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11.8:

Dampdruk

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Chemistry
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JoVE Core Chemistry
Vapor Pressure

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In een gesloten systeem, bij damp-vloeistofevenwicht, vinden condensatie en verdamping plaats met dezelfde snelheid, zonder netto verandering in de massa’s van de twee fasen. De partiële druk die wordt uitgeoefend door de gasfase in dynamisch evenwicht met de vloeistof wordt de dampspanning genoemd. Hoe meer moleculen er in de dampfase zijn, hoe hoger de dampspanning zal zijn.De dampspanning is dus een weerspiegeling van de neiging van vloeistofmoleculen om bij een bepaalde temperatuur in de dampfase te ontsnappen. Het is een meetbare hoeveelheid, beheerst door intermoleculaire krachten. Vluchtigheid beschrijft deze neiging kwalitatief, gebaseerd op de dampdrukken van vloeistoffen die onder dezelfde omstandigheden worden gehouden.Vergelijk bijvoorbeeld hexaan en water die op dezelfde temperatuur worden gehouden. Omdat hexaan zwakke dispersiekrachten vertoont en water sterke waterstofbruggen vertoont, verdampt hexaan gemakkelijker dan water. In een gesloten systeem bij evenwicht heeft hexaan een hogere dampspanning dan water:hexaan is vluchtig, terwijl water niet vluchtig is.De verdeling van thermische energieën in de vloeistoffase is een functie van temperatuur. Het verwarmen van een vloeistof verhoogt de temperatuur, wat aangeeft dat de moleculen hogere thermische energieën hebben, wat leidt tot een hogere verdampingssnelheid en hogere dampdruk. Wanneer de dampdruk gelijk is aan de externe druk, begint de vloeistof te koken en de temperatuur waarbij dit gebeurt, wordt het kookpunt van de vloeistof genoemd.Het normale kookpunt van een vloeistof is de temperatuur waarbij de dampdruk van de vloeistof gelijk is aan 1 atmosfeer. Bij een andere externe druk kookt de vloeistof echter bij een andere temperatuur dan het normale kookpunt. Op standaard zeeniveau, waar de atmosferische druk 1 atmosfeer is, kookt water bijvoorbeeld bij 100 graden Celsius.Op grotere hoogte waar de atmosferische druk minder is dan 1 atmosfeer vereist de dampfase minder moleculen om de lagere externe druk te evenaren. Dit verklaart waarom water op een lagere temperatuur kookt. In een snelkookpan vereist de hogere externe druk meer dampfasemoleculen, waardoor het water een hogere temperatuur moet hebben om te koken.

11.8:

Dampdruk

When a liquid vaporizes in a closed container, gas molecules cannot escape. As these gas phase molecules move randomly about, they will occasionally collide with the surface of the condensed phase, and in some cases, these collisions will result in the molecules re-entering the condensed phase. The change from the gas phase to the liquid is called condensation. When the rate of condensation becomes equal to the rate of vaporization, neither the amount of the liquid nor the amount of the vapor in the container changes. The vapor in the container is then said to be in equilibrium with the liquid. Keep in mind that this is not a static situation, as molecules are continually exchanged between the condensed and gaseous phases. Such is an example of dynamic equilibrium, the status of a system in which reciprocal processes (for example, vaporization and condensation) occur at equal rates.

The pressure exerted by the vapor in equilibrium with a liquid in a closed container at a given temperature is called the liquid’s vapor pressure (or equilibrium vapor pressure). The area of the surface of the liquid in contact with a vapor and the size of the vessel have no effect on the vapor pressure, although they do affect the time required for the equilibrium to be reached. The chemical identities of the molecules in a liquid determine the types (and strengths) of intermolecular attractions possible; consequently, different substances will exhibit different equilibrium vapor pressures. Relatively strong intermolecular attractive forces will serve to impede vaporization as well as favoring “recapture” of gas-phase molecules when they collide with the liquid surface, resulting in a relatively low vapor pressure. Weak intermolecular attractions present less of a barrier to vaporization, and a reduced likelihood of gas recapture, yielding relatively high vapor pressures.

Consider four compounds: ethanol (CH3CH2OH), ethylene glycol (C2H6O2), diethyl ether (C4H10O), and water (H2O).

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Diethyl ether has a very small dipole, and most of its intermolecular attractions are London dispersion forces. Although this molecule is the largest of the four under consideration, its IMFs are the weakest and, as a result, its molecules most readily escape from the liquid. It also has the highest vapor pressure. Due to its smaller size, ethanol exhibits weaker dispersion forces than diethyl ether. However, ethanol is capable of hydrogen bonding and, therefore, exhibits stronger overall IMFs, which means that fewer molecules escape from the liquid at any given temperature, and so ethanol has a lower vapor pressure than diethyl ether. Water is much smaller than either of the previous substances and exhibits weaker dispersion forces, but its extensive hydrogen bonding provides stronger intermolecular attractions, fewer molecules escaping the liquid, and a lower vapor pressure than for either diethyl ether or ethanol. Ethylene glycol has two −OH groups, so, like water, it exhibits extensive hydrogen bonding. It is much larger than water and thus experiences larger London forces. Its overall IMFs are the largest of these four substances, which means its vaporization rate will be the slowest and, consequently, its vapor pressure the lowest.

As temperature increases, the vapor pressure of a liquid also increases due to the increased average KE of its molecules. Recall that at any given temperature, the molecules of a substance experience a range of kinetic energies, with a certain fraction of molecules having sufficient energy to overcome IMF and escape the liquid (vaporize). At a higher temperature, a greater fraction of molecules have enough energy to escape from the liquid. The escape of more molecules per unit of time and the greater average speed of the molecules that escape both contribute to the higher vapor pressure.

When the vapor pressure increases enough to equal the external atmospheric pressure, the liquid reaches its boiling point. The boiling point of a liquid is the temperature at which its equilibrium vapor pressure is equal to the pressure exerted on the liquid by its gaseous surroundings. For liquids in open containers, this pressure is that due to the earth’s atmosphere. The normal boiling point of a liquid is defined as its boiling point when surrounding pressure is equal to 1 atm (101.3 kPa). At pressures greater than 1 atm, the boiling point of the liquid is higher than its normal boiling point.

This text is adapted from Openstax, Chemistry 2e, Section 10.3: Phase Transitions.