why does C4H10 have a higher BP than C2H6,C3H8, CH4
The molar mass is higher. The IM forces (London) are higher.
C4H10, also known as butane, has a higher boiling point (BP) compared to C2H6 (ethane), C3H8 (propane), and CH4 (methane) due to differences in their molecular structures and intermolecular forces.
The boiling point of a substance is mainly determined by the strength of the intermolecular forces between its molecules. These intermolecular forces are primarily due to London dispersion forces (also known as van der Waals forces) and dipole-dipole interactions.
Let's compare the molecular structures of these hydrocarbons:
- Ethane (C2H6): It consists of two carbon atoms bonded to six hydrogen atoms, forming a linear structure with no polar bonds. Ethane primarily experiences London dispersion forces. These forces are relatively weak since ethane molecules have a small number of electrons, resulting in low boiling points.
- Propane (C3H8): Propane consists of three carbon atoms bonded to eight hydrogen atoms, creating a linear structure like ethane. Similar to ethane, propane molecules also predominantly experience London dispersion forces. However, propane molecules are larger, have more electrons, and thus exhibit slightly stronger dispersion forces and higher boiling points compared to ethane.
- Methane (CH4): Methane has a tetrahedral molecular structure, where a carbon atom forms four single bonds with four hydrogen atoms. Due to its symmetrical structure, methane experiences only London dispersion forces. These forces are weaker than dipole-dipole interactions, resulting in a relatively low boiling point.
- Butane (C4H10): Butane contains four carbon atoms bonded to ten hydrogen atoms. Its molecular structure is more complex than the other hydrocarbons, featuring a branched structure. Butane has a larger surface area and more electrons compared to the other hydrocarbons, resulting in stronger London dispersion forces. Additionally, butane can also experience dipole-dipole interactions due to temporary dipoles caused by the movement of electrons within the molecule. These stronger intermolecular forces lead to a higher boiling point for butane compared to ethane, propane, and methane.
In summary, the increasing boiling points from methane to butane can be attributed to the increase in molecular size, surface area, and the strength of the intermolecular forces (primarily London dispersion forces and dipole-dipole interactions) in each molecule.