What is Vapor Liquid Equilibrium?

In chemistry and chemical engineering, vapor–liquid equilibrium (VLE) is the condition of a system in which the rates of evaporation (or vaporization) of a liquid from a surface and the condensation (or precipitation) of its vapor into the gas phase are equal. Under this condition, there is no net transfer of either phase. The concept is widely used in many branches of industrial chemistry, such as petroleum refining, where it finds use in distillation processes.

Vapor/Liquid Equilibrium Ratio Demonstration

Vapor-liquid equilibrium (VLE) is a state of dynamic equilibrium between the vapor phase and the liquid phase of a substance, in which both phases are in contact with each other. The term “equilibrium” here refers to a state of balance or stability in which the rates of vaporization and condensation are equal. When a substance is heated, its molecules gain energy and start to move around more rapidly.

At some point, the molecules will have enough energy to overcome the attractions that hold them together in the liquid state, and they will begin to vaporize. As more molecules escape from the liquid into the gas phase, the concentration of molecules in the gas phase increases. Meanwhile, as some molecules leave the liquid phase, others will be moving back into it from the gas phase.

This happens because not all molecules have enough energy to escape from the liquid; as they cool down, they lose energy and become attracted back into the liquid phase. The net result of these processes is that there is a constant exchange of molecules between the gas and liquid phases until an equilibrium is reached. At equilibrium, there is no net change in either phase; that is, neitherphase is gaining or losing molecules.

The concentrations of all componentsin both phases remain constant (although they may be different fromtheir initial concentrations). In addition,the temperature and pressureare also constant at equilibrium.

Vapor-Liquid Equilibrium Introduction

Vapor-liquid equilibrium (VLE) is a state where a liquid and its vapor are in equilibrium with each other, meaning that they are in contact with each other but neither is gaining or losing energy. This can happen when the two phases are at the same temperature and pressure. The most common type of VLE is boiling, which occurs when a liquid is heated to its boiling point and the resulting vapor is in equilibrium with the liquid.

Boiling is an example of a phase change, which is a change from one state of matter to another. The three states of matter are solid, liquid, and gas. A fourth state, plasma, can exist under certain conditions but will not be discussed here.

Solids have a definite shape and volume; liquids have a definite volume but take on the shape of their container; gases have neither a definite shape nor volume. When water freezes into ice, it changes from a liquid to a solid; when it boils and becomes steam, it changes from a liquid to a gas. In order for VLE to occur, there must be two things:

1) contact between the two phases (in this case, between theliquid water and gaseous steam), and 2) enough heat energy present so that the molecules can move around freely enough to attain equilibrium (i.e., so that they can boil). If either of these conditions is not met then VLE will not occur even if the temperature or pressure changes (as long as those changes do not also affect one of the other necessary conditions).

The reason why VLE only happens at specific temperatures and pressures has to do with how much energy each molecule has. At lower temperatures there isn’t enough thermal energy forthe moleculesto escape fromthe surfaceof aliquid intoitsgaseous form(thisiswhya potof wateronastovetopwill eventually come to aboil even thoughthe burneris turned off onceit reaches roomtemperature).At higher temperaturesmoleculeshaveenoughenergytoboilbut theyalsospeed upandmovearoundmore randomly;this makesthe likelihoodof them hittingeachotherand bouncing backinto theliquidphase ratherthan stayinginthegasphaseless likelyso thereis lessvaporproduced overall.A similarthinghappenswithpressure: increasingthepressureonthesurfaceoftheliquid increasesthenumberof collisionsbetweenmoleculesand thusreducesgasformation.

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Vapor-Liquid Equilibrium Examples

Vapor-liquid equilibrium (VLE) is a state where a chemical system consisting of two phases (vapor and liquid) are in mutual contact but in thermodynamic equilibrium with each other, meaning that there is no net change in the concentrations of the components in the two phases. The term can also refer to systems containing more than two phases. One example of vapor-liquid equilibrium is when water is heated at atmospheric pressure.

The temperature at which this occurs is called the boiling point. At the boiling point, the air above the water becomes saturated with water vapor and bubbles of vapor rise through the liquid. However, at this same temperature, the water below remains a liquid because it is in contact with air that does not contain enough vapor to saturate it.

This situation represents thermodynamic equilibrium between these two phases of matter. Another example VLE can be found when considering an alcoholic beverage such as beer or wine. When these drinks are left out exposed to air, they will eventually reach an equilibrium state where their alcohol content has partially evaporated into the surrounding air while also maintaining a constant concentration in the drink itself.

