Transcript Document

Oxygen on Earth
H2O (oceans)
O2, CO2 (atmosphere)
CO3 (rocks, coral, seashells)
SiO2, silicates (sand, clay, rocks)
Oxygen Content
Air
Earth
Crust
Made commercially by fractional distillation of air (b.p. = 90K)
ALLOTROPES OF OXYGEN
O2
Paramagnetic (why?)
O3
Higher energy form - important UV absorber
in the stratosphere
Light or
electrical discharge
3O2
2 O3
decomposition
Ozone (O3) is a strong oxidizing agent, highly toxic
Kills bacteria (replacement for Cl2 in municipal water treatment)
Irritating component of photochemical smog
OXYGEN IONS
Oxide Ion  O2
(most compounds)
e.g. Li2O = 2Li+ O2
Peroxide Ion  O22 = O – O
e.g. Na2O2 = 2 Na+ O – O 
Also, H2O2 (hydrogen peroxide)
Superoxide Ion  O2
e.g. KO2 = K+ O2
Can have positive oxidation states in combination
with fluorine
+ 2 in OF2
HYDROGEN PEROXIDE
• Strong oxidizing agent
(30-85% solutions)
e.g. bleaching wood pulp to produce white paper
• Hair bleach (~6% solution)
• Antiseptic (3% solution)
H2O2 decomposition can be explosive:
2 H2O2  2 H2O + O2
H = 200 kJ/mol
(disproportionation reaction)
HYDROGEN PEROXIDE
Reduction to H2O:
H2O2 + 2H+ + 2I  I2 + 2H2O
Oxidation to O2:
2MnO4 + 5 H2O2 + 6H+  2Mn2+ + 5O2 + 8H2O
SULFUR
Sources:
Sulfide Minerals (S2-):
FeS2 (Pyrite) Cu3FeS3 (Bornite) PbS (Galena) ZnS (Zinc Blende) -
Iron Ore (Fool’s Gold)
Source of Cu
Source of Pb
Source of Zn
Sulfate Minerals (SO42-)
e.g. Na2SO4, MgSO4
Also, CaSO4 · (H2O)2 (Gypsum)
Used for wallboard, plaster of Paris.
COMMERCIAL SOURCES OF SULFUR
1) Sulfur Mines – Along Gulf of Mexico, deposits of S8 
Frasch Process.
2) Byproduct from other manufacturing processes.
a) Production of Zn, Pb, and Cu from their sulfide ores.
b) Petroleum – 3% S.
c) Coal – 5%.
SO2 forms when coal is burned.
SO2 + H2O  H2SO3
SO2 +[O] SO3 +H2O  H2SO4
CaCO3 + H2SO4  CaSO4 + H2O + CO2
(Marble)
Acid Rain
Sulfur
Two allotropes
S8
yellow, cyclic
Polymer Sx red-brown
polymer: zigzag chains of sulfur atoms
S8(s)  S8(l) 
Melts at 113C
S
S
S
S
S
S
S
T> 150C
Sulfur
Common Oxidation States
+6
SO3; H2SO4 (sulfuric acid)
can’t be oxidized
can only be reduced
+4
SO2; H2SO3 (sulfurous acid)
can be both oxidized AND reduced
-2
H2S; S2can’t be reduced
can only be oxidized
Compounds of S
• H2SO4
most important industrial chemical
• H2S (rotten egg smell) (S2- )
source: metal sulfides + strong acid
e.g. ZnS + HCl  ZnCl2 + H2S(g)
–
–
poisonous
tarnishes Ag in presence of O2
4Ag + 2H2S + O2  2Ag2S + 2H2O
– Organic sulfides, e.g. C4H9SH
Strong odor – added to natural gas
• S2O32- (thiosulfate)
used in photography: forms water soluble complexes
with Ag
Uses of H2SO4
• Making phosphate fertilizer
Ca3(PO4)2 + 3H2SO4 
3CaSO4 + 2H3PO4
~65% of H2SO4
• Manufacture of chemicals
• Metal refining
• Petroleum refining (as catalyst)
• Strong oxidizing agent
• Drying agent
Selenium, Tellurium
Source: metal sulfides
byproducts of Cu, Pb refining
Uses: semiconductors
e.g. Se: Has low electrical conductivity in the dark which
increases in light - photoconductor
Used in photocopiers, light meters in cameras
Compounds: form covalent bonds
Oxides and hydroxides are acidic (typical of nonmetals)
Se
Te
Po
non metal
semi-metal
metal
Nitrogen
Nitrogen (N2) is very unreactive
triple bond energy = 941kJ/mol
Source
fractional distillation of air
(78% of air is N2)
KNO3
water soluble salts
NaNO3
found in deserts
Nitrogen fixation: formation of N containing
compounds from N2
N is an essenial element in proteins, nucleic acids
& necessary to maintain soil fertility
Compounds of Nitrogen
Oxidation states of 3 to +5
Compounds with H
1. NH3
(3 oxidation state)
2. N2H4 (2 oxidation state)
strong reducing agent:
N2H4  N2(g) + 2 H2(g) H = 
forms N2 readily:
S = +
3. Dimethyl hydrazine (rocket fuel)
N compounds with oxygen
N2O
colorless, odorless gas
used as anesthetic (laughing gas)
propellant in whipping cream
NO
formed in car engines: N2 +O2  2NO
N2O3 blue solid, decomposes: N2O3  NO + NO2
NO2
brown gas; component of smog
N2O4 2NO2  N2O4
N2O5 unstable, decomposes to NO2
HNO3
Produced from NH3 by Ostwald process (catalytic
oxidation).
