Talk:Enthalpy of fusion
|WikiProject Physics||(Rated Start-class, Mid-importance)|
|WikiProject Chemistry||(Rated Start-class, High-importance)|
- 1 Conflict and redundancies
- 2 Calculations not right
- 3 How to quantify enthalpy of fusion at the atomic level?
- 4 no place for explaining differences with calories on food
- 5 Regarding the image
- 6 Conceptual difficulties
- 7 Definition and explanation
- 8 Please reformat the tables to per mol
- 9 Kinetic and potential energy
- 10 Applications
- 11 Enthalpy of fusion of water salt mixture?
- 12 Kinetic and potential energy (again)
Conflict and redundancies
¿Why are there separate articles for Enthalpy of fusion (and Enthalpy of vaporization) and Latent heat, when they describe essentially the same concept? To make matters worse, the data in the tables are different (108 kJ/kg for ethanol in the article on "Latent Heat" and 109 kJ/kg in the article about enthalpy). --Gonfer00 (talk) 09:10, 25 February 2012 (UTC)
Calculations not right
1 calorie = 4.184 J.
1 calorie/gram = 4.184 kJ/kg
Something is badly wrong with the table in the article.
How to quantify enthalpy of fusion at the atomic level?
There is a fundamental problem with the presentation of this page. The Enthalpy of fusion is the amount of energy that must be REMOVED from a substance for it to transition (fuse) from a liquid to a solid. It must have a negative value, as energy must be removed to achieve the fusion. The enthalpy of melting for the same substance is the exact same number, but is the positive number, as you need to add heat to transition from solid to liquid
I was trying to explain adding salt to ice, which causes the resulting water (assuming the ice is not so cold as to stay solid) to be colder than the original ice, in terms of energy transfer.
According to the formula that average kinetic energy ∝ T, the heat going into melting the ice doesn't increase the atoms' kinetic energy. Therefore, it seems the heat of fusion energy just increases potential energy w.r.t overcoming attractive forces. However, it seems strange that KE would stop increasing while potential energy continues to increase, instead of them increasing together continuously (think of a spring with a bouncing ball attached).
no place for explaining differences with calories on food
Somebody added an explanatory note in the first paragraph explaining the differences with food calories. I don't think this is the place to discuss something so unrelated to the thema. Also, if nobody oposses, I'll proceed to deletion / move to calorie page in two weeks.
Regarding the image
The image bothers me for numerous reasons of unprofessionalism, but most strikingly is the label of the horizontal axis that says "deg K." Units Kelvin are not measured in degrees. 188.8.131.52 (talk) 04:20, 20 January 2010 (UTC)
The text says:
When thermal energy is withdrawn from a liquid or solid, the temperature falls. ... However, at the transition point between solid and liquid (the melting point), extra energy is required (the heat of fusion). For freezing, the molecules of a substance arrange in an ordered state. For them to maintain the order of a solid, extra heat must be withdrawn. ... The temperature stops falling at (or just below) the freezing point, because energy from heat is stored in the form of primarily potential energy to build the solid lattice, rather than to increase the net kinetic energy (thermal energy) of the molecules. This introduced heat that does not result in a temperature change, is called a latent heat. The temperature of the system does not fall further until the liquid has frozen completely.
[OK, let me see if I’ve got this straight. The temperature of a liquid or solid falls because thermal energy is withdrawn.] The temperature stops falling at (or just below) the freezing point, because energy from heat is stored in the form of primarily potential energy to build the solid lattice,… [No problem so far.]
rather than to increase the net kinetic energy (thermal energy) of the molecules. [Wait! Thermal energy is being withdrawn. This withdrawal *increases* the net kinetic energy (thermal energy) of the molecules? Should this phrase be deleted?]
This introduced heat that does not result in a temperature change, is called a latent heat. [Wait! Thermal energy is being withdrawn. Why do we call this “introduced” heat? Shouldn’t we call this “retained” heat or “stored” heat—at least for the process of freezing?] ThatAMan (talk) 19:54, 28 February 2011 (UTC) — Preceding unsigned comment added by ThatAMan (talk • contribs) 21:57, 27 February 2011 (UTC)
Definition and explanation
I see two problems with this article as it stands. In the first para, there is confusion about the definition. The latent heat of fusion is the enthalpy change in a body when it melts. The specific heat of fusion is the enthalpy change per unit mass, and the molar heat of fusion is the enthalpy change per mole. This could be made much clearer.
In the third para, I agree with ThatAMan that the explanation is poorly worded and even misleading. The fundamental point here is that the liquid phase has a higher internal energy than the solid phase. This means energy must be supplied to a solid body in order to melt it and energy will be released from a liquid when it freezes. I propose replacing the third para with this:
When a body passes from a solid phase to a liquid phase, its molecules experience weaker intermolecular forces. Consequently the liquid phase contains a higher potential energy. This means that energy must be supplied to a solid body in order to melt it and energy will be released from a liquid when it freezes. In addition to the change in potential energy, there may also be a much smaller change in enthalpy caused by the fact the liquid and solid phases have different densities, so that some work has to be performed to change the phase.
This is still fairly thin as a thermodynamic explanation, but at least it is not misleading.
Please reformat the tables to per mol
Kinetic and potential energy
I have changed the reference to kinetic energy to potential energy. Molecules of liquids and solids at the melting point have the same kinetic energy, because they are at the same temperature. The difference in internal energy is one of potential energy, because of the weaker bonds in the liquid phase. Dezaxa (talk) 02:17, 12 October 2011 (UTC)
- Wouldn't their heat capacities be different? elle vécut heureuse à jamais (be free) 00:18, 14 October 2011 (UTC)
It took me 12 minutes to figure out what portion of ice would cool x units of room temperature water to freezing, surely one of the main reasons to look for the heat of fusion, so I added it. Calculation: 333.55 kJ to melt ice ~= (4 * 83.6 kJ) for 4kg of water down to 0 °C degrees from 20 °C MusicScienceGuy (talk) 20:39, 10 February 2012 (UTC)
Enthalpy of fusion of water salt mixture?
I have scoured the internet searching for any info on the enthalpy of fusion of water-salt ice, but havent been able to find anything. perhaps one of you knowledgable folks can direct me to where I can find this? or post it if you think its relevant? Dazalc (talk) 00:53, 16 June 2012 (UTC)
Kinetic and potential energy (again)
I've again changed back the reference from kinetic energy to potential energy. Molecules of water and ice at the same temperature have the same kinetic energy - this follows from the kinetic theory of heat, as long as we are not talking about very low temperatures where quantum effects are signigicant. Molecules of liquid water have a higher potential energy, because energy is needed to break the intermolecular bonds that are present in the solid phase. This is similar to a bond dissociation energy, but for intermolecular forces. It is a misconception to suppose that because bonds are stronger in the solid phase that this means the potential energy is higher; the opposite is true: bonds represent a reduction in free energy in comparison to free particles. Dezaxa (talk) 18:46, 28 October 2013 (UTC)