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Iron bacteria

In the management of water-supplywells, iron bacteria are bacteria that derive the energy they need to live and multiply by oxidizingdissolved ferrous iron (or the less frequently available manganese). The resulting ferric oxide is insoluble, and appears as brown gelatinous slime that will stain plumbing fixtures, and clothing or utensils washed with the water carrying it. They are known to grow and proliferate in waters containing as low as 0.1mg/l of iron. However, at least 0.3 ppm of dissolved oxygen is needed to carry out oxidation.

Common effects of excess iron in water are a reddish-brown color, stained laundry and poor tasting coffee. An equally common but less well understood problem is infestation of water supplies with iron bacteria. Iron bacteria are a natural part of the environment in most parts of the world. These microorganisms combine dissolved iron or manganese with oxygen and use it to form rust-colored deposits. In the process, the bacteria produce a brown slime that builds up on well screens, pipes, and plumbing fixtures.

Bacteria known to feed on iron are Thiobacillus ferrooxidansandLeptospirillum ferrooxidans.

Habitat

Iron bacteria colonize the transition zone where de-oxygenated water from an anaerobic environment flows into an aerobic environment. Groundwater containing dissolved organic material may be de-oxygenated by microorganisms feeding on that dissolved organic material. Where concentrations of organic material exceed the concentration of dissolved oxygen required for complete oxidation, microbial populations with specialized enzymes can reduce insoluble ferric oxide in aquifersoils to soluble ferrous hydroxide and use the oxygen released by that change to oxidize some of the remaining organic material:

H2O + Fe2O3→ 2Fe(OH)2 + O2
(water) + (Iron[III] oxide) → (Iron[II] hydroxide) + (oxygen)

When the de-oxygenated water reaches a source of oxygen, iron bacteria use that oxygen to convert the soluble ferrous iron back into an insoluble reddish precipitate of ferric iron:

2Fe(OH)2 + O2→ H2O + Fe2O3
(Iron[II] hydroxide) + (oxygen) → (water) + (Iron[III] oxide)

Since the latter reaction is the normal equilibrium in our oxygen atmosphere while the first requires biological coupling with a simultaneous oxidation of carbon, organic material dissolved in water is often the underlying cause of an iron bacteria population. Groundwater may be naturally de-oxygenated by decaying vegetation in swamps; and useful mineral deposits of bog iron ore have formed where that groundwater has historically emerged to be exposed to atmospheric oxygen. Anthropogenic sources like landfillleachate, septic drain fields, or leakage of light petroleum fuels like gasoline are other possible sources of organic materials allowing soil microbes to de-oxygenate groundwater.

A similar reversible reaction may form black deposits of manganese dioxide from dissolved manganese, but is less common because of the relative abundance of iron (5.4 percent) in comparison to manganese (0.1 percent) in average soils. Other conditions associated with iron bacteria result from the anaerobic aqueous environment rather than the iron bacteria visibly colonizing that habitat. Corrosion of pipes is another source of soluble iron for the first reaction above and the sulfurous smell of rot or decay results from enzymatic conversion of soil sulfates to volatile hydrogen sulfide as an alternative source of oxygen in anaerobic environments.

Possible indicators

Clues which indicate that iron bacteria may be present in well water:

  • Iron bacteria often produce unpleasant tastes and odors commonly reported as
    • swampy
    • oily or petroleum
    • cucumber
    • sewage
    • rotten vegetation
    • musty
The taste or odor may be more noticeable after the water has not been used for some time.
  • Iron bacteria will usually cause yellow, orange, red, or brown stains and colored water
  • It is sometimes possible to see a rainbow colored, oil-like sheen on the water.
  • Iron bacteria produce a sticky slime which is typically rusty in color, but may be yellow, brown, or grey.
  • A feathery or filamentous growth may also be seen, particularly in standing water such as a toilet tank.

The dramatic effects of iron bacteria are seen in surface waters as brown slimy masses on stream bottoms and lakeshores or as an oily sheen upon the water. More serious problems occur when bacteria build up in well systems. Iron bacteria in wells do not cause health problems, but they can reduce well yields by clogging screens and pipes.

Prevention

Iron bacteria can be introduced into a well or water system during drilling, repair, or service. Elimination of iron bacteria once a well is heavily infested can be extremely difficult. Normal treatment techniques may be only partly effective. Good housekeeping practices can prevent iron bacteria from entering a well:

  • Water placed in a well for drilling, repair, or priming of pumps should be disinfected, and should never be taken from a lake or pond.
  • The well casing should be watertight, properly capped, and extend a foot or more above ground.
  • When pumps, well pipes, and well equipment are repaired, they should not be placed on the ground where they could pick up iron bacteria.
  • The well, pump, and plumbing should be disinfected when repaired.

