Periodic Trend: Atomic Radius: Videos & Practice Problems

Atomic radius is a measure of the size of an atom, defined as the distance from the nucleus to the valence electron shell. As you move down a group in the periodic table, atomic radius increases due to the addition of more electron shells, which can accommodate an increasing number of electrons. However, moving across a period, the number of electrons in the same shell increases, enhancing the attraction to the positively charged nucleus and causing a decrease in atomic radius. This results in a general trend where atomic radius decreases as you move from left to right across a period and up a group towards the top right corner of the periodic table. Exceptions to this trend can occur, particularly with transition metals and elements synthesized in laboratories, which may have less stable or well-defined atomic radii. Understanding this trend is crucial, rather than memorizing specific atomic radii values, which are typically measured in picometers.

Atomic Radius

Atomic radius refers to the distance from an atom's nucleus to its outermost electron shell, known as the valence shell. The nucleus, located at the center, contains protons, which are positively charged, and neutrons, which are neutral. The atomic radius, denoted as r, is crucial for understanding the size of an atom.

As we move down a group in the periodic table, the atomic radius increases. This is because additional electron shells are added, allowing the atom to accommodate more electrons. Each new shell increases the distance between the nucleus and the outermost electrons, resulting in a larger atomic radius.

Conversely, when moving across a period from left to right, the atomic radius tends to decrease. This occurs because, although the number of electrons increases, they are added to the same shell. The increased number of electrons in the same shell enhances the attraction between the electrons and the nucleus, effectively pulling the electron cloud closer to the nucleus and reducing the atomic radius.

In summary, the periodic trend indicates that while atomic radius increases down a group due to the addition of electron shells, it decreases across a period as the effective nuclear charge increases, drawing electrons closer to the nucleus. This dual behavior highlights the complexity of atomic structure and the interplay between electron configuration and nuclear charge.

Atomic Radius Trends

The atomic radius of elements exhibits a clear trend across the periodic table, which is essential for understanding their properties. As one moves from left to right across a period, the atomic radius generally decreases. This decrease occurs because, while the number of electron shells remains constant, the number of electrons in the outer shell increases. This increase in electrons leads to a stronger attraction to the nucleus, effectively pulling the electrons closer and resulting in a smaller atomic radius.

Conversely, as one moves up a group in the periodic table, the atomic radius also decreases. This is attributed to the reduction in the number of electron shells, which means that the outermost electrons are closer to the nucleus, further contributing to a smaller atomic radius.

Atomic radius is typically measured in picometers (pm). For example, hydrogen has an atomic radius of approximately 37 pm, and as one progresses to helium, there is a slight decrease in size. However, it is important to note that there are exceptions to this trend, particularly among transition metals and the heavier elements in the last row of the periodic table. These elements, often synthesized in laboratories, can be unstable, leading to less definitive measurements of their atomic radii.

In summary, the general trend is that the atomic radius decreases as one moves towards the top right corner of the periodic table. Understanding this trend is crucial for answering questions related to atomic size and its implications in chemical behavior.

Atomic Radius Example

When comparing the atomic radii of elements, it's essential to understand the general trend in the periodic table. As you move from the bottom left to the top right, atomic radius decreases. This trend is due to increasing nuclear charge, which pulls electrons closer to the nucleus, resulting in a smaller atomic size.

In the case of the elements potassium (K), rubidium (Rb), strontium (Sr), and calcium (Ca), we need to identify which has the largest atomic radius. Since rubidium is located further down in the periodic table compared to the others, it is the furthest from the top right corner. Therefore, Rb has the largest atomic radius among the given options.

To summarize, when determining atomic radius, remember that elements located lower in a group and to the left in a period will generally have larger atomic radii. In this instance, rubidium (Rb) stands out as the element with the largest atomic radius.

Arrange the following atoms in order of decreasing atomic radius:Sr, Se, Ne, Zn

• A. Sr > Zn > Se > Ne
• B. Zn > Sr > Se > Ne
• C. Ne > Sr > Zn > Se
• D. Zn > Se > Sr > Ne
• E. Se > Sr > Zn > Ne

Answer: (A)

Which alkaline earth metal has the smallest atomic radius? 

• A. Ca
• B. Rb
• C. Na
• D. Ra
• E. Fr

Answer: (A)

In moving from top to bottom in the same column on the periodic table, what trend is expected for atomic size? 

• A. Increase because the nucleus is getting stronger.
• B. Increase because electrons are being placed in higher n-values.
• C. Decrease because the nucleus is getting stronger.
• D. Increase because electrons are being added to degenerate orbitals.
• E. Decrease because electrons are being placed in higher n-values.

Answer: (B)

Here’s what students ask on this topic:

What is atomic radius?

Atomic radius refers to the size of an atom, typically measured from the center of the nucleus to the outer boundary of the surrounding cloud of electrons. Since the electron cloud doesn't have a firm edge, the atomic radius is often determined by how far apart atoms are when they bond together. There are a few different ways to define it, such as the covalent radius, the metallic radius, and the van der Waals radius, each relevant in different contexts.

The atomic radius tends to decrease across a period (left to right) on the periodic table because the number of protons increases, which pulls electrons closer due to the stronger electrostatic attraction. Conversely, the atomic radius generally increases down a group (top to bottom) because additional electron shells are added, making the atom larger. Understanding the concept of atomic radius is crucial in chemistry as it influences the atom's chemical properties, including reactivity and bonding characteristics.

Why does atomic radius decrease across a period?

Atomic radius decreases across a period on the periodic table primarily due to the increase in the number of protons in the nucleus of the atoms as you move from left to right. As the atomic number increases, the positive charge of the nucleus also increases. This greater positive charge attracts the electrons more strongly, pulling them closer to the nucleus.

Additionally, across a period, electrons are being added to the same principal energy level or shell. Since these electrons are not significantly shielded by additional layers of electrons, the increased nuclear charge has a more direct effect on the electron cloud. The effective nuclear charge (the net positive charge experienced by electrons in the valence shell) thus increases, causing the electron cloud to be drawn inwards, which decreases the atomic radius.

This inward pull overcomes the repulsion between electrons in the same energy level, resulting in a smaller atomic size despite the increasing number of electrons.

Why does atomic radius increase down a group?

Atomic radius increases down a group on the periodic table primarily due to two factors: the addition of electron shells and the effect of electron shielding.

As you move down a group, each successive element has an additional electron shell compared to the one above it. This means that the outermost electrons are found in orbitals that are further away from the nucleus, which naturally increases the size of the atom.

Additionally, with each new electron shell, there's an increase in the shielding effect. Inner shells of electrons shield the outer electrons from the attractive force of the nucleus. This shielding reduces the effective nuclear charge that the outermost electrons experience. Because these outer electrons are less strongly attracted to the nucleus, they can orbit at a greater distance, which further increases the atomic radius.

In summary, the atomic radius increases down a group due to the addition of electron shells and the increased electron shielding, which both cause the outermost electrons to be held less tightly and exist further from the nucleus.