1. Lecture 19 Aqueous Phase Chemistry Part II
ATS 621 Fall 2012
Lecture 19
Aqueous Phase Chemistry Part II

2. Solving equilibria SO2 SO2(aq) T HSO3- pH SO32- How many unknowns?
How many equations?

3. SOLVING THE SULFUR DIOXIDE / WATER EQUILIBRIUM
STEP 2’
From equilibrium we had:
Add the electroneutrality equation: [H+]=[OH-]+[HSO3-]+2[SO32- ]
If S(IV) is the only species in solution we can solve this for [H+], with one more piece of information (for example PSO2=1ppb, T=298K  pH=5.4, could then calc [S(IV)])
Sum(anions)=sum(cations)
If PSO2 = 1ppt, then pH=6.7  close to pure water because so little SO2  in the future save this for a class Q (ie. if the concentration of SO2 in the atmosphere goes down, do we expect the rainwater to be more or less acidic?)
To calc [S(IV)] either sum up individual or use H* expression
S(IV) increases with increasing pH, decreasing T or increasing PSO2
If other species are present need to modify electroneutrality equation, for example with sulfate:

5. OTHER ACID/BASE EQUILIBRIA…
STEP 2’
CO2 dissolving in a drop (same as in ocean):
CO2(g)
CO2.H2O
HCO2 = 3x10-2 M atm-1
OCEAN
electroneutrality:
[H+]= [HCO3-]+2[CO32-]+[OH-]
Kc1 = 9x10-7 M
CO2.H2O
HCO3- + H+
Can express in terms of K’s and [H+]
Find at 283K, PCO2=350ppm, pH=5.6
(rain slightly acidic)
Kc2 = 7x10-10 M
HCO3-
CO32- + H+
Ammonia (basic in solution):
NH3 (g) + H2O(l) ↔NH4OH(aq)
NH4OH(aq) ↔NH4++OH-
electroneutrality: [H+]+[NH4+]=[OH-]=kW[H+]-1
CO2 hydrolyzes, H* > H, so total amount of CO2 dissolved always exceeds that predicted by Henry’s Law alone for CO2 (for pH < 5, H*~H)  CO2 is not as strong an acid as SO2, so range of H* not as large (cannot pull in as much via ionization)
Can replace [OH-] with Kw/[H+]
Ammonia is basic in solution, probably most important neutralizing agent in the atm
(can do full ammonia derivation following S&P section  find for pH < 8, [NH3T(aq)]=[NH4+]
Salt definition: neutral compound formed by a union of an acid and a base
Salts (dissolution):
(NH4)2SO4=2NH4++SO42-
electroneutrality: [H+]+[NH4+]=2[SO42-]+[OH-]
What is the pH? If assume no exchange with the gas phase, then NH4+ equilibrates with NH3(aq). Then, [NH4+]< 2[SO42-], so [H+]>[OH-] and pH < 7

6. REMINDER: AQUEOUS PHASE REACTION MECHANISM
STEP 2’: Ionization (for some species), VERY fast
STEP 1: Diffusion to the surface
STEP 2: Dissolution X X X
 A+ + B-
STEP 4: Chemical Reaction
STEP 3: Diffusion in aqueous phase
X+Y ? X

7. AQUEOUS PHASE CHEMICAL KINETICS
STEP 4
Dissolved gases and ions in particles can undergo chemical rxn  oxidizing
Units for aqueous phase rate (Ra): M/s (moles/Lsolution/s), if multiplied by L (vol/vol) then the rate is converted to moles/Lair/s  equivalent to a gas-phase conversion rate
- occasionally see fractional rates (% per hour)
Overall kinetics are described by combining Henry’s Law dissolution and ionization with aqueous phase chemical reactions in liquid atmospheric particles.
Simplest case: if drop is in equilibrium with the gas phase (as we will assume)
Important aqueous phase reaction: conversion of S(IV) to S(VI)
- possible reagents: O2 (aq), H298=1.3x10-3
O3(aq), H298=9.4x10-3
H2O2(aq),H298=7.1x104
Important reaction in atmosphere: O3 and H2O2

