Acronym Definition
PTUM Please Tell Us More
PTUM Personal Tag-Out Users Manual
PTUM Program Technical University of Moldova
PTUM Program Technische Universität München (German University)
PTUM Personal Telephone-based Unified Messsaging
PTUM Personal Terrain Update Module
PTUM Pan Tilt Unit Management
PTUM Paper Tape Unit Management
PTUM Parallel Transfer Unit Management
PTUM Participating Test Unit Management
PTUM Police Tactical Unit (Hong Kong) Management
PTUM Polskie Towarzystwo Ubezpiecze Management
PTUM Portable Terminal Unit Management
PTUM Portable Test Unit Management
PTUM Power Transfer Unit (aircraft hydraulic backup system) Management
PTUM Prise de Terrain en U (aeronautics) Management
PTUM Process Technology Utilization Management
PTUM Propylthiouracil Management
PTUM Providence Teachers Union Management
PTUM Punjab Technical University Management
PTUM Propylthiouracil Made
PTUM Programa de Transportes Urbanos no Muncipio
An antacid is any substance, generally a base, which counteracts stomach
acidity. In other words, antacids are stomach acid neutralizers.
Action mechanism
Antacids perform a neutralization reaction, i.e. they buffer gastric acid,
raising the pH to reduce acidity in the stomach. When gastric hydrochloric acid
reaches the nerves in the gasitrointestinal mucosa, they signal pain to the
central nervous system. This happens when these nerves are exposed, as in peptic
ulcers. The gastric acid may also reach ulcers in the esophagus or the duodenum.
Other mechanisms may contribute, such as the effect of aluminum ions inhibiting
smooth muscle cell contraction and delaying gastric emptying.
Indications
Antacids are taken by mouth to relieve heartburn, the major symptom of
gastroesophageal reflux disease, or acid indigestion. Treatment with antacids
alone is symptomatic and only justified for minor symptoms. Peptic ulcers may
require H2-receptor antagonists or proton pump inhibitors.
The usefulness of many combinations of antacids is not clear, although the
combination of magnesium and aluminum salts may prevent alteration of bowel
habits.
Side effects
* Aluminum hydroxide: may lead to the formation of insoluble
aluminum-phosphate-complexes, with a risk for hypophosphatemia and osteomalacia.
Although aluminum has a low gastrointestinal absorption, accumulation may occur
in the presence of renal insufficiency. Aluminum-containing drugs may cause
obstipation.
* Magnesium hydroxide: has laxative properties. Magnesium may accumulate in
patients with renal failure leading to hypermagnesemia, with cardiovascular and
neurological complications. See Milk of magnesia.
* Carbonate: regular high doses may cause alkalosis, which in turn may result in
altered excretion of other drugs, and kidney stones. A chemical reaction between
the carbonate and hydrochloric acid may produce carbon dioxide gas. This causes
gastric distension which may not be well tolerated.
* Calcium: compounds containing calcium may increase calcium output in the
urine, which might be associated to renal stones. Calcium salts may cause
Constipation.
* Sodium: increased intake of sodium may be deleterious for arterial
hypertension, heart failure and many renal diseases.
Interactions
Altered pH or complex formation may alter the bioavailability of other drugs,
such as tetracycline. Urinary excretion of certain drugs may also be affected.
Problems with reduced stomach acidity
Reduced stomach acidity may result in an impaired ability to digest and absorb
certain nutrients, such as iron and the B vitamins. Since the low pH of the
stomach normally kills ingested bacteria, antacids increase the vulnerability to
infection.
Drug names
Examples of antacids (brand names may vary in different countries).
* Aluminum hydroxide (Amphojel®, AlternaGEL®)
* Magnesium hydroxide (Phillips’® Milk of Magnesia)
* Aluminum hydroxide and magnesium hydroxide (Maalox®, Mylanta®)
* Aluminum carbonate gel (Basaljel®)
* Calcium carbonate (Alcalak®, Calcium Rich TUMS®, Quick-Eze®, Rennie®, Titralac®,
Rolaids®)
* Sodium bicarbonate (Bicarbonate of soda, Alka-Seltzer®)
* Hydrotalcite (Mg6Al2(CO3)(OH)16 · 4(H2O); Talcid®)
* Bismuth subsalicylate (Pepto-Bismol)
* Magaldrate + Simethicone (Pepsil)
Aluminium hydroxide, Al(OH)3, is the most stable form of aluminium in normal
conditions. It is found in nature as the mineral gibbsite (also known as
hydrargillite). Closely related are aluminium oxide hydroxide, AlO(OH), and
aluminium oxide, Al2O3, differing only by loss of water. These compounds
together are the major components of the aluminium ore bauxite.
