I'll give brief answers to your questions. If you need more detail, you should ask your questions separately.
What's the difference between heat and work at the atomic level? Isn't heat simply work between particles colliding with different momentum against each other?
Treating a substance semi-classically, one can say that at the atomic level, the atoms have a certain position and momentum. Quantum mechanically, even that's dubious because position and momentum are conjugate variables. With regard to heat and work, these don't exist at the atomic level.
Heat and work are processes, not states. Atoms don't contain heat or work. Neither do individual collections of atoms. Heat and work are measures of quantities transferred amongst objects. Objects don't contain heat or work.
Does an increase of pressure also increases the temperature of the gas?
For an ideal gas being compressed adiabatically, the answer is an emphatic yes. For anything else, the answer is sometimes yes, sometimes no. The answer depends on how much heat is being transferred into or out of the gas and on the nature of the gas. If the gas is right at the triple point (ideal gases don't have a triple point), all that compressing the gas adiabatically is going to do is cause some of the gas to turn into liquid or solid.
Excluding water and other special materials, why does a increase of pressure over a solid rises is melting point?
What your teacher told you is nonsense. Increased pressure does not decrease the molecule's motion. What increasing the pressure does do is to decrease the intermolecular distance.
The reason most substances contract when they freeze is because the bonding forces that make a substance become a crystalline solid hold the atoms/molecules closer together than the intermolecular distance at the same temperature in the liquid phase. Increasing the pressure in these substances decreases the intermolecular distance, thereby making it easier for those intermolecular bonding forces that make a substance a solid to take hold.
Water is different. It expands upon freezing. The structure of ice (ice Ih) is very open thanks to the hydrogen-hydrogen bonds in ice. Because ice expands upon freezing at normal pressures, increasing the pressure reduces the freezing point. Increase the pressure beyond about 100 atmospheres and water/ice starts behaving like most other substances. Increase the pressure beyond 3000 atmospheres and something even weirder happens. Now the freezing point drops markedly with increasing pressure. Increase the pressure beyond that and something even weirder happens: The freezing point increases again, this time very sharply increasing with rising pressure. The freezing point is over 600K at a pressure of 100,000 atmospheres.
If the pressure reduces the motion of the particles, how can the inner core have material with higher temperatures (i.e. particles with higher average kinetic energy)?
What your teacher told you was wrong.
Boiling will occur when the equilibrium vapor pressure at the temperature of the liquid is equal to the total pressure of the system and heat is being provided (either externally, or by the sensible heat of the liquid itself). The gas phase in contact with the liquid can be comprised of pure vapor (in which case the total pressure is essentially equal to the equilibrium vapor pressure), or, it can be comprised of a mixture of vapor and air (in which case the total pressure essentially equal to the sum of the equilibrium vapor pressure of the boiling substance plus the partial pressure of the air).
At 40 C, it is still possible for water to boil in a container if the gas is evacuated to a pressure below the equilibrium vapor pressure at 40 C (a non-equilibrium situation). Boiling can continue until the total pressure in the head space rises to a value equal to the equilibrium vapor pressure. As long as the partial pressure of the substance in the vapor and the total pressure (if air is present) is less than the equilibrium vapor pressure of water at the liquid temperature, boiling will continue. The heat of vaporization can be supplied by the liquid itself, so the water temperature will be dropping (assuming heat is not being directly supplied). Once the equilibrium vapor pressure at the liquid temperature falls below the total pressure, boiling will stop. Once the partial pressure of the vapor in the gas phase becomes equal to the equilibrium vapor pressure, evaporation will stop.
With regard to the question about the melting point and the triple point, the melting point is very close to the triple point. At the triple point, pure water is present in all three phases (no air present in the gas phase) and the total pressure is equal to the equilibrium vapor pressure of the liquid and solid. At the melting point, air is present in the gas phase at 1 atm., and the total pressure of the system is thus essentially 1 atm. The difference between the melting point and the triple point is only about 0.01 C.
Best Answer
The boiling point is the temperature at which the vapor pressure of the liquid equals the pressure at enviornment of liquid and the liquid changes to vapor.
A liquid in a vacuum has a lower boiling point than when that liquid is at atmospheric pressure. In other words, the boiling point of a liquid varies depending upon the surrounding environmental pressure. For a given pressure, different liquids boil at different temperatures The heat of vaporization is the energy required to transform a given quantity of a substance from a liquid into a gas at a given pressure. Liquids may change to a vapor at temperatures below their boiling points through the process of evaporation. Evaporation is a surface phenomenon in which molecules of liquid escape into the surroundings as vapor.
Even water will start boiling below 100 degree C when we reduce the pressure.and by furthur decrease in pressure it will start boiling at room temperature.the liquid uses its internal energy to change its phase.