# Atomic and Ionic Radii: Trends Among Groups and Periods of the Periodic Table

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• 0:13 The Size of an Atom
• 2:14 Group Trends
• 3:43 Periodic Trends
• 5:53 Lesson Summary

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Lesson Transcript
Instructor: Kristin Born

Kristin has an M.S. in Chemistry and has taught many at many levels, including introductory and AP Chemistry.

Atoms are VERY tiny. How do we measure their size? This lesson will explain how the size of an atom is measured and teach you how to predict the relative size of an atom based on where it is located on the periodic table.

## The Size of an Atom

When you picture an atom, you probably see a bunch of protons and neutrons crammed together in a tiny little nucleus surrounded by a bunch of electrons zipping around the outside of a nucleus. It should make sense that the size of an atom is really dependent on how far away the electrons are - more specifically, how far away the outer electrons, or valence electrons, are. If they are zipping around really close to the nucleus in the first energy level, the atom will likely be very small, and if the valence electrons are flying around way out in the fifth energy level, the atom will be very large. The size of an atom is dependent on how much space the electrons take up.

But if electrons are always moving, and we never really know exactly where an electron is at any given time, how do we measure the size of an atom? You may think of an atom as being a small, hard sphere, when in reality, its outer boundaries are very difficult to define.

Measuring an atom's size is like measuring the size of a marshmallow: It depends on how it's measured. Is it apart from the rest, or is it squished into its packaging? When the size of an atom is measured, it's important to specify if it's an isolated atom, or if it's one that is bonded to something else. Typically, the atomic radius is measured as half the distance between the nuclei of two bonded atoms. This measured radius is often slightly smaller than an atom's actual radius, but because the nucleus of an atom is very well defined and easy to detect, this measurement is the most often used.

The rest of this lesson will be focused on the trends that the atoms have in size as you move down a group or across a row on the periodic table. A trend is just a tendency to change in a predictable way. We can use these trends to compare the relative sizes of two different atoms on the table.

## Group Trends

Remember that a group in the periodic table is just a vertical column, so we will only be comparing elements in the same column. As you move down a group, you will notice that the principal quantum number increases by one. This means that electrons are going to be filling energy levels farther and farther away from the nucleus. You can think of energy levels like layers in an atom. As the number of protons in an atom increases, the number of electrons will also increase. These electrons need room to move around, and each energy level can only hold so many electrons. So at the start of each row on the periodic table, a new energy level has to be 'opened' for these new electrons to be added.

For example, if we compare elements in the first column on the periodic table, hydrogen has one electron, and it is located in the first energy level. Lithium has three electrons: two of them filling the first energy level, and one of them (the valence electron) needing to be added to the newly created second energy level. Finally, let's compare this to sodium, with 11 electrons. Two of them will fill the first level, eight will fill the second level, and one (the valence electron) will need to be added to the newly created third level. Because each added level is farther and farther away from the nucleus, the atomic radius increases as you move down a group on the periodic table.

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