Name: periodic trends in atomic size
Uploaded: 2020-09-03T13:38:28.000Z
Duration: 3 min 8 s
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Atomic Radii and Effective Nuclear Charge

An atom&rsquo;s size is dictated by the electrons or their orbitals. However, orbitals do not describe a confined space, but rather the statistical probability of where an electron can be found. So how is the atomic size defined, and what influences it?

An atomic radius can be described in two ways. Nonbonding atomic radius, or van der Waals radius of an atom, is one-half of the distance between adjacent nuclei in the atomic solid.&nbsp;

Conversely, a bonding atomic radius, or covalent radius, distinguishes between metals and nonmetals. In metals, the radius is described for atoms in their crystal structure as one-half of the distance between the centers of two neighboring atoms.&nbsp;

In nonmetal, diatomic molecules, the radius is described as one-half of the distance between the centers of bonded atoms.

The periodic table depicts variations in covalent radii&nbsp;that are often called atomic radii, which are influenced by two factors; the number of principal energy levels of valence electrons, and the effective nuclear charge.&nbsp;

The trend in atomic radii for main group elements down the columns is depicted here.

Moving down a group, the principal quantum number, n, increases by one for each element. Thus, as outer electrons get farther from the nucleus, the atomic radius increases down the group.&nbsp;

For example, moving down&nbsp;group 1, the atomic radius increases from lithium to cesium. This trend is demonstrated by the entire periodic table.

Further, the plot reveals that the atomic radius is maximum for each alkali metal and falls to a minimum with each noble gas across the period. The decreasing atomic radii across a period can be explained by the effective nuclear charge.

Recall the concept of an effective nuclear charge. In any multi-electron atom, the inner shell electrons partially shield the outer shell electrons from the pull of the nucleus. Thus, the effective nuclear charge, the charge felt by an outer electron is lesser than the actual nuclear charge.

Electrons in the same valence shell do not shield one another very effectively. Across the period, the nuclear charge increases while the number of inner shell electrons remains constant.

Thus, as the effective nuclear charge increases steadily, the shielding of outer electrons becomes less, and this leads to a decrease in atomic radii.

The radii of most transition elements, however, stay roughly constant across each row. This is because the number of electrons in the outermost principal energy level is nearly constant.&nbsp; 
&nbsp;

atomic radii and effective nuclear charge

The elements in groups of the periodic table exhibit similar chemical behavior. This similarity occurs because the members of a group have the same number and distribution of electrons in their valence shells.

Going across a period from left to right, a proton is added to the nucleus and an electron to the valence shell with each successive element. Going down the elements in a group, the number of electrons in the valence shell remains constant, but the principal quantum number increases by one each time. An understanding of the electronic structure of the elements allows us to examine some of the properties that govern their chemical behavior. These properties vary periodically as the electronic structure of the elements changes.&nbsp;

<h4>Variation in Atomic Radius</h4>

The quantum mechanical picture makes it difficult to establish a definite size of an atom. However, there are several practical ways to define the radius of atoms and, thus, to determine their relative sizes that give roughly similar values.

The atomic radius in metals is one-half of the distance between the centers of two neighboring atoms. It is one-half of the distance between the centers of bonded atoms for elements that exist as diatomic molecules.

Moving across a period from left to right, generally, each element has a smaller atomic radius than the element preceding it. This might seem counterintuitive because it implies that atoms with more electrons have a smaller atomic radius. This can be explained based on the concept of an effective nuclear charge. In any multi-electron atom, the inner shell electrons partially shield the outer shell electrons from the pull of the nucleus. &nbsp;Thus, the effective nuclear charge, the charge felt by an electron, is lesser than the actual nuclear charge (Z) and can be estimated by the following:

Zeff = Z &ndash; &sigma; &nbsp; &nbsp; &nbsp; &nbsp; &nbsp;&nbsp;&nbsp; &nbsp;

where, Zeff is the effective nuclear charge, Z is the actual nuclear charge, and &sigma; is the shielding constant,&nbsp;where the shielding constant is greater than zero but smaller than Z.

Each time we move from one element to the next across a period, Z increases by one, but the shielding increases only slightly. Thus, Zeff increases as we move from left to right across a period. The stronger pull (higher effective nuclear charge) experienced by electrons on the right side of the periodic table draws them closer to the nucleus, making the atomic radii smaller.

Core electrons efficiently shield electrons in the outermost principal level from nuclear charge, but outermost electrons do not efficiently shield one another from the nuclear charge. The larger the effective nuclear charge, the stronger the hold of the nucleus on outer electrons, and the smaller the atomic radius.

However, the radii of some transition elements stay roughly constant across each row. This is because the number of electrons in the outermost principal energy level is nearly constant, and they experience a roughly constant effective nuclear charge.

Within each period, the trend in atomic radius decreases as Z increases; &nbsp;Within each group, the trend is that atomic radius increases as Z increases.

