Rank each of the below molecules' melting point and boiling point from scale of 1 (lowest) to 7 (highest) and explain why:

Question

Rank each of the below molecules' melting point and boiling point from scale of 1 (lowest) to 7 (highest) and explain why:

a) HCI

b) H₂

c) H₂O

d) NH₃

e) CCI₄

f) CO₂

e) HBr

Answer 1

Determine the imf type in each

your teacher probably wants you to draw the lewis structures

How to tell which imf type is present based on the structure

Hydrogen bonds if there's an H bonded to F, O, or N, and the F, O, or N has a lone pair

Dipole dipole if polar (asymmetrical, polar bonds)

London dispersion if nonpolar

HCl: dipole dipole, polar

H2: london dispersion, nonpolar

H2O: hydrogen bonds, H bonded to O with 2 lone pairs

NH3: hydrogen bonds, H bonded to N with 1 lone pair

CCl4: london dispersion, nonpolar

CO2: london dispersion, nonpolar

HBr: dipole dipole, polar

weakest imf --> strongest imf

london dispersion < dipole dipole < hydrogen bonds

highest boiling point = strongest imfs

For london dispersion, largest molar mass is the strongest imf

H2<CO2<CCl4

For dipole dipole, more polar bond is stronger, but these bond polarities are about the same and the internet says HBr has a higher boiling point. All molecules have london dispersion, so larger molar mass of HBr makes london dispersion forces stronger in HBr than in HCl, so HBr has a higher boiling point

HCl<HBr

Hydrogen bonds:

How many hydrogen bonds can they form per molecule? The more they can form per molecule, the stronger the hydrogen bond. This depends on the number of lone pairs on F, O, or N

NH3=1 lone pair on N, 1 hydrogen bond

H2O=2 lone pairs on O, 2 hydrogen bonds

NH3<H2O

Since

london dispersion < dipole dipole < hydrogen bonds

Then

(lowest boiling point) H2<CO2<CCl4<HCl<HBr<NH3<H2O (highest boiling point)