Uses:
fertilizer
NH3 + HNO3  NH4NO3(s)
strong acid
strong oxidizing agent.
cleaning agent
to make explosives
(e.g. nitroglycerine, TNT)
Hydrolysis of oxides
Hydrolysis: reaction with water
N is a non metal: oxides are acidic.
Oxide + H2O = hydroxide
N2O3 + H2O  2HNO2
(nitrous acid)
3NO2 + H2O  2HNO3 + NO
N2O5 + H2O  2HNO3
(nitric acid)
PHOSPHORUS
Source: Phosphate Minerals
Ca3(PO4)2 contains PO43- (tetrahedral P)
P is made by heating Ca3(PO4)2 and coke in an electric
furnace.
2Ca3(PO4)2(s) + 10C(s) + 6SiO2  6CaSiO3(s) +
10CO(g) + P4(g)
Two allot ropes:
P
Whit e P : P
4, t et rahedral P
P
P
Red P : P olymeric
P
P
P
P
P
P
PHOSPHORUS ALLOTROPES
White phosphorus (P4) burns spontaneously in air.
P4(s) + 5O2(g)  P4O10
H = 3000 kJ/mole
Red phosphorus (polymeric) is more stable. Not volatile.
Does not react with air at 25°C.
600oC
Red P
White P
Let Stand
OXIDES OF PHOSPHORUS
PHOSPHORUS OXYACIDS
P4O10 + 6H2O  4H3PO4 phosphoric acid
P4O6 + 6H2O  4H3PO3 phosphorous acid
Also H3PO2
hypophosphorous acid
USES OF PHOSPHORUS
Fertilizer
P is essential for plant growth
Ca3(PO4)2 + 3H2SO4  2H3PO4 + 3CaSO4
H3PO4 + 3NH3  (NH4)3PO4
Detergent
Complexes metal ions
Biological molecules (DNA, RNA)
Biochemical energy source (ATP)
COMPARISONS IN GROUP V
Nitrogen
N2(g) NN
NH3 is stable.
Non-metal  oxides dissolve to give acidic
solutions
N2O3 + H2O 2HNO2
N2O5 + H2O  2HNO3
3NO2 + H2O  2HNO3 + NO
PHOSPHORUS
Allotropes:
White P  P4, tetrahedral.
Red P  polymer.
PH3 burns in air.
Non-metal  oxides dissolve to give
acidic solutions:
P4O10 + 6H2O  4H3PO4
P4O6 + 6H2O  4H3PO3
ARSENIC
Allotropes:
Yellow As  As4
Gray As  brittle solid.
AsH3 ignites spontaneously in air.
As4O10 – acidic oxide:
As4O10 + 6H2O  4H3AsO4
As4O6 is amphoteric, but is more soluble in base.
ANTIMONY – Sb
Brittle gray metalloid.
Sb4O6 is amphoteric.
There is no Sb4O10.
BISMUTH - Bi
Bismuth is a metal.
Bi4O6 is basic and
dissolves only in acids.
Bi(OH)3 is basic.
Bi5+ is rare.
OXIDATION STATES
P5+ dominates.
As3+, As5+ are equally common.
Sb3+ dominates.
Bi3+ dominates.
Inert Pair Effect
HYDRIDE STABILITY
NH3 is stable.
PH3 is stable but burns in air.
AsH3 decomposes easily.
SbH3, BiH3 are very unstable.
GROUP V TRENDS
Going down the periodic table:
1)
2)
3)
4)
5)
Electronegativity decreases.
Switch from non-metallic to metallic.
Hydroxides and oxides become more basic.
Hydrides become less stable.
