Lecture Presentation
Chapter 22
Chemistry of the
Nonmetals
Sherril Soman
Grand Valley State University
© 2014 Pearson Education, Inc.
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 and N about the same as C
Electrical insulators
One day wires thinner than
a piece of hair may be
made of C nanotubes
surrounded by BN
nanotubes.
© 2014 Pearson Education, Inc.
Properties of BN and C
© 2014 Pearson Education, Inc.
Main Group Nonmetals
•
•
•
•
Groups 3A to 8A
Filling p orbitals of the valence shell
Form anions when reacting with metals
Easily reduced
– Therefore, the element is an oxidizing agent.
© 2014 Pearson Education, Inc.
Main Group Nonmetals
© 2014 Pearson Education, Inc.
The p Block
• Nonmetals on right of p block form anions in ionic
compounds.
• Elements on left side 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.
© 2014 Pearson Education, Inc.
Periodic Trends
• Atomic radius decreases across the period and
increases down the group.
• Ionization energy increases across the period and
decreases down the group.
• Metallic character decreases across the period
and increases down the group.
• Electronegativity increases across the period and
decreases down the group.
© 2014 Pearson Education, Inc.
Bonding in p Block Elements
• Bonding tendency changes across the period
• Left-side p block elements tend to form cations
and covalent bonds.
– Lower on the group = more likely cations
• Central p block elements tend to form
covalent bonds.
• Right side p block elements tend to form anions
and covalent bonds.
© 2014 Pearson Education, Inc.
Reactivity of p Block Elements
• Because the electronegativity and bonding vary so
widely across, their reactivity varies as well.
• Some of the elements on the lower left of the p
block are metallic.
– Tl, Pb
• Some of the elements in the center form covalent
network solids that are highly unreactive.
– C, Si
• Group 8A elements are highly unreactive gases.
• Some elements on the upper right of the p block
are highly reactive nonmetals.
© 2014 Pearson Education, Inc.
Silicates
• The most abundant elements of Earth’s crust are
O and Si.
• Silicates are covalent atomic solids of Si, 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.
© 2014 Pearson Education, Inc.
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
© 2014 Pearson Education, Inc.
Quartz
• A three-dimensional covalent
network of SiO4 tetrahedrons
• Generally called silica
• Formula unit is SiO2.
• When heated above 1500
°C and cooled quickly, it
forms amorphous silica,
which we call glass.
© 2014 Pearson Education, Inc.
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
© 2014 Pearson Education, Inc.
Silicates Made of Individual Units
• O of SiO4 can pick up electrons from metal to form
SiO44− ions.
• If the SiO44− is made up of 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
© 2014 Pearson Education, Inc.
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 chains
– Diopside = CaMg(SiO3)2,
where Ca and Mg occupy
lattice points between the
chains.
© 2014 Pearson Education, Inc.
Double Chain Silicates
• Some silicates have two chains bonded together
at half the tetrahedra – these are called
amphiboles.
• Often results in fibrous minerals
– Asbestos
– Tremolite asbestos = Ca2(OH)2Mg5(Si4O11)2
© 2014 Pearson Education, Inc.
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
© 2014 Pearson Education, Inc.
Mica: A Phyllosilicate
© 2014 Pearson Education, Inc.
Silicate Structures
© 2014 Pearson Education, Inc.
Boron
• Metalloid
• At least five 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
– Reduces the expansion of glass when heated
• Used in control rods of nuclear reactors
© 2014 Pearson Education, Inc.
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
© 2014 Pearson Education, Inc.
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
• Melting dissolves many metal
oxides and silicon oxides to
form glasses of different
compositions.
© 2014 Pearson Education, Inc.
Boranes: closo-Boranes
• Compounds of B and H
• Used as reagent in hydrogenation of C═C
• closo-Boranes have the formula BnHn2− and form
closed polyhedra with a BH unit at each vertex.
© 2014 Pearson Education, Inc.
