Revisiting Le Chatelier’s Principle

Le Chatelier’s Principle is a neat concept, but too often it’s expressed in a not so wonderful way. Here is a  textbook definition from a popular college textbook, fifth edition:

A change in any of the factors that determine the equilibrium conditions of a system will cause the system in such a manner as to reduce or counteract the effect of the change.

With such a definition, students often imagine the system as almost having a consciousness or a purpose. Wikipedia’s definition also leaves something to be desired:

If a chemical system at equilibrium experiences a change in concentration, temperature or total pressure, the equilibrium will shift in order to minimize that change.

They go on to use an example, of course giving the correct prediction of the disturbance, but they also mystify what is actually going on.

2 NO2(g)  ⇌ N2O4(g)

If one increases the pressure of the reactants (2 NO2(g) ) the reaction will tend to move towards the products to decrease the pressure of the reaction.

Maybe it’s not just the paucity of pretty women or (handsome guys) in freshman chemistry classes that keeps the classes so small!

An equilibrium is established in a closed system when the forward reaction rate equals that of the reverse reaction. To predict how a system in equilibrium will react to a disturbance, we could view the forward and reverse reactions as if they were in competition with one another. The prevailing reaction will be the one that benefits more from the disturbance. If one reaction is being more hampered than the other, obviously the one facing the obstacle will not win out.

Returning to the nitrogen dioxide reaction,  2 NO2(g)  ⇌ N2O4(g) , if we increase the pressure of the system by decreasing the volume, the concentration of every reactant will increase. But the forward reaction in this case will benefit more from the increased frequency of molecular collisions. Why? If compression increases the concentration of each reactant  by a factor of x then, since the forward rate can be expressed as k[NO22 , the rate will increase by a factor of x2. But the reverse rate, which equals k[N2O4(g)] will only increase by a factor of x. So increasing the pressure creates more N2O4 not because of a compensation mechanism but because compression increases the forward rate more than the reverse rate.

In the equilibrium Cr2O72-(aq) + H2O(l) ⇌ 2 H+(aq) + 2 CrO42-(aq)
the orange dichromate(Cr2O72-(aq) ) reacts with water to create acid(H+(aq) )and yellow chromate(CrO42-(aq)). Meanwhile at the same rate, acid and chromate combine to regenerate dichromate and water. So to the naked eye, it looks like nothing is going on because the work of the forward reaction is being offset by that of the reverse reaction. We could have the stable orange color of dichromate dominating.

Although the ions are all mixed in a soup and don’t have respective sides, it’s a bit like if I dump snow on my neighbor’s driveway, and he dumps snow from his driveway onto mine at the same rate. A plane’s passengers flying over us will simply see an unchanging white driveway.

But if we disturb the equilibrium by adding some base, it eliminates some H+(aq) .  Reducing the concentration of H+(aq)thus lowers the reverse rate. But since the forward reaction is unaffected by the presence of base, it proceeds at its regular rate. So what we see is the yellowish color of chromate dominating. There’s no system trying to “reorder things” or “offset anything”. It’s simply the case of a hampered reverse reaction that does not keep up and maintain the original equilibrium. If I suddenly had trouble picking up the snow while my neighbor proceeded as normal, the snow would suddenly pile up on my side. The plane passengers would see the asphalt on the neighbor’s driveway.

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Made In China: PM2.5 Pollution From Ammonia

Fill a glass with water and place it under a running tap. Water will flow out of the glass at the same rate that it flows into it. It’s not the same molecules that are in the glass at any given moment, but we maintain the same volume of water. We have what’s known as a steady state.

Nitrogen is the most common gas in the atmosphere. In terms of molecules or volume–they are proportional to each other under the same conditions of pressure and temperature–about 78% of air is N2. Air’s other main component, diatomic oxygen, has a bond dissociation energy of 494 kJ per mole. But to break nitrogen’s triple bond takes almost twice as much energy: 942 kJ per mole. This makes nitrogen unreactive within the atmosphere’s normal temperature range.

This is a problem for plants. The nitrogen atom is not only needed for proteins, which include enzymes that speed up most of life’s reactions. There are nitrogen atoms in the nucleotide bases of DNA and RNA. And nitrogen is also found in ATP. It’s the molecule that is continuously produced either through photosynthesis or with the energy-releasing breakdown of glucose and which facilitates all sorts of reactions in life.

So to survive, plants rely on bacteria which decompose animal and plant waste into reactive nitrogen-containing ions such as nitrates(NO3) and ammonium(NH4+). Some plants have set up sophisticated partnerships with certain bacteria species that have evolved ways of directly converting atmospheric nitrogen into NH4+. Other bacteria have the ability to return molecular nitrogen to the atmosphere by first converting ions that are unabsorbed by plants into nitrites. It’s this last step that normally keeps the nitrogen in the air at steady state.

But we have further complicated the nitrogen cycle. With seven billion people on the planet, 1.3 billion in China alone, we have chosen to accelerate the growth of edible plants partly by manufacturing ammonia from nitrogen and hydrogen using high pressures and temperatures and catalysts. Along with other forms of reactive nitrogen, these compounds are added to the soil. Whereas for decades so-called runoff of excess nitrogen(and phosphorus) fertilizer has drawn a lot of attention, a more subtle consequence of ammonia enrichment has been less noticed.

China uses 18.7 million tons of nitrogen fertilizer per year. There have been suggestions that because of government subsidies they overuse it, but presently they use less than the U.S. on a per capita basis. When a densely populated country like China has to focus its agriculture in only one area within its boundaries, the concentration of ammonia in the air above fertile lands increases. Ammonia then goes on to react with acids formed from industrial and vehicle emissions of sulfur and nitrogen oxides. The ammonium compounds, one being ammonium sulfate((NH4)2SO4), form very fine crystalline particles with diameters of 2.5 microns (2.5 X10-6meters) or less. Particulate matter of these dimensions is referred to as PM2.5.

If you’ve ever had to use a fire extinguisher in your home, you may be aware that most of them contain monoammonium phosphate(NH4H2PO4) and or ammonium sulfate. Cleaning up afterwards is not exactly one of life’s pleasures. The dust is extremely fine, difficult to pick up and very irritating if inhaled. But those particles are probably still not as small as those that form in the atmosphere; those are 1/40th as wide as the average human hair. Their small size gives them a longer residence time in the air due to Brownian motion, and their size allows them to lodge deeply into the lung’s air sacs. They aggravate most respiratory diseases and lead to premature deaths. The conclusion from a joint study involving Health Canada, New York State University School of Medicine, the American Cancer Society was that:

Each 10-µg/m3 elevation in fine particulate air pollution was associated with approximately a 4%, 6%, and 8% increased risk of all-cause, cardiopulmonary, and lung cancer mortality, respectively.

Other studies have specifically shown that PM2.5 is far more worrisome than PM10 or sulfate or H+.

For two days in January 2012, the city of Beijing was covered with thick smog, forcing the cancellation of almost 700 flights at airports. In the same month, the city finally decided to report analyses of PM2.5, whereas previously all of China only reported PM10. This in itself is not going to ameliorate the emissions problem in China, but it’s one small step in the right direction.

Sources:

Colorado State University Ammonia Best Management Practices

Current and Future Emissions of Ammonia in China

http://www.sciencedirect.com/science/article/pii/S1352231001003016

http://www.tandfonline.com/doi/abs/10.1080/02786820119445

http://ukpmc.ac.uk/abstract/MED/8875828/reload=0;jsessionid=4xMaEY2mISBUYJD2TRi1.110

http://www.indexmundi.com/en/commodities/minerals/nitrogen/nitrogen_t5.html