van der Vaals forces of attraction and dipole-dipole interaction
it doesn't
Given that methane is a gas at room temperature, we can see that its molecules are attracted to each other only by weak intermolecular forces. But normally, if I say "methane is a compound consisting of..." the statement concludes, one carbon atom and four hydrogen atoms. That's what the methane molecule consists of.
The only intermolecular forces in this long hydrocarbon will be dispersion forces.
If the intermolecular forces are great enough they can hold the molecules together as a liquid. If they are even stronger they will hold the molecules together as a solid. Water has nearly the same mass as methane and ammonia molecules, but the greater molecular forces between water molecules causes the water to be liquid at room temperature, while ammonia and methane, with weaker intermolecular forces, are gases at room temperature.
Methanol has the highest boiling point among methane, chloromethane, and methanol. This is because methanol has stronger intermolecular forces (hydrogen bonding) compared to methane (only dispersion forces) and chloromethane (dipole-dipole forces).
The intermolecular forces in CH4 (methane) are London dispersion forces. Methane molecules are non-polar, so the only type of intermolecular force present is the weak temporary dipole-induced dipole interactions between molecules.
In a single molecule of CH4 you would have intramolecular forces that are covalent bonds. The intermolecular forces that exist between molecules of CH4 are called dispersion forces. These forces are the only intermolecular forces that occur between non-polar molecules.
The intermolecular forces present in C2H5OH (ethanol) are hydrogen bonding, dipole-dipole interactions, and London dispersion forces.
The intermolecular forces present in diethyl ether are primarily London dispersion forces and dipole-dipole interactions.
London forces are present in chlorine molecules.
methane
Dipole forces and London forces are present between these molecules.
it doesn't
In SiF4, the intermolecular forces present are London dispersion forces. These forces arise due to temporary fluctuations in electron distribution within the molecule, leading to weak attractions between neighboring molecules.
Given that methane is a gas at room temperature, we can see that its molecules are attracted to each other only by weak intermolecular forces. But normally, if I say "methane is a compound consisting of..." the statement concludes, one carbon atom and four hydrogen atoms. That's what the methane molecule consists of.
The only intermolecular forces in this long hydrocarbon will be dispersion forces.
If the intermolecular forces are great enough they can hold the molecules together as a liquid. If they are even stronger they will hold the molecules together as a solid. Water has nearly the same mass as methane and ammonia molecules, but the greater molecular forces between water molecules causes the water to be liquid at room temperature, while ammonia and methane, with weaker intermolecular forces, are gases at room temperature.