“Inert pair effect” becomes more pronounced: +3
becomes more stable as compared to +5.
CARBON and Group IV
Carbon Sources:
1) Elemental form – coal.
2) Carbonate rocks (CO32-)
Limestone, marble, chalk = CaCO3
Dolomite = MgCO3
ALLOTROPIC FORMS OF CARBON
1) Diamond - used as abrasive, in drill
bits and cutting tools, and as a gem.
2) Graphite - used in batteries, pencils,
and lubricants.
3) Fullerenes - More recently discovered
molecules such as C60 which has the
shape of a soccer ball.
Carbon Black – Soot
Amorphous form of carbon used in
tires, inks, pigments, and carbon paper.
CARBIDES
1)
Ionic Carbides
Contain C4- or C22- (-CC-)
C4-: Be2C, Al4C3 react with water to give CH4.
C2 2-(-CC-): CaC2 reacts with water to give HCCH.
2)
Covalent Carbides
Carbon is bound covalently to a metal or metalloid.
SiC - almost as hard as diamond, does not react w/water
3)
Interstitial Carbides
Metals with carbon atoms found in between the metal
atoms in the structure.
Steel – often harder than the pure metal.
SILICON
Second most abundant element.
Found in combination with O.
Silicate Minerals: [Si2O52-]n, SiO44Sand: SiO2 (this is also quartz).
With aluminum in aluminosilicates (clay, feldspars).
Prepared by:
SiO2(s) + 2C(s)  Si(l) + 2CO(g)
sand
coke
98%
(3000C)
Very pure silicon (<1 ppb impurity) is required for electronics
applications.
GROUP IV TRENDS
Going down the periodic table:
1)
The +2 oxidation state becomes more stable than +4 due
to the “inert pair” effect.
+2 is rare for C, Si, Ge.
+2 in some compounds, +4 most common for Sn.
+4 is unstable for Pb  strong oxidizing agent (PbO2)
2)
Basicity of oxides and hydroxides increases.
CO2, SiO2, GeO2 are weakly acidic.
SnO, SnO2, PbO are amphoteric.
3)
Hydrides become less stable.
Enormous number of stable hydrocarbons.
SiH4 is stable but is spontaneously flammable.
Ge, Sn, Pb hydrides are very unstable.
Orbital Hybrids and Valence
2s
Li Be
2p
B C N O F
3s
Na Mg
3p
Al Si P S Cl
The differences between the 2nd and 3rd periods:
2nd period: Only s and p orbitals are possible with n = 2
Therefore, the maximum number of bonds is 4
(single and/or double bonds)
Examples: CH4, NF4+, BH43rd (and higher periods): can use d-orbitals to make bonds
E.g.
PF5 P atom is sp3d
SF6 S atom is sp3d2
Let’s look at valences:
N can gain 3 electrons or lose 5 to make an octet
But, N can only make 4 bonds (maxiumum for n=2)
Therefore N usually has a valence of 3
(NH3, NCl3, CH3NH2 - all have 3 bonds and one
lone pair on the N atom)
N with oxidation state 5 never has more than four bonds:
O
e.g., NO3N=O (4 bonds to N)
O
NO2+
O=N=O (4 bonds to N, like CO2)
Likewise, O usually makes 2 bonds: H2O, OF2, H2C=O
Likewise, C can gain 4 or lose 4 electrons to make an octet
(valence = 4)
So carbon always makes 4 bonds
CH4
(4 single bonds)
O=C=O
(2 double bonds)
H-CC-H
(1 single + 1 triple bond)
H2N
C=O
H2N
(2 single + 1 double bond)
(urea)
What about 3rd (and higher) periods - Si, P, S…?
For these elements, double bonds are very uncommon
(usually only single bonds)
Reason: atoms past second row are too big
C
Si
good sideways
overlap of p orbitals
(double and triple bonds OK)
C
Si
poor overlap of p orbitals
---- no multiple bonds
(can still make single bonds)
So CO2 is molecular (O=C=O, has double bonds)
But SiO2 (quartz, sand, glass…) is a 3-dimensional solid
network:
O2 is molecular (O=O, has a double bond)
But S forms rings (e.g., S8)
Nitrogen (N2) has a triple bond NN (very stable molecule)
But phosphorus is found in several forms (white, red, black),
all of which have only single bonds.
The chemistry of carbon is unique because:
• It has a valence of 4 (highest in 2nd period)
• It can make stable bonds with itself
• It can make multiple bonds to C, N, O
• The C-H bond is nonpolar, but bonds to other
elements (N, O, halogens) are polar
This is why life is based on the chemistry of carbon
(organic chemistry)