Chemist Speaks Out on Geoengineering / Chemtrails


The chemist who authored the article below has summarized exactly the soils changes and plant/tree die off being witnessed around the globe. Countless lab tests of precipitation from all over the planet prove beyond reasonable doubt that ongoing geoengineering/chemtrail programs are raining a mountain of toxic metals down on us all. Air, soils, and waters are all being systematically poisoned. “Alumina/aluminum” is only one of a long list of toxic metals used in stratospheric aerosol geoengineering spray programs. Other metals showing up in all lab tests are barium, strontium manganese, and recently, fluoride is also now apparently in the mix. Rain tests from Northern Europe are testing positive for this metal along with the rest.

We are all literally in a fight for life. The planets “life support” systems are being systematically sterilized by the experimental weather modification programs. If more scientists like the one that penned the article below had the courage to speak out, we could bring the dire crime of geoengineering to light. It is up to all of us to help with the effort of exposing the geoengineering/chemtrail spraying.

Educate yourself on this dire issue. Arm yourself with credible information. Examples of information flyers can be found at Get a copy of Michael Murphy’s “Why In The World Are They Spraying”, make copies from your original, and pass them out. Passing out credible data does far more good than frightful rants which only cause peoples defenses to go up.

Locate groups and organizations that would care about this issue if they only knew. Pass data on to them.

Again, its up to all of us.
Dane Wigington



Submitted by dwalters on Mon, 12/24/2012

Please bear with me as I explain the general consequences of spraying alumina (aluminum oxide) into the atmosphere. The chemistry required to completely understand what I will discuss is not considered to be advanced, and individuals that have completed General Chemistry II in college should have sufficient knowledge to keep up with less difficulty (if they paid attention in class). However, I have taught this information to many students, and I will attempt to be clear enough for an interested layman to learn the necessary material as they read (feel free to ask questions).


Most of you are aware that only a certain amount of a substance can be dissolved in a given amount of liquid. For instance, ten pounds of sugar can’t be completely dissolved in a single cup of water. Eventually as a substance is continuously added to a solution, a maximum concentration is reached. Any amount added beyond that maximum concentration will just fall to the bottom of the container still remaining in its solid form. Changing the temperature will change this maximum amount, but at a given temperature the maximum concentration is a constant.

Imagine for a moment that you have a clear glass filled with pure water and begin adding epsom salt (magnesium sulfate) until its maximum concentration has been reached and slightly surpassed. As a result, you now have a clear glass with a clear solution of epsom salt and a small pile of the salt sitting on the bottom of the glass.

To the naked eye, it appears that nothing is occurring in the glass. In reality, magnesium sulfate is constantly being dissolved and redeposited to and from the surface of the solid but these two are happening at the same rate – so the concentration of espom salt in the solution remains at its constant, maximum concentration. One could say that an equilibrium has been established. Since things areconstantly happening at the surface of the solid salt, this equilibrium is dynamic.

Imagine now that I devise a way to remove some of the dissolved magnesium sulfate from the solution without changing the amount of water present. For instance, I may have a rod of a certain material that adsorbs the epsom salt to its surface. I could dip this rod into the solution and remove some of the dissolved salt. If I decide to do this, what happens to the concentration of the espom salt in the solution in the few moments after submerging the adsorbing rod? For the moment, the concentration of magnesium sulfate will be decreased.

In this scenario, the equilibrium has been disturbed. The magnesium sulfate concentration is now below its possible maximum concentration (given that the temperature has remained constant). However, this won’t last for long because you have added an excess of epsom salt in the first place. Some of the solid begins to dissolve as the salt is removed by the rod, and after a little time, the maximum concentration is once again reached leaving a little less solid on the bottom of the glass than was previously there. It turns out that all chemical equilibrium processes do this. This principle is so handy that it has a common name –

Le Chatelier’s principle – If a chemical system at equilibrium experiences a change in concentration, then the equilibrium shifts to counteract the change and a new equilibrium is established.

There are a couple of more concepts that we should discuss before moving on to the use of aluminum oxide as means for reflecting sunlight by spraying it into the atmosphere – the concepts of ions and how mathematics is applied to their participation in equilibrium processes (only knowledge of multiplication and division is necessary to understand).

Ions are charged entities that result from the dissociation (coming apart) of ionic compounds. A convenient example of an ionic compound is magnesium sulfate (our epsom salt). I cannot use the necessary html to make superscripts and subscripts. So, instead we will assume that a number appearing by itself would typically be a subscript, and numbers appearing in curly brackets – {} – are intended to be superscripts. The symbolic representation of magnesium sulfate can then be written as – MgSO4 – and the ions that result from its dissociation can be represented as – Mg{2+} & SO4{2-}. The numbers appearing in the curly brackets are the units of charge on each ion, where the magnesium has a charge of “two plus” and thepolyatomic (being composed of more than one atom) sulfate ion has a charge of “two minus.”

