Vapor Pressure Hydrogen

[Catherine Nicolo]

Incidentallythere are two established storage technologies for pure hydrogen in vehicles. One involves cooling the hydrogen below its critical temperature and liquifying it. The other stores it at ambient temperature, at pressures of $5\,000$ to $10\,000\ \mathrmpsi$. In the ambient temperature cases they are not storing compressed hydrogen gas, but rather compressed supercritical hydrogen.

Hydrogen can be stored at relatively high density (in some cases actually higher density than pure liquid hydrogen) at ambient pressure, by reversible formation of electropositive metal hydrides. One of the simplest such proposed storage compounds is [magnesium hydride]. Various Synthesis methods are reported by Wikipedia:

The first method with magnesium anthracene offers espcially good reversibility, as the magnesium anthracene can be formed under mild conditions from the metal and the anthracene recovered after reaction with hydrogen to form the hydride.

Hydrogen is easily ignited. Once ignited it burns with a pale blue, almost invisible flame. The vapors are lighter than air. It is flammable over a wide range of vapor/air concentrations. Hydrogen is not toxic but is a simple asphyxiate by the displacement of oxygen in the air. Under prolonged exposure to fire or intense heat the containers may rupture violently and rocket.

The hydrogen phase diagram shows the phase behavior with changes in temperature and pressure. The curve between the critical point and the triple point shows the hydrogen boiling point with changes in pressure. It also shows the saturation pressure with changes in temperature.

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In a mixture of gases, each constituent gas has a partial pressure which is the notional pressure of that constituent gas as if it alone occupied the entire volume of the original mixture at the same temperature.[1] The total pressure of an ideal gas mixture is the sum of the partial pressures of the gases in the mixture (Dalton's Law).

The partial pressure of a gas is a measure of thermodynamic activity of the gas's molecules. Gases dissolve, diffuse, and react according to their partial pressures but not according to their concentrations in gas mixtures or liquids. This general property of gases is also true in chemical reactions of gases in biology. For example, the necessary amount of oxygen for human respiration, and the amount that is toxic, is set by the partial pressure of oxygen alone. This is true across a very wide range of different concentrations of oxygen present in various inhaled breathing gases or dissolved in blood;[2] consequently, mixture ratios, like that of breathable 20% oxygen and 80% Nitrogen, are determined by volume instead of by weight or mass.[3] Furthermore, the partial pressures of oxygen and carbon dioxide are important parameters in tests of arterial blood gases. That said, these pressures can also be measured in, for example, cerebrospinal fluid.

The symbol for pressure is usually p or pp which may use a subscript to identify the pressure, and gas species are also referred to by subscript. When combined, these subscripts are applied recursively.[4][5]

Dalton's law expresses the fact that the total pressure of a mixture of ideal gases is equal to the sum of the partial pressures of the individual gases in the mixture.[6] This equality arises from the fact that in an ideal gas, the molecules are so far apart that they do not interact with each other. Most actual real-world gases come very close to this ideal. For example, given an ideal gas mixture of nitrogen (N2), hydrogen (H2) and ammonia (NH3):

Vapor pressure is the pressure of a vapor in equilibrium with its non-vapor phases (i.e., liquid or solid). Most often the term is used to describe a liquid's tendency to evaporate. It is a measure of the tendency of molecules and atoms to escape from a liquid or a solid. A liquid's atmospheric pressure boiling point corresponds to the temperature at which its vapor pressure is equal to the surrounding atmospheric pressure and it is often called the normal boiling point.

The vapor pressure chart displayed has graphs of the vapor pressures versus temperatures for a variety of liquids.[9] As can be seen in the chart, the liquids with the highest vapor pressures have the lowest normal boiling points.

For reversible reactions, changes in the total pressure, temperature or reactant concentrations will shift the equilibrium so as to favor either the right or left side of the reaction in accordance with Le Chatelier's Principle. However, the reaction kinetics may either oppose or enhance the equilibrium shift. In some cases, the reaction kinetics may be the overriding factor to consider.

