Chapter 22 Chemistry of the Nonmetals
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Transcript Chapter 22 Chemistry of the Nonmetals
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 1500C 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 = Ca2B6O115H2O
• 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 450C
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
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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 −57C 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 = Na2CO310H2O
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
NN
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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 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
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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 42C
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 −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
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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 −112C, 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 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
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sulfur at ~150C
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sulfur at ~180C
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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 10.4C, boiling point 337C
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 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
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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
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