Special engineering consideration shall be required if the specific gravity of the liquid to be stored exceeds that of water or if the tanks are designed to contain flammable liquids at a liquid temperature below 0 F.
I always thought that pure water cannot exceed 100° Celsius at atmospheric p…
Even if you don't apply pressure, you can still have liquid water at sub-zero temperatures using additives. Additives such as salt can interfere with the chemical bonding needed to form a solid and they therefore can lower water's freezing point. Salt is composed of strong sodium and chlorine ions. When dissolved in water, the water molecules tend to stick to the salt ions instead of to each other, and they therefore don't freeze as readily. As you add more salt to water, its freezing point continues to drop until the water reaches saturation and cannot hold any more salt. If you add enough salt, the freezing point of water can be dropped as low as -21 degrees Celsius. This fact means that water at -21 degrees Celsius can still remain liquid if enough salt is added. Instead of keeping liquid water from freezing, this powerful property of salt can also be used to turn ice back into water. Sprinkling salt on icy sidewalks lowers the freezing point of the ice below the ambient temperature and the ice melts. But sprinkling salt on icy walkways won't help if the ambient temperature is below -21 degrees Celsius. The effect of salt on water's freezing point also has profound effects on earth's oceans.
Even if you don't apply pressure and don't add anything to the water, you can still have liquid water at temperatures below zero degrees Celsius. In order for water to freeze to ice, it needs something to freeze onto to start the process. We call these starting points "nucleation centers". In most situations, a little bit of dust, impurity, or even little vibrations in the water provide nucleation centers for the water to freeze onto. But if your water is very pure and very still, there is nothing for the water molecules to crystallize onto. As a result, you can cool very pure water well below zero degrees Celsius without it freezing. Water in this condition is called "supercooled". At standard pressure, pure water can be supercooled to as low as about -40 degrees Celsius. Supercooled water is kept from freezing only by the lack of nucleation centers. Therefore, once nucleation centers are provided (which could be as simple as a vibration), the supercooled water quickly freezes. Freezing rain is a natural example of supercooled liquid water. Once freezing rain hits an object on earth's surface, that object provides nucleation centers, and the rain freezes to ice.
Meteorologists routinely consider the "dewpoint" temperature (instead of, but analogous to absolute humidity) to evaluate moisture, especially in the spring and summer. The dewpoint temperature, which provides a measure of the actual amount of water vapor in the air, is the temperature to which the air must be cooled in order for that air to be saturated. Although weather conditions affect people differently, in general in the spring and summer, surface dewpoint temperatures in the 50s usually are comfortable to most people, in the 60s are somewhat uncomfortable (humid), and in the 70s are quite uncomfortable (very humid). In the Ohio Valley (including Kentucky), common dewpoints during the summer range from the middle 60s to middle 70s. Dewpoints as high as 80 or the lower 80s have been recorded, which is very oppressive but fortunately relatively rare. While dewpoint gives one a quick idea of moisture content in the air, relative humidity does not since the humidity is relative to the air temperature. In other words, relative humidity cannot be determined from knowing the dewpoint alone, the actual air temperature must also be known. If the air is totally saturated at a particular level (e.g., the surface), then the dewpoint temperature is the same as the actual air temperature, and the relative humidity is 100 percent.
Boiling is the most reliable method the public can use to disinfect their drinking water and should be the first option for on-site disinfection. However, it may not always be possible or practical to boil water. Power outages may leave consumers unable to boil, and boiling may not be practical to meet some water needs. If needs are critical and cannot be discontinued, alternate water sources or other disinfection methods may be necessary. Generally, water used by the public for drinking and food preparation during a boil water event should be obtained in the following order of preference, depending on the scope of the affected area and incident specific conditions:
Most labs use at least one type of heating device, such as ovens, hot plates, heating mantles and tapes, oil baths, salt baths, sand baths, air baths, hot-tube furnaces, hot-air guns and microwave ovens. Steam-heated devices are generally preferred whenever temperatures of 100o C or less are required because they do not present shock or spark risks and can be left unattended with assurance that their temperature will never exceed 100o C. Ensure the supply of water for steam generation is sufficient prior to leaving the reaction for any extended period of time.
Fail-safe devices can prevent fires or explosions that may arise if the temperature of a reaction increases significantly because of a change in line voltage, the accidental loss of reaction solvent or loss of cooling. Some devices will turn off the electric power if the temperature of the heating device exceeds some preset limit or if the flow of cooling water through a condenser is stopped owing to the loss of water pressure or loosening of the water supply hose to a condenser.