Boranes: nido-Boranes and
arachno-Boranes
• nido-Boranes have the formula BnHn + 4 consisting of
cage B missing one corner.
• arachno-Boranes have the formula BnHn + 6
consisting of cage B missing two or three corners.
© 2014 Pearson Education, Inc.
Carbon
• Exhibits the most versatile bonding of all the
elements.
• The crystalline forms include diamond, graphite, and
the fullerenes.
• Diamond structure consists of tetrahedral sp3
carbons in a three-dimensional array.
• Graphite structures consist of trigonal planar sp2
carbons in a two-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.
© 2014 Pearson Education, Inc.
Noncrystalline Forms of Carbon: Coal
• Coal is a mixture of hydrocarbons and carbon-rich
particles that comes from the 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
– Used to reduce iron ore to iron in a blast furnace
© 2014 Pearson Education, Inc.
Types of Coal
© 2014 Pearson Education, Inc.
Noncrystalline Forms of Carbon:
Charcoal and Soot
• Heating wood in the absence of air forms
charcoal.
• Activated carbon is finely divided charcoal
formed by heating amorphous carbon in steam.
– Used to adsorb other molecules in filters
• Soot is composed of hydrocarbons from
incomplete combustion.
• Carbon black is a finely divided form of carbon
that is a component of soot.
– Used as rubber strengthener
© 2014 Pearson Education, Inc.
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, (mΩ • cm)−1
~10−11
7.3 × 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)
© 2014 Pearson Education, Inc.
~10−14
Allotropes of Carbon: Diamond
•
•
•
•
•
•
•
•
•
Inert to common acids
Inert to common bases
Negative electron affinity
Transparent
Hardest
Best thermal conductor
Least compressible
Stiffest
Uses include jewelry,
cutting tools, nonabrasive
coatings
A covalent network of
tetrahedral C atoms in a
face-centered cubic crystal with
C in ½ the tetrahedral holes
© 2014 Pearson Education, Inc.
Allotropes of Carbon: Graphite
• Soft and greasy feeling
• Conducts electricity
• Reacts with acids and
oxidizing agents
• Uses include pencil
“lead,” solid lubricant
each layer consists of C atoms
in hexagonal rings
the layers are not bonded so
can slide past each other
© 2014 Pearson Education, Inc.
Allotropes of Carbon: Buckminsterfullerene
•
•
•
•
•
Sublimes above 800 °C
Insoluble in water
Soluble in toluene
Stable in air
Requires temps > 1000
°C to decompose
• High electronegativity
• Reacts with alkali metals
© 2014 Pearson Education, Inc.
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.
• Multiwalled nanotubes (MWNT) have concentric
layers of fused rings.
© 2014 Pearson Education, Inc.
Nanotubes
© 2014 Pearson Education, Inc.
Nanocars
© 2014 Pearson Education, Inc.
Properties and Uses of Nanotubes
• Nanotubes – 100 times the strength of steel, but
only one-sixteenth the density
• Used in applications where high strength and low
weight are important
• “Graphite” shafts on golf clubs
• Competition bicycle frames
• Composite materials with polymers to increase
their strength
• Solar cells
• Improve ultracapacitors
– For storing electrical energy
© 2014 Pearson Education, Inc.
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
© 2014 Pearson Education, Inc.
Calcium Carbide
© 2014 Pearson Education, Inc.
Cementite Fe3C Regions Found in Steel
© 2014 Pearson Education, Inc.
Oxides of Carbon: Carbon Dioxide
• CO2
• 0.04% in atmosphere
– Increased by 25% over the past century
• High solubility in water compared to most
nonpolar gases
– Due to reaction with water to form HCO3− ions
• Triple point −57 °C and 5.1 atm
– Liquid CO2 doesn’t exist at atmospheric pressure
• Solid CO2 = dry ice
© 2014 Pearson Education, Inc.