8. AQUEOUS PHASE REACTIONS: S(IV) TO S(VI), OXIDATION BY O3
STEP 4
SO32-(aq) + O3(aq)  SO42-(aq)+O2(aq)
Rate of sulfate production:
Liquid phase concentration of O3:
Using previous expression for sulfite:
Original production expression did not depend on pH, but see in final form the strong pH dependence (due to equilibrium b/w gas phase SO2 and liquid phase SO32-).
As SO32- is oxidized, it is replaced by SO2 dissolving from the gas phase. Since this is an acidic gas, both the pH and the rate of oxidation decrease as the reaction proceeds  reaction is self-limiting
All 3 S(IV) species react with ozone, the empirical rate equation is:
k2 > k1 > k0

9. AQUEOUS PHASE REACTIONS: S(IV) TO S(VI), OXIDATION BY H2O2
STEP 4
HSO3-(aq) + H2O2(aq) ↔ SO2OOH-(aq)
SO2OOH-(aq) + H+(aq)  H2SO4(aq)
Hoffman and Calvert (1985) suggest the following rate expression:
First reaction accelerated by partitioning of more S(IV) to bisulfite at higher pH.
Second reaction becomes faster as pH declines. Thus, little net effect of pH in atm.
Organic peroxides have also been proposed as potential aqueous S(IV) oxidants, however concentrations and solubilities are low enough that they are predicted to be of minor importance.
Comparison of S(IV) oxidation paths
(Seinfeld and Pandis, 2006)
Rates of sulfate production are per unit volume, and thus must be multiplied by total liquid water content (therefore more important in clouds)

10. AQUEOUS SULFUR OXIDATION
SO2
S(IV)
SO2.HO2
HSO3-
SO32-
H2O2(g)
O3(g)
S(VI)
H2SO4
HSO4-
SO42-

11. MASS TRANSFER LIMITATIONS
STEP 1/3
We have assumed that drops are in equilibrium with the gas phase, or that
this: is fast compared to this: X
X+Y ? X
This is often true because particles are small. BUT gas phase diffusion tends to limit the rates of gas-particle reactions for very large particles, such as cloud droplets.
Mass transfer limitations tend to be most important for relatively insoluble gases (for SO2 oxidation in droplets, O3 is the species most subject to mass transfer limitations)
Diffusion in liquids is relatively slow. Dissolved gases are unlikely to reach the center of the drop for fast reactions. Thus, the average liquid phase concentration is less than expected on the basis of Henry’s Law and overall rate is reduced.
EXTREME example: N2O5 hydrolysis happens so quickly that the reaction appears to happen as soon as molecule enters solution, thus rate ~Adrop
(as if occurring at the surface)
X+Y ? X

12. EXAMPLE FROM SEINFELD AND PANDIS
Pure water droplet (pH=7), L=10-6 m3water/m3air is exposed to environment :
Chemical Species
Initial
(gas phase)
After equilibrium (gas phase)
HNO3
1 ppb
10-8 ppb
H2O2
0.465 ppb O3
5 ppb
NH3
1.87 ppb
SO2
3.03 ppb
After equilibrium conditions solved with equations for Henry’s Law, dissociation equilibrium, electroneutrality and mass balance, such as:
[H+]+[NH4+]=[OH-]+[HO2-]+[HSO3-]+2[SO32-]+[NO3-]
New gas phase concentrations are effective initial conditions. pH=6.17
As reaction proceeds, will convert S(IV)S(VI), will change pH and thus effective Henry’s Law constant (for example for S(IV) dissolution). So at each time step must re-calculate electroneutrality, and now include sulfate formation as well:
[H+]+[NH4+]=[OH-]+[HO2-]+[HSO3-]+2[SO32-]+[NO3-]+[HSO4-]+2[SO42-]
Integrate chemical rate equations in time, assuming instantaneous equilibrium and either depleting (closed) or holding constant (open) the gas phase concentrations.

