Question

On average, are the intermolecular forces stronger in \(\mathrm{H}_{2} \mathrm{O}(\boldsymbol{\ell})\) or in \(\mathrm{H}_{2} \mathrm{O}(\mathrm{g})\) ? Explain.

Short answer

The intermolecular forces in liquid water (H2O(l)) are stronger than in water vapor (H2O(g)) due to the closer proximity of the water molecules in the liquid phase, which allows for more significant hydrogen bonding.

Step-by-step solution

Identify the Phase and Intermolecular Forces

First, identify the phases being compared - liquid (H2O(l)) and gas (H2O(g)). H2O, or water, is capable of forming relatively strong hydrogen bonds due to its polar nature.

Compare Intermolecular Forces

In the liquid phase, the molecules are closer together and therefore the intermolecular forces, specifically hydrogen bonds, will be stronger. In the gaseous phase, the water molecules are more spread out and can move freely, meaning the intermolecular forces will be weaker.

Draw Conclusion

So, based on these analyses, we can conclude that H2O(l) would have stronger intermolecular forces than H2O(g).

Key concepts

Phase Changes

Phase changes occur when a substance transitions between different states of matter - solid, liquid, and gas. These changes can happen through processes like melting, freezing, vaporization, condensation, sublimation, and deposition. Each process requires energy changes and usually involves the uptake or release of heat.
When we think about phase changes, it helps to remember that:

• Melting and vaporization require energy input to overcome intermolecular forces.
• Freezing and condensation release energy as molecules slow down and form stronger interactions.

In the context of water, from liquid to gas (vaporization), energy is required to break the hydrogen bonds holding the water molecules closely packed. As water vaporizes, it moves from a state of higher density and stronger molecular interactions (liquid) to one of low density and weaker interactions (gas). This fundamental change in intermolecular forces significantly affects properties like density, viscosity, and even solubility during phase transitions.

Hydrogen Bonding

Hydrogen bonding is a special type of dipole-dipole interaction that occurs when hydrogen, covalently bonded to a highly electronegative atom like oxygen, nitrogen, or fluorine, experiences attraction towards another electronegative atom.
This interaction is especially powerful in water molecules. In water, each hydrogen atom can potentially form a bond with a nearby oxygen atom from another water molecule. This is why water can create an exceptionally strong network of connections, making it quite unique among other substances.
Hydrogen bonds are responsible for many of water's unusual properties, such as its high boiling point relative to its molecular size and its ability to stay liquid over a wide range of temperatures. This network of hydrogen bonds also means that a lot of energy is required to break these interactions to convert water from liquid into gas, as seen in phase changes.

Liquid vs Gas Phases

Understanding the difference between liquid and gas phases helps to comprehend how substances behave differently under varying conditions.
In the liquid phase, water molecules are tightly packed, ensuring that the attractive forces, such as hydrogen bonds, are maximized due to proximity. This results in water's relatively high density and surface tension in its liquid form.
On the other hand, in the gaseous phase, molecules spread out and occupy more space. The intermolecular forces in this phase are much weaker because the molecules are far apart and move rapidly. This difference influences many physical properties:

• Liquids have fixed volume but can change shape, while gases can expand to fill a container.
• Density in liquids is higher than in gases because of the close molecular packing.
• The viscosity of liquids is significantly greater than that of gases due to stronger intermolecular forces in the liquid state.

These characteristics explain why water in its liquid state (H₂O(l)) has stronger intermolecular forces compared to its gas state (H₂O(g)).