Oxides of Carbon: Carbon Monoxide
• CO
• Colorless, odorless, tasteless gas
– Boils at −192 °C
• Toxic
– Binds to hemoglobin
• Low solubility in water
• Reactive compared to CO2
– 2 CO + O2 → 2 CO2
– Burns with a blue flame
• Strong reducing agent
– CO + CuO → Cu + CO2
– CO + Cl2 → COCl2 (phosgene)
© 2014 Pearson Education, Inc.
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
– Heating drives off waters of hydration
• All carbonate solutions are basic in water.
– Due to CO32− + H2O ⇔ OH− + HCO32−
• Baking soda = NaHCO3
– Decomposes on heating to Na2CO3, H2O, and CO2
– Baking powder is a mixture of NaHCO3 and an acid that
is activated when water is added.
© 2014 Pearson Education, Inc.
Elemental Nitrogen, N2
• 78% of atmosphere by volume
• MP −210 °C; BP −196 °C
• Purified by distillation of liquid air, or filtering air
through zeolites
• Very stable, very unreactive
– N≡N
– Heating with O2 gives NO in low yields.
– Heating with H2 gives NH3 in low yields.
• Used as “inert” atmosphere for reactions,
where oxygen causes side reactions
© 2014 Pearson Education, Inc.
Elemental Phosphorus
• First isolated by distilling urine
– Major source of P for many years
• Today the major source is heating the mineral
apatite in an electric furnace with sand and coke.
2 Ca3(PO4)2(s) (apatite) + 6 SiO2(s) + 10 C(s) → P4(g, wh) +
6 CaSiO3(l) + 10 CO(g)
© 2014 Pearson Education, Inc.
Three Allotropic Forms of Phosphorus:
White Phosphorus, P4
• White, soft, waxy solid
• Flammable and toxic
• Stored under water to prevent spontaneous
combustion
• Tetrahedron with small angles
– 60°
© 2014 Pearson Education, Inc.
Three Allotropic Forms of Phosphorus:
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
© 2014 Pearson Education, Inc.
Three Allotropic Forms of Phosphorus:
Black Phosphorus
• Formed by heating white P under pressure
• Most thermodynamically stable form; therefore,
least reactive
• Layered structure similar to graphite
© 2014 Pearson Education, Inc.
Compounds of Nitrogen
• Nitrogen forms binary
compounds with most
nonmetals and with
some metals.
• Nitrogen exhibits oxidation
states ranging from −3 up to
+5 in its covalent compounds.
– Only shows positive oxidation
states when combined with O or F.
© 2014 Pearson Education, Inc.
Hydrides of Nitrogen: Ammonia, NH3
• Pungent gas
• Basic
– NH3 + H2O ⇔ NH4+ + OH−
• Reacts with acids to make NH4+ salts
– The major source of chemical fertilizers
• Made by fixing N from N2 using the Haber–
Bosch process
– Fixing nitrogen means converting N2(g) into NH3(g).
– Nitrogen fixation is also accomplished by some bacteria
that live on legumes.
© 2014 Pearson Education, Inc.
Haber–Bosch Process
• N2(g) + 3 H2(g) → 2 NH3(g)
• Equilibrium constant small and
reaction slow at 25 °C
• The process uses a form of
magnetite, Fe3O4, as a catalyst
• ΔH = −46.2 kJ/mol NH3
• Pressure raised to about 200 atm
• Temperature 300 to 550 °C
• 15% yield on each pass
• Gas mixture cooled and NH3
condensed to remove it, then
unreacted N2 and H2 recycled
© 2014 Pearson Education, Inc.
Temperature,
°C
K,
×10−5
300
434
400
16.4
450
4.51
500
1.45
550
0.538
600
0.225
Hydrides of Nitrogen: Hydrazine, N2H4
• Colorless liquid
– MP = 2 °C, BP = 114 °C
• Trigonal pyramid around each N
• Weakly basic
– N2H4 + H2O ⇔ N2H5+ + OH−
• ΔHf ° = +51 kJ/mol
– Unstable
• Powerful reducing agent
– Used as fuel for the Apollo spacecraft
© 2014 Pearson Education, Inc.
