Evaporation of substances

[Reven]

Hi

in a vacuum water would sublimate when heated up without ever going through the liquid phase first. As I understand it intuitively, this is because up to melting temperature, the heat energy goes into separating the molecules from their bonds, but it is the air pressure that then prevent them from just flying off. which means they start swimming among each other without yet being able to escape the downward push force of the atmospheric air pressure.

So adding enough heat to increase the temperature by 100 deg C, is then what is needed to give the molecules enough energy to overcome the air pressure and turn into gas (boil. I know some turn into gas before that temp.). Is this understanding correct?

If it is correct. My main question is this; if water is able to fight atm. Air pressure with only a 100 deg C difference in temp/energy. Why does a metal for instance need a temperature/energy difference from 1500 deg C (melting) to 2500 deg C (boil) in order to counter the same atm. air pressure and become a gas?

[scooke]

Dear Ronny,

We’re always glad to see curious people exploring chemistry! We do try to give relatively short explanations for some questions. Others may really be a good springboard to further study. Your question takes a commonly observable phenomenon that has a more complex chemical explanation.

The short, ‘easy’ version is that the energy holding different types of molecules and atoms together is different for all different compounds. Thus, the energy required to get them to change phase is different. So, if you look at the energy in terms of temperature (thermal energy), the hotter something becomes, the more rapidly the individual atoms of the compound move.

When that movement is sufficiently large it can “break away” from the inter-atomic attractions that make a substance solid or liquid in bulk form, and “change” the phase to liquid or vapor. A metal has much stronger internal attractions (not always a true ‘bond’) than most organics or water. Thus, it requires much higher temperatures (more energy) to get them to separate, forming a liquid from a solid phase, or evaporating as a vapor from a liquid phase.

The pressure dependence is really just an additional restrictive force on the fundamental motion of the atoms. Besides the internal attractions that must be exceeded, any atmospheric pressure must also be overcome in order for a compound to leave a bulk liquid and enter the surrounding vapor atmosphere. That is why things boil at lower temperatures under reduced pressure or vacuums, and will require higher temperatures under higher pressures (the principle behind a pressure-cooker).

You can look up other discussions on the effect of pressure, temperature and chemical structure on freezing/melting and condensation/boiling points of different substances. You may be able to try some experiments of your own with common compounds and simple equipment. I would offer that one of the reasons we do have these references, as well as formal programs of study, is that none of us has the time to recreate all of the previous experiments that provided the wealth of chemical knowledge that we have available today.

So, STAY curious. Keep asking questions. And look for good sources of information on the topics that interest you. In all cases, be very careful about any SAFETY considerations when dealing with different compounds or energetic materials (including flames)

Best regards,

Steven

[karenwiz]

Just wanted to some comments to Steven's excellent answer.

First of all I'd like to say that I really like your image of heated molecules "swimming" among each other as their intermolecular bonds are weakened .  Gives an idea of the motion of molecules past each other that I never visualized before!

You did phrase something a little awkwardly: "separating the molecules from their bonds" - molecules are separated from other molecules by breaking or weakening the bonds between them - you don't separate molecules from bonds. 

Intermolecular bonds are vastly weaker than the intramolecular bonds that hold atoms together in the molecule but are also really interesting because of how they influence things like melting and boiling (as you discussed) but also other characteristics like surface tension and crystallization. 

You have used the example of water in your thought experiment.  Water is one of the most amazing chemicals because of its extremely strong hydrogen bonding which is an intermolecular bond.  Most molecules of the size of water (think for example about carbon dioxide or hydrogen sulfide) boil at such low temperatures (-78 C and -60 C respectively) that they are gases at room temperature! These compounds have lower intermolecular bonds, and are also of low mass which makes it easy for there to be enough ambient energy at ordinary conditions for them to be gases. (This is the opposite example to yours of the metals where the interatomic or intermolecular attractions are very high.) Water is kind of an anomalous compound because, despite being such a small molecule, it is a liquid at 25 C and 1 atm air pressure. This is due to the strength of its intermolecular bonding.

Lastly I'd like to second Steven's comments about you staying curious and continuing to think about how things work. Your diligence in actively seeking answers to these questions you have is to be commended.

Karen

[hlmark]

I think both Steve and Karen missed an important point, due to the way Ronny stated her question. What they say about intra- versus inter- molecular forces is correct, but I think Ronny sowed confusion in the way she phrased her question. The important point that was missed is whether the gas above the water in question consists of pure water (i.e., water vapor or steam), or a mixture of water and air (or other gas).

