCl  2 ClO2 + 2 H2O + 2 NaCl
Tro, Chemistry: A Molecular Approach
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