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Deviation of real gas pressure is at not too high pressure and not too low temperature often negligible. … This has the opposite, pressure decreasing effect. …

The total work ( and its value for the constant force case ) for pressure $$p=p_\mathrm{force} + p_\mathrm{atm}$$ $$W_\mathrm{tot} = - \int_{V1}^{V2}{p \cdot \mathrm{d}V} $$ is shared between the source …

If we consider an idealised gas behaviour and the gas state equation $$pV = nRT$$ where $p$ is pressure $n$ is molar amount $R$ is universal gas constant $T$ is absolute temperature $V$ is gas volume … As the reaction is highly exothermic, temperature and therefore pressure significantly increase, as $$ p= \left( \frac {nR}{V} \right) T$$ …

Rather as "...the pressure necessary...", not force. The full speed osmosis means zero differential pressure on the membrane. … The particular value of the counteracting pressure, that would cause the zero net water flow, is called the osmotic pressure of the solution. …

then the amount of vapor and its pressure accomodate to external pressure. (**) If the external pressure is kept greater than saturated vapor pressure for given temperature, all vapour will eventually … gas+vapor mixtures, liquid saturated vapor pressure increases little with the total pressure. …

The gas with liquid vapour, previously in equilibrium with liquid, with suddenly increased pressure: Vapour gets over-saturated and condenses until the saturated vapour pressure is reached again. … Gas starts dissolving and gas partial pressure decreases, until the new equilibrium between its partial pressure and it's molar fraction in liquid is achieved ( Henry's law ). …

. (*) If air is being compressed, the partial vapor pressure increases. It it reaches the saturated vapor pressure at given temperature, water vapor starts to condensate. … and therefore vapor capacity of space very slighly increases with total pressure. …

}$, the total pressure raises to $\pu{10 atm}$, with partial pressures : $\ce{Ar} : \pu{5 atm}$ $\ce{N2} : \pu{4 atm}$ $\ce{O2} : \pu{1 atm}$ Note that pressure is macroscopic quantity. … Partial pressure of any gas is determined by frequency of collisions with a wall unit area, and by passed impulse by 1 molecule. …

At $\pu{25 ^{\circ}C}$, all water eventually evaporates, if pressure is kept below vapor pressure at $\pu{25 ^{\circ}C}$, or condenses if pressure is kept above the vapor pressure. … Water at room temperature would evaporate to ambient air, if water vapor partial pressure is lower than (saturated) vapor pressure, in spite of the total pressure being much higher than vapor pressure. …

Vapour pressure is intensive property. It depends for pure liquids (if we neglect minor secondary effects) only on temperature. …

The principle means that external change has smaller impact on the system ( in sense of the values of changed parameter, typically temperature or pressure, or change of composition), than it would have …

So to answer your question, you can get as high pressure as you can produce and the container can withstand, as there is no condensation reducing the pressure. …

Reducing pressure causes temperature drop, as part of thermal energy of air is spent on mechanical work during air expansion. That is why there is cold air at high altitudes. …

If you took out vapor at some moment and if you put it in a container of a fixed volume then its pressure would linearly grow with temperature, as in the Charles' law. … It grows roughly exponentially with temperature and so does the pressure. …

But the reaction quotient, (formally the same expression as for the equilibrium constant, but for any reaction state), decreases with increasing pressure, if volume of reactants is bigger than of products … changes of both temperature and pressure changes both values $K$ and $Q$. …