The simpler case is when the gas above the liquid water consists only of water vapor (also known as steam). Assuming that the system is in equilibrium and at steady-state (i.e., water is not actively evaporating or condensing), then at constant volume the pressure of the system is determined only by  the temperature of the water (which is the same as that of of the water vapor). It doesn't matter how much vapor or liquid water there is, as long as both are present and equilibrium is maintained. If heat is added so that the  temperature rises, then either the pressure or the volume of steam (or both) will increase, until equilibrium is attained at the new conditions. The full behavior of the liquid water + steam is described by what's known as a Phase Diagram.

This discussion assumes that only liquid water and steam are present , and does not consider any other effects, such as the temperature dropping so low that the water freezes, for example. Then things get much more complicated, the same as happens if the gas above the liquid water consists of a mixture of water vapor + air.  A full phase diagram for water includes the behavior for mixtures of all three phases: gas, liquid & solid.

Ronny's confusion comes in when she asks about the effect of atmospheric air pressure. That implies that there IS external atmospheric air pressure. That brings in a whole set of other effects beyond what pure water (in different phases) exhibits. I'll try to simplify it here. As long as the liquid water is between 0 and 100 degrees Celsius then some water molecules will leave the surface and enter the gas above it (which consists of a mixture of water + air). The higher the temperature of the water, the more water molecules evaporate and enter the gas phase. This effect is known as the vapor pressure of the water. The higher the temperature, the higher the vapor pressure. When the vapor pressure equals the external atmospheric pressure (which occurs when the temperature  reaches 100 degrees C, at normal atmospheric pressure) then the molecules can't leave the surface of the water fast enough,  and the water boils.

[hlmark]

I think both Steve and Karen missed an important point, due to the way Ronny stated her question. What they say about intra- versus inter- molecular forces is correct, but I think Ronny sowed confusion in the way she phrased her question. The important point that was missed is whether the gas above the water in question consists of pure water (i.e., water vapor or steam), or a mixture of water and air (or other gas).

The simpler case is when the gas above the liquid water consists only of water vapor (also known as steam). Assuming that the system is in equilibrium and at steady-state (i.e., water is not actively evaporating or condensing), then at constant volume the pressure of the system is determined only by  the temperature of the water (which is the same as that of of the water vapor). It doesn't matter how much vapor or liquid water there is, as long as both are present and equilibrium is maintained. If heat is added so that the  temperature rises, then either the pressure or the volume of steam (or both) will increase, until equilibrium is attained at the new conditions. The full behavior of the liquid water + steam is described by what's known as a Phase Diagram.

This discussion assumes that only liquid water and steam are present , and does not consider any other effects, such as the temperature dropping so low that the water freezes, for example. Then things get much more complicated, the same as happens if the gas above the liquid water consists of a mixture of water vapor + air.  A full phase diagram for water includes the behavior for mixtures of all three phases: gas, liquid & solid.

Ronny's confusion comes in when she asks about the effect of atmospheric air pressure. That implies that there IS external atmospheric air pressure. That brings in a whole set of other effects beyond what pure water (in different phases) exhibits. I'll try to simplify it here. As long as the liquid water is between 0 and 100 degrees Celsius then some water molecules will leave the surface and enter the gas above it (which consists of a mixture of water + air). The higher the temperature of the water, the more water molecules evaporate and enter the gas phase. This effect is known as the vapor pressure of the water. The higher the temperature, the higher the vapor pressure. When the vapor pressure equals the external atmospheric pressure (which occurs when the temperature  reaches 100 degrees C, at normal atmospheric pressure) then the molecules can't leave the surface of the water fast enough,  and the water boils.

[jbozin]

Something is being missed here, without reading all the comments.  

1. To "sublime", the water has to be in the solid state(aka ice). 

2. Since phase change is a state function, one has to put all the energy to melt, heat-up, and evap to turn it to pure vapor from the solid, it's just occurring at a lower temperature...  

The system pressure affects the temperature at which it will occur.   In this case, in vacuo.

3. The air is not applying pressure on the water and holding it back, as described, the water and air are both exerting their own pressure.  Ptotal= Sum {Pvp}i  (chemical editing not good on this site.)   So when one removes the air pressure, via vacuum, only the water exerts vapor pressure.   Vapor pressure, Pvp, is a function of temperature = f(T).

I know it's confusing for some, but pick up a good chemical thermodynamics book and it's all spelled out in there.

4. As for metals, they have much higher molecular weight, so require much more energy to turn to vapor.   Therefore, the melt points are very high, and the boiling points even greater.   You can sublimate some metals in processes known as vapor deposition.   Requires high vacuum and plasma temperatures, or precursor molecules that decompose.

Sorry to say, the air is not playing the role as described in the question.