Chemistry
Gibbsite has a typical metal hydroxide structure with hydrogen bonds. It is
built up of double layers of hydroxyl groups with aluminum ions occupying
two-thirds of the octahedral holes between the two layers.
Aluminum hydroxide is amphoteric. It dissolves in acid, forming Al(H2O)63+ or
its hydrolysis products. It also dissolves in strong alkali, forming Al(OH)4-.
Aluminum hydroxide is produced in the Bayer process as an intermediate in the
production of aluminum metal.
Pharmacology
Pharmacologically, this compound is used as an antacid under names such as Alu-Cap,
Aludrox or Pepsamar. The hydroxide reacts with excess acid in the stomach,
reducing its acidity. This decrease of acidity of the contents of the stomach
may in turn help to relieve the symptoms of ulcers, heartburn or dyspepsia. It
can also cause constipation and is therefore often used with magnesium
carbonate, which has counterbalancing laxative effects. This compound is also
used to control phosphate levels in the blood of people suffering from kidney
failure.
Aluminium hydroxide is included as an adjuvant in some vaccines, since it
contributes to induction of a good antibody (Th2) response. However, it has
little capacity to stimulate cellular (Th1) immune responses, important for
protection against many pathogens (Petrovsky and Aguilar, 2004).
Because the brain lesions found in Alzheimer's disease contain aluminium, there
is concern that consumption of excess aluminium compounds may cause or
contribute to the development of this and other neurodegenerative diseases
(Perl, 2006, Kawahara, 2005). In addition, elevated aluminium levels in blood,
resulting from kidney dialysis with well water containing high aluminium, result
in dementia that is similar to but probably different from that of Alzheimer's
disease (Carpenter, 2001). However, this hypothesis is controversial
Magnesium hydroxide, Mg(OH)2, otherwise known as milk of magnesia, is commonly
used as an antacid or a laxative. The mineral form of magnesium hydroxide is
known as brucite. Magnesium hydroxide interferes with the absorption of folic
acid and iron. The diarrhea caused by magnesium hydroxide carries away much of
the body's supply of potassium, and failure to take extra potassium will lead to
muscle cramps.
It has very low solubility in water, and has a Ksp value of 1.5x10-11.
It does not completely dissociate and therefore is a weak base.
Preparation
From a Magnesium salt:
In words:
Solid Magnesium nitride + Water = Ammonia + Aqueous Magnesium Hydroxide
From a Magnesium base:
In words:
Solid Magnesium oxide + Water = Aqueous Magnesium Hydroxide
From pure Magnesium metal:
In Words:
Solid Magnesium + Water = Aqueous Magnesium Hydroxide + Hydrogen Gas
Uses
Magnesium hydroxide is used as an antacid to neutralize stomach acid. In
industries, it is used as a non-hazardous alkali to neutralise acidic
wastewaters. In addition, magnesium hydroxide, better known as its common name
Milk of Magnesia can be used as an non antiperspirant armpit deodorant. It is
also used in bleaching solutions to whiten clothes.
Simethicone, is an oral anti-foaming agent used to reduce bloating, discomfort
and pain caused by excess gas in the stomach or intestinal tract. It is a
mixture of polydimethylsiloxane and silica gel.
Chemical action
Simethicone is an anti-foaming agent that reduces the surface tension of gas
bubbles, causing them to combine into larger bubbles in the stomach that can be
passed more easily by burping. Simethicone does not reduce the quantity of gas
in the digestive tract, it only increases the rate at which it exits the body,
and is ineffective in the intestine. Simethicone is not absorbed by the body
into the bloodstream, and is therefore considered relatively safe, with sources
reporting the worst side effects as bloating, constipation, diarrhea, gas and
heartburn. Simethicone solutions of differing concentration also have industrial
applications for reducing foaming in certain chemical processes.