Hydrazine as a Reducing Agent
© 2014 Pearson Education, Inc.
Hydrides of Nitrogen: Hydrogen Azide, HN3
• Liquid
– BP = 37 °C
• Weakly acidic (hydrazoic acid)
– HN3 + H2O ⇔ H3O+ + N3−
•
•
•
•
Azide ion N3− isoelectronic with CO
Linear with both N—N bonds = 116 pm
Azide ion thermodynamically unstable
Ionic compounds of metals with azide decompose
explosively to the elements.
– Explosive decomposition of 2 NaN3 → 2 Na + 3 N2 is
used to deploy air bags in cars.
© 2014 Pearson Education, Inc.
Oxides of Nitrogen
• Formed by reaction of N2 or NOx with O2
• Oxidation of N2 requires high temperatures
because of the strength of the N≡N.
– N2(g) + O2(g) → 2 NO(g) H = +91 kJ/mol NO
– Reaction both slow and unfavorable by
thermodynamics
• K = 10−30 at 25 °C and 2 × 10−2 at 2300 °C
– Required conditions found in lightning storms and
car engines
• All NOx are unstable and will eventually
decompose into N2 and O2.
– But the decomposition is often slow.
© 2014 Pearson Education, Inc.
Oxides of Nitrogen
Name
Physical
State
BP, °C
N2O
dinitrogen
monoxide
(nitrous oxide)
colorless gas
−151.8
NO
nitrogen
monoxide
(nitric oxide)
colorless gas
+3
N2O3
dinitrogen
trioxide
deep blue
liquid at −80
°C
+4
NO2
nitrogen dioxide
brown gas
+4
N2O4
dinitrogen
tetroxide
white solid
21.2
+5
N2O5
dinitrogen
pentoxide
white solid
47
Oxidation
Formula
State
+1
+2
© 2014 Pearson Education, Inc.
−88.5
?
Oxides of Nitrogen: Nitrogen monoxide, NO
•
•
•
•
•
Also known as nitric oxide
Colorless, reactive gas
Very slightly soluble in water
Toxic
Important in living systems
–
–
–
–
–
–
Neurotransmitter
Dilates blood vessels
Aids memory
Important in digestion
Plays major role in inducing erections
Plays major role in inducing uterine contractions
• Free radical
• Bond order = 2.5; bond length = 115 pm
© 2014 Pearson Education, Inc.
Oxides of Nitrogen: Nitrogen monoxide, NO
• Prepared industrially by the oxidation of NH3
4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g)
– Part of the Ostwald process
– Requires a Pt–Rh catalyst heated to about 900 °C
• Because of its low oxidation state (+2), NO is a
good reducing agent.
NO(g) + CO2(g) → NO2(g) + CO(g)
2 NO(g) + Cl2(g) → 2 NOCl (g)
• But can also be an oxidizing agent
2 NO(g) + 2 H2(g) → N2(g) + 2 H2O(l)
P4(s) + 6 NO(g) → P4O6(s) + 3 N2(g)
© 2014 Pearson Education, Inc.
Oxides of Nitrogen: Nitrogen dioxide, NO2
• Reddish brown gas
• Found in equilibrium with colorless gas N2O4
2 NO2(g)
N2O4(g)
– ΔH = −57 kJ/mol N2O4
– At 21 °C, the resulting mixture is a reddish brown liquid.
• Free radical
• N—O bond length = 120 pm, bond angle = 134°
© 2014 Pearson Education, Inc.
Oxides of Nitrogen: Nitrogen dioxide, NO2
• Prepared by the oxidation of NO(g)
2 NO(g) + O2(g) → 2 NO2(g)
• Or the thermal decomposition of metal nitrates
2 Pb(NO3)2(s) → 2 PbO(s) + 4 NO2(g) + O2(g)
• NO2 is a strong oxidizing agent.
• Hydrocarbons will combust when heated with it.