Control

Treatment techniques which may be successful in removing or reducing iron bacteria include physical

Metallic hydrogen

Metallic hydrogen is a state of hydrogen which results when it is sufficiently compressed and undergoes a phase transition; it is an example of degenerate matter. Solid metallic hydrogen is predicted to consist of a crystal lattice of hydrogen nuclei (namely, protons), with a spacing which is significantly smaller than the Bohr radius. Indeed, the spacing is more comparable with the de Broglie wavelength of the electron. The electrons are unbound and behave like the conduction electrons in a metal. In liquid metallic hydrogen, protons do not have lattice ordering; rather, it is a liquid system of protons and electrons.

History

Theoretical predictions

Metallization of hydrogen under pressure

Though at the top of the alkali metal column in the periodic table, hydrogen is not, under ordinary conditions, an alkali metal. In 1935 however, physicists Eugene Wigner and Hillard Bell Huntington predicted that under an immense pressure of ~ ( or ), hydrogen atoms would display metallic properties, losing hold over their electrons. Since then, metallic hydrogen has been described as "the holy grail of high-pressure physics".

The initial prediction about the amount of pressure needed was eventually proven to be too low. Since the first work by Wigner and Huntington the more modern theoretical calculations were pointing toward higher but nonetheless potentially accessible metallization pressures. Techniques are being developed for creating pressures of up to , higher than the pressure at the center of the Earth, in hopes of creating metallic hydrogen.

Liquid metallic hydrogen

Helium-4 is a liquid at normal pressure and temperatures near absolute zero, a consequence of its high zero-point energy (ZPE). The ZPE of protons in a dense state is also high, and a decline in the ordering energy (relative to the ZPE) is expected at high pressures. Arguments have been advanced by Neil Ashcroft and others that there is a melting point maximum in compressed hydrogen, but also that there may be a range of densities (at pressures around ) where hydrogen may be a liquid metal, even at low temperatures.

Superconductivity

In 1968, Ashcroft put forward that metallic hydrogen may be a superconductor, up to room temperature (~), far higher than any other known candidate material. This stems from its extremely high speed of sound and the expected strong coupling between the conduction electrons and the lattice vibrations.

Possibility of novel types of quantum fluid

Presently known "super" states of matter are superconductors, superfluid liquids and gases, and supersolids. It was predicted by Egor Babaev that, if hydrogen and deuterium have liquid metallic states, they may have ordered states in quantum domains which cannot be classified as superconducting or superfluid in usual sense but represent two possible novel types of quantum fluids: "superconducting superfluid" and "metallic superfluid". These were shown to have highly unusual reactions to external magnetic fields and rotations, which might represent a route for experimental verification of these possible new states of matter. It has also been suggested that, under the influence of magnetic field, hydrogen may exhibit phase transitions from superconductivity to superfluidity and vice-versa.

Lithium doping reduces requisite pressure

In 2009, Zurek et al. predicted that the alloy LiH6 would be a stable metal at only of the pressure required to metallize hydrogen, and that similar effects should hold for alloys of type LiHn and possibly other alloys of type ?Lin.

Experimental pursuit

Metallization of hydrogen in shock-wave compression

In March 1996, a group of scientists at Lawrence Livermore National Laboratory reported that they had serendipitously produced, for about a microsecond and at temperatures of thousands of kelvins and pressures of over a million atmospheres (>100 GPa), the first identifiably metallic hydrogen. The team did not expect to produce metallic hydrogen, as it was not using solid hydrogen, thought to be necessary, and was working at temperatures above those specified by metallization theory. Previous studies in which solid hydrogen was compressed inside diamond anv

Thermoelectric effect

This page is about the thermoelectric effect as a physical phenomenon. For applications of the thermoelectric effect, seethermoelectric materials, thermoelectric generator, and thermoelectric cooling.

The thermoelectric effect is the direct conversion of temperature differences to electric voltage and vice versa. A thermoelectric device creates a voltage when there is a different temperature on each side. Conversely when a voltage is applied to it, it creates a temperature difference (known as the Peltier effect). At atomic scale (specifically, charge carriers), an applied temperature gradient causes charged carriers in the material, whether they are electrons or electron holes, to diffuse from the hot side to the cold side, similar to a classical gas that expands when heated; hence, the thermally induced current.