Dosage
Simethicone comes in many different oral forms, which have differing usual
dosages. Though there are standard dosages for adults and teens, dosages for
children should be determined by a doctor.
* Capsules or tablets:
: Adults and teenagers: 60 to 125 milligrams (mg) four times a day, after meals
and at bedtime. No more than 500 mg should be taken in twenty-four hours.
* Chewable tablets:
:Adults and teenagers: 40 to 125 mg four times a day, after meals and at bedtime
or the dose may be 150 mg three times a day, after meals. No more than 500 mg
should be taken in twenty-four hours.
* Suspension:
:Adults and teenagers: 40 to 95 mg four times a day, after meals and at bedtime.
No more than 500 mg should be taken in twenty-four hours.
* Dissolving membrane:
:A thin dissolving membrane dose that is placed on the tongue and dissolves with
saliva
Availability
Simethicone is generally available over the counter under many trade names in
varying dosage sizes, including:
* Air-X (Thailand)
* Baby's Own Infant Drops
* Espumisan
* Flatulex
* Gas Relief
* Gas-X
* GasAid
* Gasvan (Serbia)
* Genasyme
* Imodium Advanced
* lnfacol
* lnfacon
* Kremil-S
* Lefax (Germany)
* Little Tummies
* Maalox Anti-Gas
* Maalox Max
* Minifom (Norway, Sweden)
* My Baby Gas Relief Drops
* Mylanta Gas
* Mylanta Gas Relief
* Mylicon Drops
* Ovol
* Phazyme
* Simeco
* Triaerom (Peru)
* WindEze (UK)
Other uses
While not indicated on the label, it has been reported that Simethicone is
sometimes used by patients before a gastroscopy or a radiography of the bowel.
Simethicone is also used in some detergents when foaming is unwanted.
Despite being useful in the treatment of gas, simethicone does not appear to be
useful in the treatment of infant colic.
Magaldrate (INN), is a common antiacid drug that is used for the treatment of
duodenal and gastric ulcers, esophagitis from gastroesophageal reflux.
Formula
Al5Mg10(OH)31(SO4)2 H2O
Available forms
Magaldrate is available in the form of oral suspension or tablets.
Interactions and adverse reactions
Magaldrate may negatively influence drugs like tetracyclines, benzodiazepines,
and indomethacin. High doses or prolonged usage may lead to an increment of
defecation and a reduction in feces consistance. In some cases it can alter the
functionality of the gastrointestinal tract, occasionally provoking constipation
or diarrhea.
Bismuth subsalicylate
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Bismuth subsalicylate, with a chemical formula C7H5BiO4, is the active
ingredient in the popular medication Pepto-Bismol that is used to treat nausea,
heartburn, indigestion, upset stomach, diarrhea, and other temporary discomforts
of the stomach and gastrointestinal tract. It is also the main ingredient of
Kaopectate (since 2003, replacing attapulgite).
It displays anti-inflammatory action (due to salicylic acid) and also acts as an
antacid and mild antibiotic.
It can also cause a black tongue and black stools in some users of the drug,
when it combines with trace amounts of sulfur in their saliva and
gastrointestinal tract. This discoloration is temporary and harmless.
Children should not take medication with Bismuth subsalicylate while recovering
from the flu or chicken pox, as epidemiologic evidence points to an association
between the use of salicylate containing medications during certain viral
infections and the onset of Reye's syndrome
Sodium bicarbonate is the chemical compound with the formula NaHCO3. Because it
has long been known and is widely used, the salt has many other names including
sodium hydrogencarbonate, sodium bicarb, baking soda, bread soda, cooking soda,
bicarb soda or bicarbonate of soda. The word saleratus, from Latin sal ?ratus
meaning "aerated salt", was widely used in the 19th century for both sodium
bicarbonate and potassium bicarbonate. The term has now fallen out of common
usage.
It is soluble in water. Sodium bicarbonate is a white solid that is crystalline
but often appears as a fine powder. It has a slight alkaline taste resembling
that of sodium carbonate. It is a component of the mineral natron and is found
dissolved in many mineral springs. The natural mineral form is known as
nahcolite. It is also produced artificially.