Gases will dissolve in liquids to an extent that is determined by the equilibrium between the undissolved gas and the gas that has dissolved in the liquid (called the solvent).[10] The equilibrium constant for that equilibrium is:

The form of the equilibrium constant shows that the concentration of a solute gas in a solution is directly proportional to the partial pressure of that gas above the solution. This statement is known as Henry's law and the equilibrium constant k \displaystyle k is quite often referred to as the Henry's law constant.[10][11][12]

The minimum safe lower limit for the partial pressures of oxygen in a breathing gas mixture for diving is 0.16 bars (16 kPa) absolute. Hypoxia and sudden unconsciousness can become a problem with an oxygen partial pressure of less than 0.16 bar absolute.[15] Oxygen toxicity, involving convulsions, becomes a problem when oxygen partial pressure is too high. The NOAA Diving Manual recommends a maximum single exposure of 45 minutes at 1.6 bar absolute, of 120 minutes at 1.5 bar absolute, of 150 minutes at 1.4 bar absolute, of 180 minutes at 1.3 bar absolute and of 210 minutes at 1.2 bar absolute. Oxygen toxicity becomes a risk when these oxygen partial pressures and exposures are exceeded. The partial pressure of oxygen also determines the maximum operating depth of a gas mixture.[14]

Narcosis is a problem when breathing gases at high pressure. Typically, the maximum total partial pressure of narcotic gases used when planning for technical diving may be around 4.5 bar absolute, based on an equivalent narcotic depth of 35 metres (115 ft).

The effect of a toxic contaminant such as carbon monoxide in breathing gas is also related to the partial pressure when breathed. A mixture which may be relatively safe at the surface could be dangerously toxic at the maximum depth of a dive, or a tolerable level of carbon dioxide in the breathing loop of a diving rebreather may become intolerable within seconds during descent when the partial pressure rapidly increases, and could lead to panic or incapacitation of the diver.[14]

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Many chemicals used in the industrial field present risks, which differ depending on their chemical properties. Additionally, their various physicochemical properties change considerably with concentration. Many chemicals are used in customized processes in factories in the form of different aqueous solutions. The Korean Chemicals Control Act evaluates "hazardous chemicals," describes their risks to the public, and regulates their concentration. To prepare against chemical accidents, factories construct models of potential damage radius, which is greatly influenced by a chemical's vapor pressure. This study selected substances with widely varying vapor pressures (hydrogen fluoride, hydrogen chloride, aqueous ammonia, and hydrogen peroxide) and compared the results of different modeling programs (KORA, ALOHA, PHAST, and RMP*Comp) for various aqueous solution concentrations. The results showed that damage radius and vapor pressure increased similarly for each substance. Damage radius was negligible at low concentrations for all substances studied. Damage radius of ammonia solution increased with vapor pressure. Hydrogen fluoride is not found in aqueous solution at concentrations of less than 37%, and hydrogen peroxide does not show a large damage radius at low concentrations. However, the Chemicals Control Act strictly regulates hydrogen fluoride concentration beginning at 1%, hydrogen chloride and aqueous ammonia at 10%, and hydrogen peroxide at 6%. To effectively prepare against chemical accidents, we must examine scientifically-based, suitable regulations based on physicochemical properties.

In many cases, the amount of gas evolved by a reaction is ofinterest. Since gases have such small densities, it is usuallynot practical to collect the gas and find its mass. For gasesthat are not particularly soluble in water, it is possible tocollect the evolved gas by displacement of water from acontainer.

The setup for the collection of a gas over water involves acontainer in which the reaction takes place and a gas collectioncontainer filled with water and inverted in a reservoir of water.The gas evolved from the reaction is collected by attaching oneend of a hose to the reaction container and inserting the other upinto the inverted gas collection bottle. As the gas is created,it will displace water from the bottle. The volume of gas can bedetermined by the amount of water that was displaced by thegas.

During the collection, the water level in the container willadjust so that the pressure inside and outside the container arethe same. Because of this, if we know the atmoshperic pressure,we also know the pressure of the gas inside the bottle.

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