Environmental Chemistry
Chapter 3:
The Detailed Chemistry of the Atmosphere
Copyright © 2007 DBS
Review: How to Draw Lewis Structures
1.
Determine the sum of valence electrons
2.
Use a pair of electrons to form a bond between each pair of
bonded atoms
3.
Arrange the remaining electrons to satisfy octet rule (duet rule
for H)
4.
Assign formal charges (valence – directly surrounding e-)
Methane, CH4
CH4 and H2O
Water, H2O
but unlike methane, two e- pairs are
bonding and two are non-bonding
The non-bonding e- pairs
take up more space than
bonding pairs, so the Hto-O-to-H bond angle is
compressed
VSEPR
No. e- pairs around
central atom
Shape of molecule
Bond angle
4 pairs, all bonding:
CH4, CF4, CF3Cl, CF2Cl2
Tetrahedral
109.5°
4 pairs, three bonding,
one non-bonding:
NH3, PCl3
Triangular pyramid
~107°
4 e- pairs, two bonding,
two non-bonding:
H2O, H2S
Bent
~105°
Valence Shell Electron Pair Repulsion Theory assumes that the most stable
molecular shape has the electron pairs surrounding a central atom as far away
from one another as possible
Lewis Structures of Free Radicals
•
•
•
Free radicals possess an unpaired eThe unpaired e- is not in actual use as a bonding eCarbon centered radical in which the carbon atom has one unpaired e- forms 3
bonds rather than four
•
H―C―H
|
H
•
Oxygen forms one rather than 2 bonds:
•O – H
•
A halogen forms no bonds:
Cl•
Lewis Structures of Free Radicals
• Choice of assigning the unpaired e• Hydroperoxy radical, HO2:
H-O-O•
• Complicated in molecules containing multiple bonds
• For hydroxy formyl, HOCO a reasonable structure is:
•
H-O-C=O
Does not go to O
since C must have 4
valence
Lewis Structures of Free Radicals
A simple formula, ClO• does not indicate which atom carries the e-.
Draw Lewis structures for:
OH•
CF2Cl•
ClO•
NO•
•O – H
•
F – C –F
|
Cl
Cl – O•
•N = O
Hydroxyl Radical: The Atmosphere’s Detergent
•OH is the prominent oxidizing species in the atmosphere
• Despite very low atmospheric concentrations, currently
estimated at 106 molecules cm-3, corresponding to a mean
tropospheric volume mixing ratio of 4 x 10-8 ppmv
• The lifetimes of most atmospheric gases are, therefore, largely
determined by [OH] and the corresponding reaction coefficients
• Radical reactions that are spontaneous produce stable products
with strong bonds
Hydroxyl Radical: The Atmosphere’s Detergent
The major route for the formation of the hydroxyl radical in the
troposphere is:
NO2• + h ( < 400 nm) → NO• + O•
O• + O2 + M → O3
O3 + h ( < 320 nm) → O2 + O*
O* + H2O → 2 •OH
NO2• + H2O → NO• + 2 •OH
Others:
O* + CH4 → OH + CH3OH
HNO2 → OH + NO
H2O2 + h → 2OH
Interactions with Hydroxyl Radical
•
•
•
Usually it reacts by adding itself to a molecule at the multiple bond
It can also abstract hydrogen atom to produce carbon centered radicals
•OH addition does not occur to O=O bonds since the bonding that
would result will be weak
For example, in the case of SO2, the OH radical adds to the sulfur atom
forming a strong bond but not to an oxygen atom
•
•
•
Hydroxyl radicals do not add to CO2 since C=O bonds are very strong
However, it adds to CO, the addition favors conversion of triple bond to
stable double bond
Radicals React with O2 to produce Peroxy
and Hydroperoxy Radicals
•
Predominant fate is ‘add-on’ reaction with O2,
e.g. •CH3 + O2 → CH3OO•
H3C – O – O•
Successive reactions will
completely oxidize the
organic compound
HOO• / HO2• (hydroperoxy) and CH3OO• are called peroxy radicals
- Less reactive than other radicals - Do not readily abstract H
- Do not react with O due to low conc.