Production
M Solvay process
NaHCO3 is mainly prepared by the Solvay process, which is the reaction of sodium
chloride, ammonia, and carbon dioxide in water. It is produced on the scale of
about 100,000 ton/year (year: 2001).
Commercial quantities of baking soda are also produced by this method: soda ash,
mined in the form of the ore trona, is dissolved in water and treated with
carbon dioxide. Sodium bicarbonate precipitates as a solid from this method:
Na2CO3 + CO2 + H2O → 2 NaHCO3
Chemistry
Acid-base reactions
NaHCO3 is a salt which consists of the ions Na+ and the bicarbonate anion,
HCO3-. It has a pKa of 6.3 in water which causes aqueous solutions to be mildly
alkaline:
HCO3- + H2O ? H2CO3 + OH-
Reaction of sodium hydroxide with carbon dioxide
NaHCO3 may be obtained by the reaction of carbon dioxide with an aqueous
solution of sodium hydroxide :
The initial reaction produces sodium carbonate:
CO2 + 2NaOH → Na2CO3 + H2O
Further addition of carbon dioxide produces sodium bicarbonate, which at
sufficiently high concentration will precipitate out of solution:
Na2CO3 + CO2 + H2O → 2NaHCO3
Decomposition
Treatment of sodium bicarbonate with acids, releases carbon dioxide and water:
NaHCO3 + HCl → NaCl + H2O + CO2 (gas)
NaHCO3 + CH3COOH → CH3CO2Na + H2O + CO2 (gas)
Thermal decomposition
Above 60 °C, it does gradually decomposes into sodium carbonate, water and
carbon dioxide. The conversion is fast at 200 °C:
2NaHCO3 → Na2CO3 + H2O + CO2
Most bicarbonates undergo this dehydration reaction. Further heating converts
the carbonate into the oxide:
Na2CO3 → Na2O + CO2
These conversions are relevant to the use of NaHCO3 as a fire-suppression agent
("BC powder") in some dry powder fire extinguishers.
Applications
Cooking
M leavening agent
Sodium bicarbonate is primarily used in cooking (baking) where it reacts with
other components to release carbon dioxide, that helps dough "rise." The acidic
compounds that induce this reaction include cream of tartar, lemon juice, yogurt
etc., hence sodium bicarbonate can be substituted for baking powder provided
sufficient acid reagent is also added to the recipe. . Many forms of baking
powder contain sodium bicarbonate combined with cream of tartar.
* A small amount can be added to a beef stew to make tough meat tenderize
faster.
* It is used by the Chinese to make meat tender for stir fries.
* Was formerly used as a source of carbon dioxide for soda water.
* Can be used when preparing tomato sauce to neutralize the tomato's acidity.
* It is added to beans in water to prevent flatulence produced by digesting
them.
* It is effective in extinguishing grease fires which may occur when deep
frying.
* In Thailand, soaking insects in baking soda for three to five hours prior to
cooking produces a sweeter, more mushroom-like taste in the finished cuisine.
* Freshly cut fruit can be soaked in sodium bicarbonate solution for a short
while to prevent yellowing.
For neutralization of acids and bases
The reaction of acids with sodium bicarbonate is a common method for
neutralizing acid and base spills. The advantage of sodium bicarbonate is that
it is amphoteric, hence reacts with acids and bases. For example, with sulfuric
acid:
2 NaHCO3 + H2SO4 → Na2SO4 + 2 H2O + 2 CO2
With sodium hydroxide:
NaHCO3 + NaOH → Na2CO3 + H2O
Furthermore, it is relatively innocuous, that there is no harm in using excess
sodium bicarbonate.
A wide variety of applications follow from its neutralization properties
including ameliorating the effects of white phosphorus in incendiary bullets,
from spreading inside a soldier's afflicted wounds.
* It is commonly used to increase the pH and total alkalinity of the water for
pools and spas. Sodium bicarbonate can be added as a simple solution for
restoring the pH balance of water that has a high level of chlorine.
* It is sometimes used in septic tanks to control pH and bacteria.
Miscellaneous and domestic uses
Baking soda has many uses.