Recall before that we used an excess of MgSO4 when performing the above thought experiment to explain Le Chatelier’s principle. Also recall that the equilibrium was dynamic; there was constant dissolving and redeposition occurring at the surface of the salt that rested at the bottom of the glass. How, then, can this chemical equilibrium be represented using our symbolic representation? Well…

MgSO4(solid) <—> Mg{2+}(aqueous) + SO4{2-}(aqueous)

where the two-sided arrow (<—>) means that the reaction goes both ways, and the word aqueous simply means that the magnesium and sulfate ions are dissolved in a water solution. Typically, an “s” is used to stand for “solid,” and “aq” is used to shorten the word “aqueous.”

To make the concept of equilibrium more useful to a scientist, mathematics must be applied to it. What results is called an equilibrium expression that has a particularequilibrium constant (usually represented as capital K) associated with each unique equilibrium process. For instance, the equilibrium expression for the above reaction is:

K = [Mg{2+}][SO4{2-}]

Here the square brackets – [] – represent the concentration (a number) of the given ion (or undissociated compound) in the solution. What this mathematical expression says is that – the multiplication of the magnesium ion concentration with the concentration of the sulfate ion is equal to the constant number K (the equilibrium constant). Solids (and solvents) do not appear in equilibrium expressions, and so, MgSO4(s) doesn’t enter into the math (if a reactant did appear in the expression it would divide into the product of the product concentrations).

The equilibrium constant, K, is temperature dependent (as we would expect from our observation that more or less of a substance will dissolve with changes in temperature), and Le Chatlier’s principle is obeyed by the equilibrium expression. For instance, if another source of magnesium (such as MgCl2) is added to our saturatedsolution of MgSO4, the product – [Mg{2+}][SO4{2-}] – increases above the value of K, and the equilibrium must re-establish (in this case by increasing the amount of solid MgSO4 in the bottom of the glass until the product of those two ion concentrations once again equals the K associated with the dissociation of the epsom salt). Another point that can be realized with a little more consideration is that if K is large a relatively larger amount of products will be present (in this case the maximum epsom salt concentration would be higher), and as K gets smaller, a relatively smaller amount products will be present at equilibrium.

Now, we may discuss the spraying of alumina into our atmosphere. There are many natural gases that are present in the atmosphere. All of them are participating in dynamic equilibria (plural), and all are subject to Le Chatelier’s principle.

One major implication of reactions occurring in the atmosphere is that they are not contained (such as epsom salts dissolved in a glass of water). When substances are added to the atmosphere, and they participate in any of these equilibria, products and reactants are constantly swept around non-uniformly by wind currents. As a result, these equilibria must constantly shift in efforts to re-establish equilibrium concentrations in the immediate vicinity. If the products get swept one way, the reaction in their vicinity will move towards reactants while the reactions occurring in the vicinity of the reactants strewn the other way will move towards products. In the case of substances naturally occurring in the atmosphere, this jostling about of local equilibria are stable (on average) or regulated by natural, cyclic processes when there are no outside influences.

Alumina (aluminum oxide – Al2O3) is relatively inert when contained. For example, one probably shouldn’t spend their life worrying about drinking from aluminum cans. The Al2O3 layer inside of the can is so insoluble that the concentration in the solution of soda will likely never reach harmful levels (unless special conditions are present). In other words, the equilibrium constant (K) for the oxide dissolving in the soda is a very small number.

However, the smallness of the equilibrium constant becomes much less relevant when the reaction is not contained in a can. By consideration of Le Chatelier’s principle, we can imagine that if we constantly extracted the minute amount that is dissolved in the soda, we could eventually dissolve the entire can (at least until a hole was eaten in the container and the soda spilled). Thus, when a substance such as alumina (that can & WILL participate in atmospheric equilibria) is released into the air, it doesn’t really matter how small the K is. The consequences of such actions could be as severe as major depletion of the ozone layer leading to dramatic increases in UV radiation reaching deeper into the atmosphere that will not only cause severely increased risks of cancer, but in addition, the increased presence of UV radiation will also drive reactions that previously didn’t have enough energy to occur (to any significant degree) under normal conditions.

In the immediate time after the release of alumina, I would expect pH sensitive equilibria to be the first affected because Al2O3 can react with water in the following way:

Al2O3 + H20 <—> Al2O2OH{1+} + OH{1-}

As a result, any equilibrium that had water or the hydroxide ion (OH{1-}) as a participant would immediately be shifted. Such a shift would have to be adjusted for, and this would have a butterfly effect on all associated equilibria that occur “down the line.” The results will likely be disastrous.

The atmosphere is not the only part of the earth system that will be affected. The Al2O3 will eventually fall to the surface either as alumina or as aluminum hydroxide – Al(OH)3. As these substances reach the surface of the earth, the pH of the soil will inevitably increase leading to severe losses in natural vegetation which act to remove CO2 from our air and produce oxygen. As the foliage dies, the soil will be unprotected and desertification will result. There will likely be severe dust storms for many many years until enough natural erosion can occur to replace the contaminated soil and once again restore suitable growing conditions. Further, whatever toxicological conditions that result from increased intake of aluminum will manifest in younger generations that will survive long enough to incur the effects.

If alumina is being sprayed into the atmosphere, it should be discontinued immediately. It’s high time for the government and others with clout to quit only looking at short-term effects and consider the much more important long-term consequences of their actions. We all may be dead in the long run, but our descendents will continue to pay for our actions.


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