– Sometimes explosively
CH4(g) + 4 NO2(g) → CO2(g) + 2 H2O(g) + 4 NO(g)
© 2014 Pearson Education, Inc.
Oxides of Nitrogen: Dinitrogen
tetroxide, N2O4
• Colorless gas that condenses at 21 °C and
freezes at 11 °C to a white solid.
• Toxic
– Poisoned astronauts on Apollo–Soyuz Test Project
during re-entry
•
•
•
•
Diamagnetic
N—O = 119 pm, N—N = 178 pm
O—N—O bond angle = 134°
Oxidizer in rocket fuel
– Oxidizes hydrazine, N2H4
© 2014 Pearson Education, Inc.
•
•
•
•
•
Oxides of Nitrogen: Dinitrogen
monoxide, N2O
Also known as nitrous oxide
Colorless gas, pleasant smell, sweet taste
Nonflammable
Anesthetic (laughing gas)
Produced by heating ammonium nitrate
NH4NO3(g) → N2O(g) + H2O(l)
• Oxidizing agent
Mg(s) + N2O(g) → N2(g) + MgO(s)
– Used as oxidizer in hybrid rocket and race car fuels
• Decomposes on heating
2 N2O(g) → 2 N2(g) + O2(g)
• Used to pressurize food dispensers
© 2014 Pearson Education, Inc.
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 a bottle reacts with H2O to form NO2.
– Main use to produce fertilizers and explosives
NH3(g) + HNO3(aq) → NH4NO3(aq)
© 2014 Pearson Education, Inc.
Nitrates and Nitrites
• NO3− = nitrate
– ANFO = ammonium nitrate fuel oil
• Used as explosive in Oklahoma City
– Ammonium nitrate (as well as some other nitrates) can
decompose explosively
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
© 2014 Pearson Education, Inc.
Hydrides of Phosphorus: 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 white 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
© 2014 Pearson Education, Inc.
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
• Phosphorus halide and oxyhalides are key starting
materials in the production of many P compounds.
– Fertilizers, pesticides, oil additives, fire retardants,
surfactants
© 2014 Pearson Education, Inc.
Phosphorus Oxides
• P4 reacts with O2 to make P4O6(s) or P4O10(s).
– Get P4O10 with excess O2
© 2014 Pearson Education, Inc.
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(aq)
– Used in rust removal, fertilizers, detergent additives, and food
preservation
• Sodium pyrophosphate = Na4P2O7
• Sodium tripolyphosphate = Na5P3O10
© 2014 Pearson Education, Inc.
Use of Phosphates in Food
© 2014 Pearson Education, Inc.
Oxygen
• 2s22p4
– Six 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
© 2014 Pearson Education, Inc.
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)
© 2014 Pearson Education, Inc.
Elemental Oxygen, O2
• 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
© 2014 Pearson Education, Inc.
Oxides
• Reacts with most other elements to form oxides
– Both metals and nonmetals
• Oxides containing O2− with −2 oxidation state
– Most stable for small cations with large charge
• Oxides containing O22− with −1 oxidation state
• Oxides containing O2− with −½ oxidation state
– Most stable for large cations with small charge
© 2014 Pearson Education, Inc.
Ozone, O3
•
•
•
•
Toxic, pungent, blue, diamagnetic gas
Denser than O2
Freezing point −112 °C, where it becomes a blue liquid
Synthesized naturally from O2 when it is activated by
ultraviolet light: 3 O2(g) → 2 O3(g)
– Mainly in the stratosphere
– Protecting the living Earth from harmful UV rays
• Will spontaneously decompose 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.
© 2014 Pearson Education, Inc.
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
© 2014 Pearson Education, Inc.
Natural Sulfur Deposit
© 2014 Pearson Education, Inc.
Frasch Process
© 2014 Pearson Education, Inc.
Allotropes of Sulfur
• Several crystalline forms
• The most common naturally occurring allotrope
has S8 rings.
– Most others also have 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.
© 2014 Pearson Education, Inc.
Molten Sulfur
© 2014 Pearson Education, Inc.