Pressure cookers are equipped with a valve that lets gas escape when the pressureinside the pot exceeds some fixed value. This valve is often set at 15 psi, which meansthat the water vapor inside the pot must reach a pressure of 2 atm before it can escape.Because water doesn't reach a vapor pressure of 2 atm until the temperature is 120oC,it boils in this container at 120oC.
Based on air bubbles trapped in mile-thick ice cores and other paleoclimate evidence, we know that during the ice age cycles of the past million years or so, atmospheric carbon dioxide never exceeded 300 ppm. Before the Industrial Revolution started in the mid-1700s, atmospheric carbon dioxide was 280 ppm or less.
Although abundant on earth as an element, hydrogen is almost always found as part of another compound, such as water (H2O) or methane (CH4), and it must be separated into pure hydrogen (H2) for use in fuel cell electric vehicles. Hydrogen fuel combines with oxygen from the air through a fuel cell, creating electricity and water through an electrochemical process.
"The water that we generate is much cleaner than anything you'll ever get out of any tap in the United States," says Carter. "We certainly do a much more aggressive treatment process (than municipal waste water treatment plants). We have practically ultra-pure water by the time our water's finished."
Despite Mars' thin atmosphere, the Red Planet still exhibits a dynamic climate and extreme weather events including impressive dust storms and even snow! But Mars hasn't always been this way. NASA's MAVEN mission scientists reported that Mars once had a thick atmosphere (opens in new tab) that could have supported surface liquid water on the surface for extended periods of time.
At times, it even snows on Mars. The Martian snowflakes, made of carbon dioxide rather than water, are thought to be very small particles that create a fog effect rather than appearing as falling snow. The north and south polar regions of Mars are capped by ice, much of it made from carbon dioxide, not water.
Things are more flammable in pure oxygen than in the 21% oxygen of the normal atmosphere (seehere for a demonstration). Even iron, which does not burn in 21% oxygen, will burn in 100% oxygen (seehere). This led to a disaster in the early days of the American space program. The capsules used in the Mercury, Gemini, and Apollo program operated in space with an atmosphere of pure oxygen at a pressure of 5 psi. (Normal atmospheric pressure is about 14.7 psi. The reduced pressure oxygen environment was used to eliminate the need to carry nitrogen tanks into space for the astronauts to breathe, in the belief that a two-component gas system would be more difficult to manage, since it would require the oxygen/nitrogen ratio to be calibrated precisely at all times to prevent the astronauts from suffocating.) Since the spacecraft were to operate in a pure oxygen environment in space, they were tested on the ground in a pure oxygen environment. There were no complications with these tests during the Mercury and Gemini programs of the early 1960s, but on January 27, 1967, a fire broke out during a routine test in the command module of what was to become the first mission of the Apollo program; the door opened inward, and rapidly become impossible to pull open against the pressure of the gases being generated in the fire. Within 17 seconds, the astronauts Virgil Grissom, Ed White, and Roger Chaffee were killed. As a result of the disaster, the Command Module was extensively redesigned to prevent such a tragedy from occurring again.
The discovery of oxygen is an extremely tangled story, partially because of questions of priority, and partially because of misunderstandings about the nature of combustion and the gas phase. For thousands of years, air was considered to be an "element," and it was not recognized that air was actually a mixture of many different gases. The nature of combustion was also hotly debated (pun intended); many scientists believed that flammable substances contained a material called phlogiston, which was released when a substance burned. When nitrogen was discovered in 1772, it was referred to as "phlogisticated air," since an atmosphere of pure nitrogen (actually, nitrogen plus carbon dioxide) did not support combustion. (It was thought that this "air" had absorbed the maximum amount of phlogiston.) Oxygen was discovered by the Swedish chemist Carl Wilhelm Scheele in 1772, but his account of his experiment was not published until 1777. The English chemist Joseph Priestley produced oxygen in 1774 by heating a sample of merucry(II) oxide, HgO, and collecting the oxygen gas it produced over water. He called the gas "dephlogisticated air," since it supported combustion more vigorously that "normal" air, and therefore presumably was more capable of "pulling" phlogiston out of other substances. The French chemist Antoine Lavoisier claimed to have produced oxygen in 1774, independently of Priestley, but Priestley had visited him a few months before and told him of his experiment. Lavoisier did, however, correctly interpret the significance of Priestley's result: that combustion is the not release of phlogiston from a substance, but the combination of the substance with oxygen in the air, to produce oxides (as well as heat and light). Lavoisier believed that the new element was an essential component of all acids, and proposed that it be called "oxygen," from the Greek words oxy, "acid" and genes "forming." (However, not all acids contain oxygen; for example, hydrochloric acid, HCl.) 2ff7e9595c
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