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Because hydrogen bonds are rather strong intermolecular interactions, breaking them requires a great deal of energy. This is why water has a high boiling point: breaking the intermolecular interactions that hold water molecules together requires a lot of heat (energy).
The melting and boiling temperatures of compounds with hydrogen bonds are unusually high. The fact that more energy is required to break hydrogen bonds accounts for the compound’s high melting and boiling points.
The existence of hydrogen bonding accounts for the particularly high boiling point of hydrogen fluoride among the halogen acids.
ii. At ordinary temperatures, H2O is a liquid, while H2S, H2Se, and H2Te are all gases.
Water’s boiling point is higher than that of other chemicals due to hydrogen bonding, which generates links in the water molecules.
iii. Ammonia has a greater boiling point than PH3 due to hydrogen bonding, which does not exist in PH3.
Because there are hydrogen bonds in ethanol, it has a higher boiling point than diethyl ether.
The unusually high boiling point of hydrogen fluoride
The electronegativity of fluorine is the highest of any element. That it enjoys attracting electrons. When fluorine interacts with hydrogen, its polarity becomes so strong that it exhibits hydrogen bonding, which is essentially just an extreme dipole. Chlorine lacks the attraction of its halogen equivalent and is incapable of forming a polar molecule.
Because of hydrogen bonding, hydrogen fluoride molecules tend to stick together in practice, requiring more energy to separate them and transition them from a liquid to a gas. Because hydrogen chloride molecules do not have hydrogen bonds, it is easier to separate them.
Hydrogen bonds can be formed between HF molecules due to the high electronegativity of fluorine.
– It takes more energy to break hydrogen bonds than it does to shatter London Forces. Other halogens are less electronegative, thus they can’t form hydrogen bonds with other molecules. The London Forces are the only ones to be constituted. As a result, breaking the intermolecular interactions in HF takes more energy than breaking them in other hydrogen halides, resulting in a higher boiling point.
Ethanol has a higher boiling point than diethyl ether because there is hydrogen bonding in the ethanol
H-bonding exists in ethanol, but none exists in diethyl ether (since the O-atom is connected to the C-atom), and the H-bonds created by the N-atom in ethylamine are weaker than those formed by the O-atom.
The hydroxyl group, which is strongly polar, is found in alcohols as a functional group.
As a result, the shared electron pair of the OH bond is drawn to the oxygen of the OH group.
As a result, the OH group’s oxygen takes on a partial negative charge, whereas hydrogen takes on a partial positive charge.
One molecule’s negative oxygen reacts with another molecule’s positive hydrogen.
Hydrogen bonding is the term for this interaction.
As a result, a significant number of alcohol molecules are firmly linked.
As a result, a significant amount of energy is required to break the connection.
As a result, the boiling point of alcohol is greater.
Such hydrogen bonding does not exist in the case of ethers.
As a result of the lesser intermolecular interaction, ethers boil at lower temperatures than other organic compounds.
As a result, the boiling point of ethanol is greater than the boiling point of diethyl ether.
Ammonia has a higher boiling point than PH3 because there is hydrogen bonding in NH3 but not in PH3
Because of the polarity of N-H interactions, molecules in ammonia NH3 are involved in intermolecular hydrogen bonding. However, because the polarity of the P-H bond is minimal relative to the N-H bond, this is not achievable in Phosphine PH3 molecules. Because phosphorus has a lower electronegativity than nitrogen, this occurs. Ammonia has a higher boiling point than phosphine, as a result.
Group 15 elements are both N and P.
B.P = -33°C for NH3, whereas B.P = -83°C for PH3.
Explanation: We already know that electronegativity diminishes as we move down the Periodic Table. This is due to the atomic radius growing in size.
The intermolecular H bonding between NH3 molecules is quite strong. Larger PH3 molecules, however, cannot establish hydrogen bonds between themselves due to their lower electronegativity. Instead, they’re separated by the Weak London dispersion force.
Because NH3 molecules have a stronger affinity, they are bound together strongly. Thus, NH3 has a higher boiling point than PH3.
Conclusion
At a standard pressure of 1 atmosphere, a compound’s melting point is the temperature at which its solid phase transforms into its liquid phase. A compound’s melting point is a physical attribute that can be used to identify a compound, similar to solubility, density, colour, and electronegativity. Compound melting points are stated in a range because determining the exact temperature at which a compound begins to melt is difficult. The temperature at which the first drops of liquid appear is considered the melting point range’s lower limit. The temperature at which the solid phase has completely transitioned to liquid phase is the range’s upper limit.
The boiling and melting points of bigger molecules are usually higher. Take a look at the boiling temperatures of several hydrocarbons as they get bigger. With more carbons and hydrogens, London forces have a larger surface area to work with, resulting in higher boiling points.