13. CLOSED VS. OPEN SYSTEM
In an open system, partial pressures remain constant (gases are replenished)
In a closed system, as gas is transferred to the liquid phase it is depleted
[Seinfeld & Pandis]

15. The image shows samples of cloud / fog water collected by Prof
The image shows samples of cloud / fog water collected by Prof. Collett’s group (CSU ATS) during a field study on a mountaintop in China (Taishan). The water was found to contain large concentrations of organic species, most likely due at least in part to emissions from agricultural burning.

16. Atmospheric deposition
Pollutants affect the chemical composition of precipitation and lead to environmental impacts
Nitrogen deposition
Mercury deposition
Deposition of organochlorine compounds (persistent organic pollutants, POPs)
Acidic deposition:
Acid rain
Acid precipitation (e.g., snow)
Dry deposition of acidic gases
Acid fog
Acid clouds
Occult deposition

18. Detrimental effects of acidity http://www. scar. utoronto
Acidification of streams and lakes and effects on aquatic ecosystems
Effects on soil
Mineral leaching; killing of microorganisms
Crop and tree damage
first noticed in the 1970’s in Germany’s Black Forest, in coniferous and deciduous trees of all ages - “Waldsterben”
Damage to buildings and artwork
Limestone, marble, and sandstone dissolved by strongly acidic solutions

20. Acidic deposition http://www. internat. naturvardsverket. se/index
Acidic deposition Also Seinfeld and Pandis (1998), Chapter 20.6
H+ increases in precipitation known for over a century
Acidic rain linked to coal burning in England in 1852
Major scientific attention:
Scandinavian countries in the middle 1960s
1961: Swedish chemist establishes Scandinavian network to monitor surface water chemistry
Showed acid rain was a large-scale regional problem
North America appreciates scope in mid and late 1970s
US and Canadian governments officially recognized effects in 1986
National Acid Precipitation Program (NAPAP),

21. Measure acidity as concentration of H+
pH = - log [H+]
pH = log [1/H+]
Expression of molar
concentration
M = pH 4
M = pH 5
pH ranges from
0 to 14
7 = neutral
< 7 = acidic
> 7 = alkaline

22. What environmental pH is “acidic”?
Carbon dioxide is ubiquitous in the atmosphere, and is slightly soluble in water
It is an ACID in solution:
For background CO2 levels, compute a water pH of ~ 5.6
Theoretical clean atmosphere? There are natural emissions of acidic species as well; should take these effects into account
Remote locations observed / predicted pH = 5.0 or even lower
Areas affected by human activities in Eastern U.S.: observations
Lowest values ~3.0
Average values ~
Usual demarcation for precipitation:
pH > 5.6 uninfluenced
5 < pH < 5.6 may be influenced
pH < 5 ACID PRECIPITATION

23. For moderate pH inputs, can actually have fertilization effects
pH < 3.8: all effects negative

24. Rain acidity

25. What acids are important?
We now know from years of data that the primary causes of acidity are sulfuric acid and nitric acid
Northeastern U.S.: Sulfuric acid (60%); nitric acid (35%)
Western states: Sulfuric acid (10%); Nitric acid (90%)
In North America and Europe, sulfur emissions more than doubled 1900 to 1980
Nitrogen emission are harder to quantify, but may be as high as 20x the 1900 level
In some locations, there are contributions from HCl and from organic acids (e.g., Okefenokee swamp in SE U.S.)
In Katharine, Australia, formic and acetic acids accounted for more than half the rainwater acidity (Keene et al., 1983)
S in fuels  sulfur dioxide, SO2(g) + OH 
sulfuric acid, H2SO4(p)
nitrogen oxides, NO2(g) + OH 
nitric acid, HNO3(g)
Also NO2(g) + NO3(g) 2 HNO3(g)
Note NO + O3 NO2 + O2
NO + HO2 NO2 + OH