As a deodorizer
* An absorbent for moisture and odors e.g; an open box can be left in a
refrigerator for this purpose. However, according to one source, baking soda
does not actually absorb odors well when used in a refrigerator.
* To help relieve itching due to bacterial infections
Medical uses
* It is used as an antacid to treat acid indigestion and heartburn.
* Mixed with water and drunk, it can relieve cystitis.
* Mixed with water in a 10% solution can soften earwax for removal.
* In paramedicine, sodium bicarbonate 7.9% is administered intravenously for
cases of acidosis and overdoses of acidic toxic substances, such as tricyclic
antidepressants and aspirin.
* Adverse reactions to emergency administration include congestive heart
failure, with edema secondary to sodium overload, and the metabolic complication
of hyperosmolarity, metabolic acidosis, and hypernatremia.
* Aids in itch relief from poison ivy rashes.
* Added to a bath or made into a paste it can be used to relieve the itching
caused by chicken pox.
* For local injections of anesthetics subdermally or subcutaneously, it may be
added to lessen the burning sensation of the anesthetic to the patient. For
example. 9 milliliters of lidocaine mixed with 1 milliliter of sodium bicarb in
a 10 cc syringe will greatly lessen the feeling of burning, pressure and overall
pain from the injection.
* Relieves mosquito bites and bee stings (but not wasp stings).
* Recently there have been claims that it cures many types of cancer.
(http://www.cancerfungus.com/)
Cosmetic uses
* It is marketed as a whitener because of its abrasive properties in some
toothpaste brands.
* Is an essential ingredient in the production of Bath Bombsuk ingredients
supplier along with Citric Acid. Used in ratios of 3:1 or 2:1 the "dry"
ingredients are bound together with oil, water and or fragrance or essential
oils and formed into a ball or "bomb" shape usually with the aid of a mould.
When placed in a tub or bath of water the ingredients react with the water and
fizzing occurs whilst fragrance is released. A natural alternative to bubble
bath!
As a cleaning agent
* A paste from baking soda can be very effective when used in cleaning and
scrubbing.
* A solution in warm water will remove the tarnish from silver when the silver
is in contact with a piece of aluminum
foilhttp://www.instructables.com/id/EWU42JVP9EEV2ZBWZL/.
* With water, it cleans the impurities on contact lenses. Rinse completely
before wearing contacts to avoid stinging residue.
* Cleans brushes and combs to prevent residues.
* Use to clean juice, wine, and coffee stains.
* Pouring 1 cup of baking soda down a drain and following with 1/2 gallon of
vinegar will degrease the drain.
* Baking soda was the primary cleaning agent in the restoration of the statue of
liberty.
Other uses
* It is used as a fabric softener in laundry.
* Sodium bicarbonate has been used as a performance enhancer for sprinters, by
countering build up of lactate through induced metabolic alkalosis.
* Baking soda can be used as a low-cost alternative to raise pH in swimming
pools.
* Sodium bicarbonate is used in BC Dry Chemical fire extinguishers as an
alternative to the corrosive ammonium phosphate in ABC extinguishers. The alkali
nature of sodium bicarbonate makes it the only dry powder, excluding Purple-K,
agent allowed for use on commercial deep fat fryers, the agent forms a crust
over the surface similar to the effects of a wet chemical.
* Sodium bicarbonate is often used in the pharmaceutical industry as an additive
to cell culture media. It acts as a weak buffer.
* It is also used in a process for cleaning paint called sodablasting.
* It is used as a base in the production of crack cocaine.
Calcium carbonate is a chemical compound, with the chemical formula CaCO3. It is
a common substance found as rock in all parts of the world, and is the main
component of shells of marine organisms, snails, and eggshells. Calcium
carbonate is the active ingredient in agricultural lime, and is usually the
principal cause of hard water. It is commonly used medicinally as a calcium
supplement or as an antacid.
Occurrence
Calcium carbonate is found naturally as the following minerals and rocks:
* Aragonite
* Calcite
* Vaterite or (μ-CaCO3)
* Chalk
* Limestone
* Marble
* Travertine
To test whether a mineral or rock contains calcium carbonate, strong acids, such
as hydrochloric acid, can be added to it. If the sample does contain calcium
carbonate, it will fizz and produce carbon dioxide and water. Weak acids such as
acetic acid will react, albeit less vigorously. All of the rocks/minerals
mentioned above will react with acid.
