Question

Ordinary ice (solid water) melts at \(0^{\circ} \mathrm{C},\) whereas dry ice (solid carbon dioxide) melts at a much lower temperature. Explain the differences in the melting points of these two substances on the basis of the intermolecular forces involved.

Short answer

The differences in melting points between ordinary ice and dry ice can be attributed to the different intermolecular forces present in each substance. Water experiences the strongest type of intermolecular force, hydrogen bonding, due to the presence of hydrogen atoms directly bonded to oxygen, a highly electronegative atom. In contrast, carbon dioxide only experiences weaker London dispersion and dipole-dipole forces. This leads to a higher melting point for water, which melts at \(0^{\circ} \mathrm{C}\), while dry ice melts at a much lower temperature due to the weaker intermolecular forces involved.

Step-by-step solution

Identify the types of intermolecular forces in Water and Carbon Dioxide

First, let's identify the types of intermolecular forces that are present within water (H2O) and carbon dioxide (CO2). Both substances experience the two weaker intermolecular forces: London Dispersion (the weakest) and Dipole-Dipole (the middle strength one), but only water experiences hydrogen bonding, which is the strongest type of intermolecular forces. The hydrogen bonding occurs because water has hydrogen atoms directly bonded to oxygen, a highly electronegative atom.

Understand the impact of intermolecular forces on melting points

The strength of the intermolecular forces in a substance is directly related to the melting point of that substance. Substances with stronger intermolecular forces generally have higher melting points because it requires more energy to break the bonds between the molecules, while substances with weaker intermolecular forces have lower melting points as less energy is needed to separate the molecules.

Compare the intermolecular forces in water and carbon dioxide

Water has hydrogen bonding, which is the strongest type of intermolecular force, while carbon dioxide does not. Carbon dioxide mainly experiences London dispersion and dipole-dipole forces, which are generally weaker than hydrogen bonding. This difference in intermolecular forces between water and carbon dioxide plays a crucial role in determining their respective melting points.

Relate the intermolecular forces to the melting points of ice and dry ice

The stronger hydrogen bonding present in water leads to a higher melting point than carbon dioxide, which only experiences the weaker London dispersion and dipole-dipole forces. As a result, ordinary ice melts at \(0^{\circ} \mathrm{C}\), while dry ice melts at a much lower temperature due to the weaker intermolecular forces involved. In conclusion, the differences in melting points between ordinary ice (solid water) and dry ice (solid carbon dioxide) can be explained by the differences in the intermolecular forces with water experiencing stronger hydrogen bonding while carbon dioxide experiencing weaker London dispersion and dipole-dipole forces.

Key concepts

Melting Points

Melting points are crucial in understanding the physical properties of a substance.
They indicate the temperature at which a solid changes to liquid. This transformation requires enough energy to overcome the forces that hold the solid's molecules together.
More robust intermolecular forces mean more energy—and thus a higher temperature—is needed to break these bonds during the melting process. This implies that the stronger the intermolecular forces in a substance, the higher its melting point will be.
For example, ordinary ice, which consists of water molecules, has a melting point of 0°C due to strong hydrogen bonding, while dry ice, composed of carbon dioxide, has a much lower melting point.
This is because the intermolecular forces in dry ice are not as strong.

• Stronger forces = higher melting point
• Weaker forces = lower melting point

Hydrogen Bonding

Hydrogen bonding is a powerful type of intermolecular force, usually recognized as the strongest among the common ones.
This occurs when hydrogen is directly bonded to a highly electronegative atom such as oxygen, nitrogen, or fluorine.
In water, hydrogen bonding is responsible for its relatively high melting and boiling points, as well as other unique properties like surface tension and high heat capacity.
Because water molecules form extensive networks of these hydrogen bonds, a substantial amount of energy is needed to disrupt them, which results in water's higher melting point.
Imagine these bonds like a network of strong handshakes keeping water molecules together, requiring significant energy to break apart during melting.

London Dispersion Forces

London dispersion forces are the weakest of all intermolecular forces.
They arise due to the temporary dipoles that occur when electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles.
These forces are present in all molecules, regardless of whether they are polar or nonpolar.
In gases like carbon dioxide, dispersion forces are the primary means of attraction.
Since these forces are weak, substances like dry ice, which relies heavily on them, have relatively low melting points.
Compared to water's hydrogen bonding, dispersion forces in carbon dioxide allow molecules to separate more readily, hence its much lower melting point.

• Present in all molecules
• Weak, temporary interactions

Dipole-Dipole Interactions

Dipole-dipole interactions are stronger than dispersion forces but weaker than hydrogen bonds.
They occur in polar molecules where there is an unequal distribution of electron density.
The positive end of one molecule attracts the negative end of another, providing a degree of stability to the molecular structure.
In carbon dioxide and water, dipole-dipole interactions occur due to their polar molecules.
However, in carbon dioxide, these interactions are not as pronounced as in other substances due to its linear molecular shape.
Consequently, these interactions are overshadowed by stronger forces like hydrogen bonding in water, significantly affecting melting points. - Occur in polar molecules - Position and shape affect strength - Contribute to physical properties