This effect can be used to generate electricity, to measure temperature, to cool objects, or to heat them or cook them. Because the direction of heating and cooling is determined by the polarity of the applied voltage, thermoelectric devices make very convenient temperature controllers.

Traditionally, the term thermoelectric effect or thermoelectricity encompasses three separately identified effects, the Seebeck effect, the Peltier effect, and the Thomson effect. In many textbooks, thermoelectric effect may also be called the Peltier–Seebeck effect. This separation derives from the independent discoveries of French physicist Jean Charles Athanase Peltier and Estonian-German physicist Thomas Johann Seebeck. Joule heating, the heat that is generated whenever a voltage is applied across a resistive material, is somewhat related, though it is not generally termed a thermoelectric effect (and it is usually regarded as being a loss mechanism due to non-ideality in thermoelectric devices). The Peltier–Seebeck and Thomson effects can in principle be thermodynamically reversible, whereas Joule heating is not.

Seebeck effect

The Seebeck effect is the conversion of temperature differences directly into electricity.

Seebeck discovered that a compass needle would be deflected when a closed loop was formed of two metals joined in two places with a temperature difference between the junctions. This is because the metals respond differently to the temperature difference, which creates a current loop, which produces a magnetic field. Seebeck, however, at this time did not recognize there was an electric current involved, so he called the phenomenon the thermomagnetic effect, thinking that the two metals became magnetically polarized by the temperature gradient. The Danish physicist Hans Christian Ørsted played a vital role in explaining and conceiving the term "thermoelectricity".

The effect is that a voltage, the thermoelectric EMF, is created in the presence of a temperature difference between two different metals or semiconductors. This causes a continuous current in the conductors if they form a complete loop. The voltage created is of the order of several microvolts per kelvin difference. One such combination, copper-constantan, has a Seebeck coefficient of 41 microvolts per kelvin at room temperature.

In the circuit:

(which can be in several different configurations and be governed by the same equations), the voltage developed can be derived from:

V = \int_{T_1}^{T_2} \left( S_\mathrm{B}(T) - S_\mathrm{A}(T) \right) \, dT.

SA and SB are the Seebeck coefficients (also called thermoelectric powerorthermopower) of the metals A and B as a function of temperature, and T1 and T2 are the temperatures of the two junctions. The Seebeck coefficients are non-linear as a function of temperature, and depend on the conductors' absolute temperature, material, and molecular structure. If the Seebeck coefficients are effectively constant for the measured temperature range, the above formula can be approximated as:

V = (S_\mathrm{B} - S_\mathrm{A}) \cdot (T_2 - T_1).

The Seebeck effect is commonly used in a device called a thermocouple (because it is made from a coupling or junction of materials, usually metals) to measure a temperature difference directly or to measure an absolute temperature by setting one end to a known temperature. A metal of unknown composition can be classified by its thermoelectric effect if a metallic probe of known composition, kept at a constant temperature, is held in contact with it. Industrial quality control instruments use this Seebeck effect to identify metal alloys. This is known as thermoelectric alloy sorting.

Several thermocouples connected in series are called a thermopile, which is sometimes constructed in order to increase the output voltage since the voltage induced over each individual couple is small.

This is also the principle at work behind thermoelectric generators (such as radioisotope thermoelectric generators or RTGs) which are used for creating power from heat differentials.

The Seebeck effect is due to two effects: charge carrier diffusion and phonon drag (described below).

Thermopower

The thermopowe


From Yahoo Answers

Question:For example, what would be the chemical reaction (as simple as can be) of an iron nail rusting?

Answers:The rusting of iron is an electrochemical process that begins with the transfer of electrons from iron to oxygen.The rate of corrosion is affected by water and accelerated by electrolytes, as illustrated by the effects of road salt (calcium chloride) on the corrosion of automobiles. The key reaction is the reduction of oxygen: O2 + 4 e- + 2 H2O 4 OH- Because it forms hydroxide ions, this process is strongly affected by the presence of acid. Indeed, the corrosion of most metals by oxygen is accelerated at low pH. Providing the electrons for the above reaction is the oxidation of iron that may be described as follows: Fe Fe2+ + 2 e The following redox reaction also occurs in the presence of water and is crucial to the formation of rust: 4 Fe2+ + O2 4 Fe3+ + 2 O2 Additionally, the following multistep acid-base reactions affect the course of rust formation: Fe2+ + 2 H2O -> Fe(OH)2 + 2 H+ Fe3+ + 3 H2O -> Fe(OH)3 + 3 H+ as do the following dehydration equilibria: Fe(OH)2 -> FeO + H2O Fe(OH)3 ->FeO(OH) + H2O 2 FeO(OH) Fe2O3 + H2O