Amorphous Sulfur
© 2014 Pearson Education, Inc.
Other Sources of Sulfur: Hydrogen Sulfide
•
•
•
•
•
•
•
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)
© 2014 Pearson Education, Inc.
Other Sources of Sulfur: 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 to stop dandruff in shampoo
FeS2 (iron pyrite)
– Roasted in absence of air forming FeS(s) and S2(g)
© 2014 Pearson Education, Inc.
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
• 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
© 2014 Pearson Education, Inc.
Sulfuric Acid, H2SO4
• Oily, dense liquid at room temperature
•
•
•
•
•
•
Melting point 10.4 °C; boiling point 337 °C
Most produced chemical in the world
Strong acid
Good oxidizing agent
Dehydrating agent
Reacts vigorously and exothermically with water
–“You always oughter(sic) add acid to water.”
• Used in production of fertilizers, dyes,
petrochemicals, paints, plastics, explosives,
batteries, steel, and detergents
© 2014 Pearson Education, Inc.
Dehydration of Sucrose
C12H22O11(s) + H2SO4(l) → 12 C(s) + 11 H2O(g) + H2SO4(aq)
© 2014 Pearson Education, Inc.
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 concentrated 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)
© 2014 Pearson Education, Inc.
Halogens
• Most reactive nonmetal group, never found in
elemental form in nature
• Source is dissolved salts in seawater
– Except fluorine, which comes from the 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, while
others have oxidation states ranging from −1 to +7.
© 2014 Pearson Education, Inc.
Properties of the Halogens
Halogen
F2
Cl2
MP,
Color and
BP, °C
State
°C
−219
−101
−188
pale yellow
gas
72
4.0
−34
yellowish
green gas
99
3.0
113
2.8
133
2.5
Br2
−7
60
reddish
brown liquid
I2
114
185
metallic black
solid
© 2014 Pearson Education, Inc.
Atomic
ElectroRadius,
negativity
pm
Fluorine, F2
• F2 is a yellowish green toxic gas
• Most reactive nonmetal
– Forms binary compounds with every element except He,
Ne, and Ar
– Including XeF2, XeF6, XeOF4, KrF2
• Produced by the electrolysis of HF
© 2014 Pearson Education, Inc.
Fluorine, F2
• 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.
© 2014 Pearson Education, Inc.
Elemental Fluorine and Hydrofluoric Acid
© 2014 Pearson Education, Inc.
Hydrofluoric Acid, HF
• Produced by the reaction of fluorospar with H2SO4
CaF2(s) + H2SO4(l) → CaSO4(s) + 2 HF(g)
• Crystalline HF is zigzag chains.
• HF is weak acid, Ka = 6.8 × 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
© 2014 Pearson Education, Inc.
Elemental Chlorine, Cl2
• Most abundant of the halogens
• Produced by the electrolysis of seawater
2 NaCl(aq) + 2 H2O(l) → Cl2(g) + NaOH(aq) + H2(g)
– Eighth most abundant substance produced
• Also produced as a byproduct of metal processing
MgCl2(l) → Mg(s) + Cl2(g)
• Used in bleaching, making PVC, insecticides,
freons, manufacture of Br2
KBr(aq) + Cl2(g) → KCl(aq) + Br2(l)
© 2014 Pearson Education, Inc.
Halogen Compounds
• Halogens form ionic compounds with metals and
molecular compounds with nonmetals having
covalent bonds.
• Most halogen oxides are unstable.
– Tend to be explosive
– OF2 only compound with O = +2 oxidation state
– ClO2(g) is strong a 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
© 2014 Pearson Education, Inc.
Interhalogen Compounds
• Halogens can also form compounds with other
halogens, called interhalides.
• For interhalides, the larger halogen atom has lower
electronegativity; Therefore, 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 compounds with B = F can be AB5
– IF7 only known AB7
• Only ClF3 used industrially
– To produce UF6 in nuclear fuel enrichment
© 2014 Pearson Education, Inc.