Preparation
The vast majority of calcium carbonate used in industry is extracted by mining
or quarrying. Pure calcium carbonate (e.g. for food or pharmaceutical use), can
be produced from a pure quarried source (usually marble) or it can be prepared
by passing carbon dioxide into a solution of calcium hydroxide: the calcium
carbonate precipitates out, and this grade of product is referred to as a
precipitate (abbreviated to PCC).
Ca(OH)2 + CO2 → CaCO3 + H2O
Chemical properties
See also: Carbonate
Calcium carbonate shares the typical properties of other carbonates. Notably:
1. it reacts with strong acids, releasing carbon dioxide:
CaCO3 + 2HCl → CaCl2 + CO2 + H2O
2. it releases carbon dioxide on heating (to above 840 °C in the case of CaCO3),
to form calcium oxide, commonly called quick lime:
CaCO3 → CaO + CO2
Calcium carbonate will react with water that is saturated with carbon dioxide to
form the soluble calcium bicarbonate.
CaCO3 + CO2 + H2O → Ca(HCO3)2
This reaction is important in the erosion of carbonate rocks, forming caverns,
and leads to hard water in many regions.
Uses
The main use of calcium carbonate is in the construction industry, either as a
building material in its own right (e.g. marble) or limestone aggregate for
roadbuilding or as an ingredient of cement or as the starting material for the
preparation of builder's lime by burning in a kiln . A common contaminant is
magnesium carbonate.
Calcium carbonate is widely used as an extender in paints, in particular matte
emulsion paint where typically 30% by weight of the paint is either chalk or
marble.
Calcium carbonate is also widely used as a filler in plastics. Some typical
examples include around 15 to 20% loading of chalk in uPVC drain pipe, 5 to 15%
loading of stearate coated chalk or marble in uPVC window profile. Fine ground
calcium carbonate is an essential ingredient in the microporous film used in
babies' diapers and some building films as the pores are nucleated around the
calcium carbonate particles during the manufacture of the film by biaxial
stretching.
Calcium carbonate is also used in a wide range of trade and DIY adhesives,
sealants, and decorating fillers. Ceramic tile adhesives typically contain 70 to
80% limestone. Decorating crack fillers contain similar levels of marble or
dolomite. It is also mixed with putty in setting Stained glass windows, and as a
resist to prevent glass from sticking to kiln shelves when firing glazes and
paints at high temperature.
Calcium carbonate is widely used medicinally as an inexpensive dietary calcium
supplement, antacid, and/or phosphate binder. It is also used in the
pharmaceutical industry as a base material for tablets of other pharmaceuticals.
Calcium carbonate is known as whiting in ceramics/glazing applications, where it
is used as a common ingredient for many glazes in its white powdered form. When
a glaze containing this material is fired in a kiln, the whiting acts as a flux
material in the glaze.
Used in swimming pools as a pH corrector for maintaining alkalinity "buffer" to
offset the acidic properties of the disinfectant agent.
It is commonly called chalk as it has been a major component of blackboard
chalk. Chalk may consist of either calcium carbonate or gypsum, hydrated calcium
sulfate CaSO4·2H2O.
In North America, calcium carbonate has begun to replace kaolin in the
production of glossy paper. Europe has been practicing this as alkaline
papermaking or acid-free papermaking for some decades. Carbonates are available
in forms: ground calcium carbonate (GCC) or precipitated calcium carbonate
(PCC). The latter has a very fine and controlled particle size, on the order of
2 micron in diameter, useful in coatings for paper.
As a food additive, it is used in some soy milk products as a source of dietary
calcium.
In 1989, a researcher introduced CaCO3 into the Whetstone Brook in Massachusetts
. His hope was that the calcium carbonate would counter the acid in the stream
from acid rain and save the trout that had ceased to spawn. Although his
experiment was a success, it did increase the amounts of aluminum ions in the
area of the brook that was not treated with the limestone. This shows that CaCO3
can be added to neutralize the effects of acid rain in river ecosystems.
Nowadays, calcium carbonate is used to neutralise acidic conditions in both soil
and water.