Question:and how do you represent them in a word equation ? Pleasee helpp as soon as possiblee :)

Answers:Rusting of iron EXPLANATION OF THE RUSTING OF IRON When iron rusts a spontaneous redox reaction occurs, between the oxygen and iron. If water is added the rusting occurs more rapidly. Iron (s) + Oxygen(g)----------> Iron (III) oxide(s) When water is added to iron, an electrochemical cell is created that has a distinct anode and cathode. If an iron nail is placed in agar or gel in which ferric cyanide ions and phenolphthalein indicator have been placed, the ends of the nail turn blue and the middle of the nail turns red. The blue colour is caused by ferricyanide indicator reaction with the iron ions, and the red colour (pink) is due to reaction between hydroxide ions and phenolphthalein. How are these ions produced. At one spot on the nail (the Anodic site of our electrochemical cell) Iron loses electrons (is oxidized) to form iron (II) ions. Fe (s) ------> Fe2+ (aq) + 2e- At another spot on the nail the oxygen in the air combines with water and forms hydroxide ions. 1/2 O2 (g) + H2O (l) + 2e- --------> 2OH- (aq) In the presence of oxygen the iron further oxidizes at the anode (loses electrons) to become iron (III) ions. Fe 2+ (aq) ------> Fe3+ (aq) + e- The iron (III) ions and the hydroxide combine to form rust ( flaky brown substance) . 2 Fe 3+ (aq) + 6OH- (aq) -----> Fe2O3 (s) + 3 H2O (l) Notice that water is required for the reaction at the cathode but produced in the overall reaction. It therefore is acting like a homogeneous catalyst, to speed up the rusting of iron. In essence the water and oxygen make it easier for iron to rust. As with an electrochemical cells the electrons flow from the anode to the cathode. Oxygen and water are both need to speed the rusting process in metals.

Question:Can someone tell me all the metals are on the periodic table that can rust when in contact with oxygen? And if aluminum was a superhero, what would be his weakness? thanks If aluminum were a superhero, coiuld oxygen and mercury be a villans? Mercury makes sense, but qxygen also seems like he could be a sidekick. I was thinking that aluminum's friends could be iron and so he saves them from oxygen. What do you think?

Answers:The other answerer is correct - most metals will react with oxygen except for a few "noble" metals like gold, silver, and platinum. Alkali metals are extremely reactive and will easily corrode in air (some ignite spontaneously in moist air). However, there are some metals that form a thin layer of hard, impervious oxide when exposed to air and then stop corroding, a phenomenon known as passivation. Aluminum, titanium, niobium, zirconium, and several other metals fall in this category. Despite being very reactive, these metals resist corrosion well due to this oxide layer, which stops corrosive chemicals from reaching fresh metal. If formation of this oxide layer is disrupted somehow, the metal will corrode, often quickly. This brings us to the question about aluminum's weakness. If the aluminum was a liquid rather than a solid, the very thin oxide layer could easily be removed or broken. To make aluminum into a liquid, it must be melted. However, mercury and gallium - both liquid at or near room temperature - will dissolve solid aluminum metal to form a liquid alloy. This alloy will not be protected by an oxide layer and will easily be corroded by oxygen or water - the aluminum will be converted to white, crumbly aluminum oxide (forming clumps instead of a protective layer). Acids and alkalis will also dissolve aluminum, but the answer that your teacher is looking for is probably mercury.

Question:Hey! I am working on corrosion in chem at the moment. I don't quite follow how salt speeds up the rate of corrosion iron... I figure there is some large significance since the nails i had in a higher salinity lost more weight then those in a lower salinity... I am just not sure what it is. Help please!

Answers:When metals rust, it forms metal oxides. But by definition, Catalyst Speeds up the process of the Rxn but does not get used up. However as you can see in the equation below, Oxygen is consumed with metal to for metal oxide. Therefore Oxygen CAN'T be a catalyst. It is a reactant. Fe + O2 ---> Fe2O3 The Correct answer is SALT. (Sodium in NaCl for an example) Rust is the process of metal oxidizing into metal oxide. The oxidation needs a transfer medium to occur, such as water. Salt in the water speeds up the oxidation process by acting as a catalyst. The salt is never used up and the problem with salt is that even when water is gone, it remains and wait to restart the forming of rust once more moisture is present. This is why metals rust much easily in salty areas such as beaches and cold states where they use salt to melt the icy roads.