•
Main reactions:
HOO• + NO• → •OH + NO2•
R-OO• + NO• → RO• + NO2•
(where R = carbon chain)
H Atom Abstraction by O2 from Nonperoxy Radicals
CH3-O• + O2 → H2C=O + HOO•
•
H-C=O + O2 → C=O + HOO•
Gases that undergo decomposition by
absorbing UV-A or visible light can
generate free radicals.
e.g., formaldehyde
H-abstraction occurs
provided a new bond
is formed
H2CO + UV-A (<338 nm) → H• + HCO•
If there is no suitable hydrogen atom for O2 to abstract then it adds-on
peroxy radical
HNO3, HCl, NH3, etc
Fate
H2CO
Decision tree
illustrating
the fate of gases
emitted
into the air
CH4 + OH•  H2O + CH3•
ROO· + NO  NO2 + RO·
Fate of Free Radicals
Decision tree
illustrating the
fate of airborne
free radicals
CH3· + O2  CH3OO·
Oxidation of CH4
•
•
•
•
•
Produced in inefficient (anaerobic) burning of fuels
Predominant HC in atmosphere
No multiple bonds
Not soluble in water, does not absorb sunlight
Slow oxidation initiated by hydroxyl radical
(hydrogen abstraction reaction)
CH4 + •OH → •CH3 + H2O
•CH3 + O2 → •CH3OO•
CH3OO• + NO → CH3O• + NO2
CH3O• + O2 → H2CO + HOO•
abstraction
O2 adds forming peroxy
transfer of O
O2 absracts H
…conversion of methane to formaldehyde
H2CO + UV-A (338 nm) → H• + HCO•
H• + O2 → HOO•
HCO• + O2 → CO +HOO•
unstable
O2 abstracts
O2 abstracts
Note: CO is a stable intermediate and can further
undergo transformations
C
O + OH• → HO-C=O
H-O-C=O + O2 → O=C=O + HOO•
….. Production of CO2 as the final product
CH4 + 5O2+ NO + 2OH• + UV-A →
CO2 + H2O + NO2 + 4HOO•
Notice the radicals consumed
and produced.
What happens to the HO2
produced?
What happens to the NO2
produced?
(see fate of free radicals)
Reaction
intermediates
during hydride
oxidation
Problems
• 3-4
• 3-5
• 3-6
Part 2
Photochemical Smog
Oxidation of Reactive Hydrocarbons
Saturated hydrocarbons such as CH4
react with hydroxyl radical by hydrogen
abstraction
Hydrocarbons with double bond (e.g.,
ethene) react with •OH by addition
because of lower activation energy
Energetics favor
addition over
abstraction
…formation of carbon centered radical
Photochemical
Carbon centered
radical reacts with O2Smog
to produce a peroxy
radical which in turn oxidizes NO to NO2
Photochemical decomposition of
NO2 to NO and O and formation of
ozone results in photochemical smog
NO2 → NO + O
(1)
O + O2 → O3
(2)
NO + O3 → 2NO2 + O2
(3)
NO2 is the only
significant source of O
Formation of Aldehydes
Decomposition of carbon centered radical
Original C=C
is split into 2
aldehydes
Aldehydes further decompose in sunlight
RHCO + sunlight → R• + HCO•
….further increase in the number of radicals
Overall
RHC=CHR + OH• + 2O2 + NO• → 2RHC=O + HOO• + NO2•
Mechanism of the RHC=CHR oxidation process in the smog
Energetics favor addition over abstraction
Addition of O2 to radical center
NO is oxidized by the C-O-O·
Cleavage allows formation of aldehyde double bond
Reaction with O2 allows 2nd aldehyde to form
Photolysis follows.
Peroxy radicals are formed.