Calcination Equilibrium
Equilibrium Pressure of CO2 over CaCO3
550 °C 0.055 kPa
587 °C 0.13 kPa
605 °C 0.31 kPa
680 °C 1.80 kPa
727 °C 5.9 kPa
748 °C 9.3 kPa
777 °C 14 kPa
800 °C 24 kPa
830 °C 34 kPa
852 °C 51 kPa
871 °C 72 kPa
881 °C 80 kPa
891 °C 91 kPa
898 °C 101 kPa
937 °C 179 kPa
1082 °C 901 kPa
1241 °C 3961 kPa
Calcination of limestone using charcoal fires to produce quicklime has been
practiced since antiquity by cultures all over the world. The answer to the
question, "how hot does the fire have to be?" is usually given as 825 °C, but
stating an absolute threshold is misleading. Calcium carbonate exists in
equilibrium with calcium oxide and carbon dioxide at any temperature. At each
temperature there is a partial pressure of carbon dioxide that is in equilibrium
with calcium carbonate. At room temperature the equilibrium overwhelmingly
favors calcium carbonate, because the equilibrium CO2 pressure is only a tiny
fraction of the partial CO2 pressure in air, which is about 0.035 kPa. At
temperatures above 550 °C the equilibrium CO2 pressure begins to exceed the CO2
pressure in air. So above 550 °C, calcium carbonate begins to outgas CO2 into
air. But in a charcoal fired kiln, the concentration of CO2 will be much higher
than it is in air. Indeed if all the oxygen in the kiln is consumed in the fire,
then the partial pressure of CO2 in the kiln can be as high as 20 kPa. The table
shows that this equilibrium pressure is not achieved until the temperature is
nearly 800 °C. For the outgassing of CO2 from calcium carbonate to happen at an
economically useful rate, the equilibrium pressure must significantly exceed the
ambient pressure of CO2. And for it to happen rapidly, the equilibrium pressure
must exceed total atmospheric pressure of 101 kPa, which happens at 898 °C.
Solubility of calcium carbonate in water
Solubility in pure water with varying CO2 pressure
Calcium carbonate is poorly soluble in pure water. The equilibrium of its
solution is given by the equation (with dissolved calcium carbonate on the
right):
:
CaCO3 ? Ca2+ + CO32– Ksp = 3.7×10–9 to 8.7×10–9 at 25 °C
where the solubility product for [Ca2+][CO32–] is given as anywhere from Ksp =
3.7×10–9 to Ksp = 8.7×10–9 at 25 °C, depending upon the data source. What the
equation means is that the product of molar concentration of calcium ions (moles
of dissolved Ca2+ per liter of solution) with the molar concentration of
dissolved CO32– cannot exceed the value of Ksp. This seemingly simple solubility
equation, however, must be taken along with the more complicated equilibrium of
carbon dioxide with water (see carbonic acid). Some of the CO32– combines with
H+ in the solution according to:
:
HCO3– ? H+ + CO32– Ka2 = 5.61×10–11 at 25 °C
HCO3– is known as the bicarbonate ion. Calcium bicarbonate is many times more
soluble in water than calcium carbonate -- indeed it exists only in solution.
Some of the HCO3– combines with H+ in solution according to:
:
H2CO3 ? H+ + HCO3– Ka1 = 2.5×10–4 at 25 °C
Some of the H2CO3 breaks up into water and dissolved carbon dioxide according
to:
:
H2O + CO2(dissolved) ? H2CO3 Kh = 1.70×10–3 at 25 °C
And dissolved carbon dioxide is in equilibrium with atmospheric carbon dioxide
according to:
:
where kH = 29.76 atm/(mol/L) at 25°C (Henry constant), being the CO2 partial
pressure.