This happens because as alcohol molecules escape from the surface of the liquid they become more likely to encounter other alcohol molecules which then condense back into the liquid phase.

Vapor-Liquid Equilibrium in Distillation

Vapor-liquid equilibrium (VLE) is an important concept in distillation, as it determines the composition of the vapor and liquid phases in contact with each other. VLE data can be used to predict the behavior of mixtures during distillation and other separation processes. In a system at VLE, the vapor and liquid phases are in equilibrium with each other, meaning that they are in thermodynamic balance.

The composition of the two phases is also constant throughout the system. This relationship between the vapor and liquid phases is governed by their respective chemical potentials. The fugacity coefficient is a measure of how much a substance deviates from ideality.

It is defined as the ratio of the actual partial pressure of a substance to its partial pressure if it were ideal. For a pure substance, its fugacity coefficient is always equal to one. However, for mixtures, the fugacity coefficients can be different from each other and from one.

The activity coefficient (γ) is another measure of non-ideality which quantifies how much a solute departs from Raoult’s law behavior. In general, γ will be less than one for dilute solutions but greater than one for more concentrated solutions. The closer γ is to one, the more ideal the solution behaves.

For example, if γ = 0 . 5 {\displaystyle \gamma =0{.}5} , then only half as much solvent will evaporate as would be expected based on Raoult’s law (i . e . , 50% relative volatility).

The relation between fugacity coefficients and activity coefficients can be expressed mathematically as follows:

Vapor-Liquid Equilibrium Pdf

Vapor-liquid equilibrium (VLE) is a state of chemical thermodynamic equilibrium between vapor and liquid phases, in which both phases are in contact with each other. The term can also refer to the equilibrium condition achieved when a solution is heated at a constant temperature below its boiling point. The existence of two coexisting phases in thermodynamic equilibrium with each other implies that there is no net transfer of matter or energy between them.

In order for VLE to occur, all components must be present in both the vapor and liquid phases in the same proportions as they are in the mixture. If one component is present in only one phase, then VLE cannot occur. The pressure exerted by the vaporphase on the liquid phase is called the vapor pressure, and it increases with temperature according to the Clausius–Clapeyron relation:

\ln \left( \frac{P}{P_0} \right) = \frac{\Delta H_{vap}}{R} \left( \frac{1}{T} – \frac{1}{T_0}\right), where P is the absolute pressure, P_0 is a reference pressure (usually atmospheric pressure), T is absolute temperature, T_0is a reference temperature (usually 298.15 K or 25 °C), ΔH vapis molar enthalpy of vaporization, Ris universal gas constant (). This relation shows that at any given temperature above absolute zero, there will always be some pressure at which Vapor-Liquid Equilibrium will occur.

Vapor-Liquid Equilibrium Formula

In a binary mixture, the vapor-liquid equilibrium (VLE) occurs when the chemical potentials of the two components in the gas and liquid phases are equal. The VLE can be described by the following formula: X_g = X_l \frac{\mu _2 – \mu _1}{RT}

where Xg is the mole fraction of component 2 in the gas phase, Xl is the mole fraction of component 2 in the liquid phase, μ1 and μ2 are the chemical potentials of components 1 and 2 respectively, R is the universal gas constant, and T is absolute temperature. The VLE can be used to predict whether a given mixture will separate into phases or not. If the difference between μ1 and μ2 is positive, then component 2 will preferentially vaporize and a gas-liquid separation will occur.

If this difference is negative, then component 2 will preferentially stay in the liquid phase. If this difference is zero (i.e., if μ1=μ2), then there is no driving force for separation and equilibrium between phases exists.

What is Meant by Vapour Liquid Equilibrium?

Vapour-liquid equilibrium (VLE) is a state where the vapour and liquid phases of a substance are in equilibrium with each other, meaning that there is no net transfer of either phase. The pressure, temperature and composition of the two phases are all equal. The concept of VLE is important in many fields, including chemical engineering, pharmacy and thermodynamics.

It can be used to determine the optimum conditions for separating a mixture into its component parts, or for carrying out chemical reactions. When a substance is heated, it will first start to evaporate. This process will continue until the vapour and liquid phases reach equilibrium.

The temperature at which this happens is known as the boiling point. The composition of the vapour will be different from that of the liquid, because some components will vaporize more readily than others. The VLE curve shows how the composition of the vapour changes as the temperature is increased.