NO is oxidized to NO2
Radicals formed:
HO2 (2); OH (1), RO (1), NO2 (3)
Also CO
Problems
• 3-8
The Fate of Free Radicals
Rate of reaction between two radicals increase as the radical
concentration increases
R• + R’• → R-R’ stable molecule
e.g.,
OH• +NO2• → HNO3
sunlight
OH• + NO• → HNO2
OH• + NO•
(HONO accumulates only in the night)
When the concentration of NOx is low,
2OH• → H2O2
2HOO• → H2O2 + O2
Starts AM cycle
Fate of Other Radicals
•
O2 + R-C=O
Peroxyacetylnitrate is eye irritant
and toxic to plants
(peroxyacetylnitrate)
Thus in the afternoon hours a build up of oxidizing agents such as
nitric acid, hydrogen peroxide and PAN is encountered
Hourly Variation of Concentration of Gases
HC → Aldehydes
•
NO → NO2
•
•
•
Source: http://jan.ucc.nau.edu/~doetqp-p
Early morning traffic increases the
emissions of both nitrogen oxides
and VOCs as people drive to work
Later in the morning, traffic dies
down and the nitrogen oxides and
volatile organic compounds begin
to react forming nitrogen dioxide,
increasing its concentration
As the sunlight becomes more
intense later in the day, nitrogen
dioxide is broken down and its
byproducts form increasing
concentrations of ozone
As the sun goes down, the
production of ozone is halted. The
ozone that remains in the
atmosphere is then consumed by
several different reactions
Role of NO3•
• Nitrate radical produced from NO2 and O3
NO2• + O3 → NO3• + O2
• Photolysis yields NO2 and O
• Abstracts H from RH during evening
NO3• + RH → HNO3 + R•
Similar to OH
Part 3
Oxidation of SO2 (g)
Addition of OH• followed by the formation of SO3
SO3 + H2O(g) → H2SO4(g)
H2SO4(g) + nH2O → H2SO4 (aq)
Oxidation of SO2 (aq)
Determination of total sulfur content in water
SO2 is soluble in water. It exists in the dissolved form if there is significant cloud or mist in the
atmosphere. The oxidation to sulfuric acid occurs in the aqueous phase after rain drops
accumulate on earth.
SO2 (g) + H2O (aq) ⇌ H2SO3 (aq)
Typically SO2 conc. is 0.1 ppm or (0.1/106) =1 x 10-7 atm
From Henry’s law, KH = 1 M atm-1 = [H2SO3]/P
 [H2SO3] = 1 M atm-1 x 1x10-7 atm = 1 x 10-7 M (or moles/L)
But H2SO3 dissociates readily with a dissociation constant of K = 1.7 x 10-2 M-1
H2SO3 ⇌ H+ + HSO3As HSO3 dissociates, more of SO2 dissolves until it reaches an equilibrium with H+ and HSO3
1.7 x 10-2 M-1 (or K) = [H+][HSO3-]/[H2SO3]
1.7 x 10-2 M-1 (or K) = [HSO3-]2/[H2SO3] = [HSO3-]2 / 1 x 10-7 M
…[H+] = [HSO3-]
[HSO3-]2 = 17 x 10-10 M2 = 4 x 10-5 M
 Total dissolved S is 4 x 10-5 M
Oxidation of SO2 (aq)
•
•
•
•
Dissolved SO2 is oxidized by trace
amounts of H2O2 and O3
Sunlight is a dominant factor in forming
O3 and H2O2
If strong acids are present in the droplet
they control the pH.
Any freshly dissolved SO2 has no effect
[HSO3-] = K x [H2SO3]/[H+]
=1.7 x 10-2 x 10-7/[H+]
=1.7x10-9/[H+]
…inversely proportional to H+
Since strong acids dissociate readily,
[H+] concentration controls the overall
concentration of HSO3-
•
•
Acidity of the droplet has effect on the rate of SO2 oxidation
At pH below 5 H2O2 dominates oxidation and above pH 5 ozone or other
catalytic reactions dominate the oxidation
•
•
Hydrogen abstraction reactions dominate chemistry in both stratosphere
and troposphere
…….but the radicals that dictate the chemistry are different
Stratosphere: •OH, •O, •Cl, and •Br abstract H atom from stable
molecule such as CH4
Troposphere: hydroxyl and NOx radicals are the primary reactants
Processes Involving Loosely Bound
Oxygen Atoms
A Y-O structure from which O atom can be detached readily
Examples of “Loose O Atom” Reactions
Reaction with atomic oxygen
Y―O → Y + O2
Photochemical decomposition
Y―O + sunlight → Y + O
Reaction with NO
Y―O + NO → Y + NO2
Abstraction of oxygen from Ozone
Y―O + X → Y + XO
O2―O + Cl → O2 + ClO