Calcium Ion Solubility
as a function of CO2 partial pressure at 25 °C
(atm) pH [Ca2+] (mol/L)
10?12 12.0 5.19 × 10?3
10?10 11.3 1.12 × 10?3
10?8 10.7 2.55 × 10?4
10?6 9.83 1.20 × 10?4
10?4 8.62 3.16 × 10?4
3.5 × 10?4 8.27 4.70 × 10?4
10?3 7.96 6.62 × 10?4
10?2 7.30 1.42 × 10?3
10?1 6.63 3.05 × 10?3
1 5.96 6.58 × 10?3
10 5.30 1.42 × 10?2
For ambient air, is around 3.5×10–4 atmospheres (or equivalently 35 Pa). The
last equation above fixes the concentration of dissolved CO2 as a function of ,
independent of the concentration of dissolved CaCO3. At atmospheric partial
pressure of CO2, dissolved CO2 concentration is 1.2×10–5 moles/liter. The
equation before that fixes the concentration of H2CO3 as a function of [CO2].
For [CO2]=1.2×10–5, it results in [H2CO3]=2.0×10–8 moles per liter. When [H2CO3]
is known, the remaining three equations together with
:
H2O ? H+ + OH– K = 10–14 at 25 °C
(which is true for all aqueous solutions), and the fact that the solution must
be electrically neutral,
:2[Ca2+] + = [HCO3–] + 2[CO32–] + [OH–]
make it possible to solve simultaneously for the remaining five unknown
concentrations (note that the above form of the neutrality equation is valid
only if calcium carbonate has been put in contact with pure water or with a
neutral pH solution; in the case where the origin water solvent pH is not
neutral, the equation is modified).
The table on the right shows the result for [Ca2+] and (in the form of pH) as a
function of ambient partial pressure of CO2 (Ksp = 4.47×10?9 has been taken for
the calculation). At atmospheric levels of ambient CO2 the table indicates the
solution will be slightly alkaline. The trends the table shows are
1) As ambient CO2 partial pressure is reduced below atmospheric levels, the
solution becomes more and more alkaline. At extremely low , dissolved CO2,
bicarbonate ion, and carbonate ion largely evaporate from the solution, leaving
a highly alkaline solution of calcium hydroxide, which is more soluble than
CaCO3.
2) As ambient CO2 partial pressure increases to levels above atmospheric, pH
drops, and much of the carbonate ion is converted to bicarbonate ion, which
results in higher solubility of Ca2+.
The effect of the latter is especially evident in day to day life of people who
have hard water. Water in aquifers underground can be exposed to levels of CO2
much higher than atmospheric. As such water percolates through calcium carbonate
rock, the CaCO3 dissolves according to the second trend. When that same water
then emerges from the tap, in time it comes into equilibrium with CO2 levels in
the air by outgassing its excess CO2. The calcium carbonate becomes less soluble
as a result and the excess precipitates as lime scale. This same process is
responsible for the formation of stalactites and stalagmites in limestone caves.
Two hydrated phases of calcium carbonate, monohydrocalcite, CaCO3.H2O, and
ikaite, CaCO3.6H2O, may precipitate from water at ambient conditions and persist
as metastable phases.
Solubility at atmospheric CO2 pressure with varying pH
We now consider the problem of the maximum solubility of calcium carbonate in
normal atmospheric conditions ( = 3.5 × 10?4 atm) when the pH of the solution is
adjusted. This is for example the case in a swimming pool where the pH is
maintained between 7 and 8 (by addition of NaHSO4 to decrease the pH or of
NaHCO3 to increase it). From the above equations for the solubility product, the
hydratation reaction and the two acid reactions, the following expression for
the maximum [Ca2+] can be easily deduced:
showing a quadratic dependence in . The numerical application with the above
values of the constants gives
pH 7.0 7.2 7.4 7.6 7.8 8.0 8.2 8.27 8.4
[Ca2+]max (10-4mol/L or °F) 1590 635 253 101 40.0 15.9 6.35 4.70 2.53
[Ca2+]max (mg/L) 6390 2540 1010 403 160 63.9 25.4 18.9 10.1
Comments:
* decreasing the pH from 8 to 7 increases the maximum Ca2+ concentration by a
factor 100
* note that the Ca2+ concentration of the previous table is recovered for pH =
8.27
* keeping the pH to 7.4 in a swimming pool (which gives optimum HClO/OCl- ratio
in the case of "chlorine" maintenance) results in a maximum Ca2+ concentration
of 1010 mg/L. This means that successive cycles of water evaporation and partial
renewing may result in a very hard water before CaCO3 precipitates. Addition of
a calcium sequestrant or complete renewing of the water will solve the problem.
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