At lower temperatures, most of the substance will be in the liquid phase and only a small amount will be vaporized. As the temperature increases, more and more of the substance will vaporize until at very high temperatures almost all of it will be present in the vapour phase. To achieve VLE, both phases must be in contact with each other so that they can exchange molecules.

What is Liquid-Vapour Equilibrium Give Example?

In thermodynamics, liquid–vapour equilibrium or saturation equilibrium is the condition of a mixture of liquid and vapour of a single substance in contact with each other at a given temperature and pressure when thermal equilibrium has been reached. The chemical potentials of the liquid and vapour are equal at this temperature and pressure; thus there is no net flow of either phase into or out of the other. A familiar example is the case of water at its boiling point (100 °C, 1 atm).

If a pan containing only water (and not connected to any heat source) is left on a stovetop, eventually all the water will vaporize into steam, leaving no liquid water behind. However, if the pan is removed from the heat before all the water has vaporized, some steam will condense back into liquid water; both phases will be present in dynamic equilibrium. The same situation applies to any pure substance that can exist as both a liquid and a gas under specific conditions: for example ammonia (boiling point −33 °C), mercury (boiling point 357 °C) or chlorine (boiling point −34 °C).

With most substances that can exist as liquids or gases at room temperature however, such as air, nitrogen or carbon dioxide, very little condensation occurs because their vapours are much less dense than their liquids.

How Do You Calculate Vapour Liquid Equilibrium?

In order to calculate vapour liquid equilibrium, one must first understand the concept of partial pressures. The partial pressure of a given compound is the pressure that would be exerted by that compound if it were present alone in a container. When two or more gases are present in a container, their total pressure is equal to the sum of their partial pressures.

For example, consider a container with nitrogen and oxygen gas at a total pressure of 1 atmosphere (atm). If the partial pressure of nitrogen is 0.5 atm and the partial pressure of oxygen is 0.4 atm, then we know that the concentration of nitrogen molecules must be higher than the concentration of oxygen molecules. This is because when calculating vapour liquid equilibrium, we use mole fractions rather than concentrations – and the mole fraction of a given compound is simply its Partial Pressure divided by the Total Pressure.

In our example above, the mole fraction of nitrogen would be 0.5/1 = 0.5, while the mole fraction of oxygen would be 0.4/1 = 0.4 – meaning that there are more nitrogen molecules than oxygen molecules in our container despite both gases being present at equal pressures! Now that we understand how to calculate partial pressures, let’s move on to vapourliquid equilibrium itself. Vapourliquid equilibrium occurs when a system reaches chemical balance between its gaseous and liquid phases – meaning that there is no net change in either phase over time.

In order to calculate whether or not vapourliquid equilibrium has been reached, we need to compare two key properties: The Partial Pressure of each Compound in the Gas Phase – this represents how much each compound contributes tothe overall gas-phasepressure; and The Mole Fractionof each Compoundin boththe Gas Phaseand Liquid Phase – this represents how concentrated eachcompoundis within each phase respectively.

If we have these two values for each compound present in our system, we can then use themto calculate something calledthe Equilibrium Constant(Kp). This value tells us whetheror notvapourliquidequilibrium has been reached – if Kp>1 then vapour-rich conditions existand if Kp<1 then liquid-rich conditions exist;if Kpequals1 exactlythenvapouriquid equilibrium has been reachedfor all compoundsin threesystem!

What Happens When a Liquid And Vapour are in Equilibrium?

In a closed system, when a liquid and vapour are in equilibrium there is no net transfer of mass between the two phases. This means that the rate of condensation (liquid to vapour) is equal to the rate of evaporation (vapour to liquid). The temperature at which this occurs is called the boiling point.

At the boiling point, the vapour pressure of the liquid is equal to the atmospheric pressure. Below this temperature, the liquid can still evaporate, but it requires more energy to do so because its vapour pressure is lower than atmospheric pressure. Above the boiling point, the liquid can still condense, but again it requires more energy because now its vapour pressure is higher than atmospheric pressure.

Conclusion

Vapor liquid equilibrium (VLE) is a state where a liquid and its vapor are in thermodynamic equilibrium with each other. This means that the vapor pressure of the liquid is equal to the partial pressure of the vapor in the gas phase. The VLE can be used to determine the composition of a mixture of liquids and vapors.