Question

The boiling points of the noble gas elements are listed below. Comment on the trend in the boiling points. Why do the boiling points vary in this manner? $$\begin{aligned} &\begin{array}{llll} \mathrm{He} & -272^{\circ} \mathrm{C} & \mathrm{Kr} & -152.3^{\circ} \mathrm{C} \end{array}\\\ &\begin{array}{llll} \text { Ne } & -245.9^{\circ} \mathrm{C} & \text { Xe }-107.1^{\circ} \mathrm{C} \end{array}\\\ &\begin{array}{llll} \mathrm{Ar} & -185.7^{\circ} \mathrm{C} & \mathrm{Rn} & -61.8^{\circ} \mathrm{C} \end{array} \end{aligned}$$

Short answer

The boiling points of noble gases increase as we move down the group in the periodic table, from helium to radon. This trend is due to the increasing atomic size and mass, resulting in a greater number of electrons and a larger electron cloud. Stronger van der Waals forces occur between atoms with more electrons in their electron cloud, as larger electron clouds lead to a higher likelihood of electron density fluctuations that generate temporary dipoles. Consequently, more energy is required to overcome these stronger van der Waals forces, leading to higher boiling points.

Step-by-step solution

Identify the trend in boiling points

Observe the boiling points of the noble gases: Helium (He): \(-272^\circ C\) Neon (Ne): \(-245.9^\circ C\) Argon (Ar): \(-185.7^\circ C\) Krypton (Kr): \(-152.3^\circ C\) Xenon (Xe): \(-107.1^\circ C\) Radon (Rn): \(-61.8^\circ C\) As we move down the group in the periodic table (from He to Ne to Ar, and so on), the boiling points increase, signifying that the amount of energy required for the substance to transition from liquid to vapor also increases.

Explain the trend

Now, let's give a reason for the observed trend in noble gas boiling points. Noble gases feature complete valence electron shells, resulting in limited chemical reactivity. This means the primary force responsible for their transition to a gaseous state is the van der Waals forces between atoms, which are induced dipole-dipole interactions. As you move down the group from He to Rn, the atomic size and mass increase. This results in a greater number of electrons and thus a larger electron cloud. A larger electron cloud causes stronger van der Waals forces between the noble gas atoms due to a higher likelihood of electron density fluctuations that generate temporary dipoles. Hence, as you move down the group, the boiling points increase because stronger van der Waals forces between atoms require more energy to break. The larger number of electrons in the electron cloud contributes to stronger induced dipole-dipole interactions, which in turn result in higher boiling points.

Key concepts

van der Waals Forces

To begin, let's delve into the van der Waals forces, a fundamental concept in understanding the boiling points of different substances, particularly inert gases such as the noble gases. Van der Waals forces are weak electrical forces that cause attraction between molecules or atoms. These forces play a crucial role in determining the physical properties of matter, including boiling and melting points.

For noble gases, which are characterized by their complete valence shells and lack of strong chemical bonds like ionic or covalent bonds, van der Waals forces are the only significant intermolecular attractions. As the main source of these forces, fluctuations in the electron cloud surrounding the atom create temporary dipoles, and these dipoles induce attractions between neighboring atoms.

In the context of noble gases, the size of the atom—and consequently, its electron cloud—becomes larger as you travel down the periodic table. This increased size allows for a greater surface area over which the van der Waals forces can act, leading to stronger attractions between particles and thus higher boiling points.

Induced Dipole-Dipole Interactions

Inside each noble gas atom, the electron movement is random and can lead to a temporary uneven distribution of charge, known as a 'dipole'. When this happens, nearby atoms respond similarly, resulting in a 'dance' of fluctuating charges that hold the atoms together. The strength of this interaction depends on the number of electrons interacting and the distance between them; this is why heavier noble gases, with more electrons and larger electron clouds, have stronger van der Waals forces and therefore higher boiling points.

Periodic Trends

Understanding the periodic trends can explain a variety of physical and chemical properties observed among elements. As it pertains to boiling points, particularly of noble gases, we observe a trend that correlates with their position on the periodic table.

As you move down the group of noble gases in the periodic table, from Helium down to Radon, the atomic number increases. With each step down, the atoms gain an additional electron shell. This increase in the number of shells leads to a larger atomic radius, which in turn means a more expansive electron cloud around the nucleus.

Atomic Size and Boiling Points

As atomic size increases, the van der Waals forces become more pronounced due to the larger area over which the temporary dipoles can influence other atoms. This trend is neatly demonstrated by the incremental increase in boiling points among the noble gases. However, it's crucial to understand that these trends have exceptions and are influenced by multiple factors, including the atom's electron configuration and the types of interactions between atoms. The trend of noble gases provides an uncomplicated example since they are not involved in chemical bonding, which allows the periodic trend for boiling points to be more apparent.

States of Matter

Exploring the states of matter gives us insight into how substances change between solid, liquid, and gas phases. These phases are determined largely by temperature and pressure, with the boiling point being a critical temperature at which a substance moves from the liquid to the gas phase.

When we discuss noble gases, their boiling points correspond to the temperature at which the balance between kinetic energy, which propels particles apart, and van der Waals forces, which attract particles together, shifts in favor of the kinetic energy.

Transitions Between States

At the boiling point, the kinetic energy within the noble gas atoms becomes sufficient to overcome the van der Waals attractions that hold the atoms close in the liquid phase. As such, the boiling point is a direct indicator of the strength of these intermolecular forces. The higher the boiling point, the more energy is required for the atoms or molecules to escape the liquid state into the gaseous state, which is why Radon, with the strongest van der Waals forces among the noble gases, has the highest boiling point of the group.