O₂ has a nonpolar covalent bond. Nonpolar covalent bonds are bonds formed between atoms that have the same (or nearly the same) electronegativity. Since both oxygen atoms have the same electronegativity, they will have a nonpolar covalent bond between them. 

The intramolecular bonds within a water molecule are polar covalent bonds. O-H bonds are polar covalent since there is more electron density around the oxygen atom than there is around the hydrogen atom, making the oxygen atom slightly negative and the hydrogen atom slighty positive. This polarity allows for intermolecular hydrogen bonding; water does not exhibit intramolecular hydrogen bonding.

This bond is not ionic since the oxygen does not completly steal the electron away from the hydrogen (the electronegativity difference between the two atoms is between 0.4 and 1.7 on the Pauling scale).

 is an ionic compound, while  is a nonpolar covalent compound. Remember that polarity results from a difference in the electronegativities of the atoms involved in the bond. Too great of a difference will result in an ionic bond; two of the same atoms will have zero difference, resulting in a nonpolar bond.

, , and  contain polar covalent bonds. In the first two, oxygen will carry a slight negative charge, leaving sulfate and phosphorus with slight positive charges. In , chlorine will carry a slight negative charge, leaving the carbon slightly positive. In  and , symmetry helps to balance the polar bonds, resulting in an overall nonpolar molecule, even though the individual bonds are nonpolar.

Covalency is a form of electron sharing that lets an atom fulfill the octet rule. The sharing may be unequal, in which case the more electronegative species more strongly attracts the electrons than the weaker, less electronegative species, creating a polar covalent bond. In the Lewis dot model, however, we draw the electrons in freeze-frame, equally distributed between the various atoms.

A sigma bond is a single covalent bond, involving an electron pair located between the two bonding atoms. A pi bond occurs when the p orbitals above and below the bonding atoms overlap, or when the p orbitals to the left and right overlap. In any covalent bond, the first bond formed is a sigma bond and any additional bonds must be pi bonds. Initial orbital overlap always comes from the sigma, or s, subshell; subsequent overlap comes from the pi, or p, subshells.

This question is testing your ability to draw Lewis dot structures and your knowledge of how resonance effects bond length. The N-O bond with the greatest pi-bond character will be the shortest; thus, we are looking for a double- or triple-bond between nitrogen and oxygen.

Hydroxylamine () only contains single bonds, which have the least pi-bond character.

The nitrite and nitrate ions both have a double bond between the nitrogen and oxygen, but also one or more single bonds between these elements. This means that, because of resonance, the N-O bonds in these molecules will be averaged and our average bond order will be somewhere between single and double. Nitrate will have one double bond and two single bonds, for an average bond order of 1.33. Nitrite has one single bond and one double bond, for a bond order of 1.5.

The nitrosyl ion, however, will contain a triple bonds between the nitrogen and oxygen atoms, giving it the greatest pi-bond character. This bond will contain more energy and be shorter than the bonds in the other answer options.

Water exhibits hydrogen bonding. Each hydrogen atom is bond to a highly electronegative oxygen atom, resulting in a slight positive dipole on the hydrogen and a slight negative dipole on the oxygen. Within a solution of water, these dipoles can align, causing attractive forces between the hydrogens of one water molecule and the oxygen of another. Hydrogen bonding is a specialized form of dipole moment; since water has a permanent dipole moment, dipole-dipole attractions are present. Finally, like all molecules, water exhibits London dispersion intermolecular forces as well.

Phase phenomena, such as vapor pressure, are best explained as a product of intermolecular forces. Strong intermolecular forces, like hydrogen bonds or dipole interactions, would make us expect weak vapor pressure, increasing attraction between molecules in the same phase and preventing transition. Large vapor pressures are more likely due to weak van der Waals intermolecular interactions, rather than stronger forces.

London dispersion forces are weak, temporary attracting forces caused when electrons in adjacent atoms move into asymmetrical arrangements about their nuclei, forming temporary dipoles. Since electrons are constantly moving in any atom, both polar and nonpolar substances can develop temporary dipoles and experience London dispersion forces. These forces occur randomly as electrons form spontaneous polarized distributions, before immediately dissipating as the electrons move to new positions.

The passage states that jar B has a boiling point of , the lowest of all three molecules; therefore, jar B contains molecules that have the weakest intermolecular forces. Of all the intermolecular forces, London dispersion forces (or van der Waals forces) are the weakest. These forces occur in all molecules because the movement of electrons inside molecules causes brief polarities that can be used to form intermolecular forces. Since they are so brief, London dispersion forces are very weak and easy to break.

Many nonpolar molecules, such as propane, only contain London dispersion forces and lack all forms of dipole interaction. This results in very weak intermolecular interactions for these compounds, and will affect physical properties like vapor pressure, boiling point, and surface tension. Weak intermolecular forces will increase vapor pressure, decrease boiling point, and decrease surface tension.

Remember that polar molecules, such as methanol and dichloromethane, also contain London dispersion forces. They just contain much stronger forces (hydrogen bonding and dipole-dipole interactions) that make the London dispersion forces negligible.