Scanning down a group, the principal quantum number, n, increases by one for each element. Thus, the electrons are being added to a region of space that is increasingly distant from the nucleus. Consequently, the size of the atom (and its atomic radius) must increase as we increase the distance of the outermost electrons from the nucleus. This trend is illustrated for the atomic radii of the halogens in the table below.

<table style="margin-left: auto; margin-right: auto;">
	<tbody style="border: 1px solid black;">
		<tr>
			<td colspan="3" style="color: black; border: 1px solid black; text-align: center;">Atomic Radii of the Halogen Group Elements</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Atom</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Atomic radius (pm)</td>
			<td style="color: black; border: 1px solid black; text-align: center;">Nuclear charge, Z</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">F</td>
			<td style="color: black; border: 1px solid black; text-align: center;">64</td>
			<td style="color: black; border: 1px solid black; text-align: center;">9+</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Cl</td>
			<td style="color: black; border: 1px solid black; text-align: center;">99</td>
			<td style="color: black; border: 1px solid black; text-align: center;">17+</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">Br</td>
			<td style="color: black; border: 1px solid black; text-align: center;">114</td>
			<td style="color: black; border: 1px solid black; text-align: center;">35+</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">I</td>
			<td style="color: black; border: 1px solid black; text-align: center;">133</td>
			<td style="color: black; border: 1px solid black; text-align: center;">53+</td>
		</tr>
		<tr>
			<td style="color: black; border: 1px solid black; text-align: center;">At</td>
			<td style="color: black; border: 1px solid black; text-align: center;">148</td>
			<td style="color: black; border: 1px solid black; text-align: center;">85+</td>
		</tr>
	</tbody>
</table>

&nbsp;

This text is adapted from <a href="https://openstax.org/books/chemistry-2e/pages/6-5-periodic-variations-in-element-properties">Openstax Chemistry 2e, Section 6.5: Periodic Variations in Element Properties.</a>

The elements in groups of the periodic table exhibit similar chemical behavior. This similarity occurs because the members of a group have the same number ...

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Periodic Properties of the Elements

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periodic trends in atomic size

Periodic trends in atomic size help explain why atoms get larger or smaller across the periodic table. Elements in the same group show similar chemical behavior because they have the same number and arrangement of valence electrons. Across a period, each new element adds one proton to the nucleus and one electron to the valence shell. Down a group, the number of valence electrons stays the same, but the principal quantum number increases by one.

Atomic radius is not easy to define with perfect precision because atoms do not have a hard edge. Scientists use practical measurements instead. In metals, atomic radius is one-half the distance between the centers of two neighboring atoms. For elements that form diatomic molecules, it is one-half the distance between the centers of bonded atoms.

Moving from left to right across a period, atomic radius generally decreases. That may seem surprising because atoms on the right have more electrons. The reason is effective nuclear charge, or the charge an electron feels from the nucleus. It can be estimated as Z eff = Z – σ, where Z is the actual nuclear charge and σ is the shielding constant.

As Z increases by one from element to element across a period, shielding rises only a little. This means Z eff also increases from left to right. A stronger pull from the nucleus draws outer electrons closer, so the atomic radius becomes smaller. Core electrons shield outer electrons well, but outermost electrons do not shield one another very effectively. Some transition elements keep nearly the same radius across a row because the number of electrons in the outermost principal energy level changes very little, so the effective nuclear charge stays roughly constant.

The trend within a group is the opposite. Atomic radius increases as Z increases down a group. Each step down adds a principal energy level, so the outermost electrons are farther from the nucleus. That greater distance makes the atom larger. The halogens show this pattern clearly, with atomic radii increasing from fluorine to astatine.

Periodic Trends in Atomic Size

periodic classification of the elements

How Valence Electrons Shape the Periodic Table

The periodic table organizes elements by increasing atomic number, and it also shows why elements in the same group behave alike. As the electron configurations are placed into the table, a repeating pattern appears in the outer shells of atoms. These outer electrons are called valence electrons, and they play the biggest role in chemical reactions. They are the highest-energy electrons in an atom, so they are easier to lose or share than core electrons. Valence electrons also help determine...

Periodic trends in atomic size help explain why atoms get larger or smaller across the periodic table. Elements in the same group show similar chemical behavior because they have the same number and arrangement of valence electrons. Across a period, each new element adds one proton to the nucleus and one electron to the valence shell. Down a group, the number of valence electrons stays the same, but the principal quantum number increases by one. Atomic radius is not easy to define with perfect...

ionization energy

Ionization Energy Trends in the Periodic Table

Ionization energy is the energy needed to remove an electron from a gaseous atom or ion in its ground state. The first ionization energy, or IE1, removes the most loosely bound electron and forms a 1+ cation. The second ionization energy, IE2, removes the next electron, and each later ionization energy removes another electron from the atom or ion. Ionization always requires energy, so these processes are endothermic. That means ionization energy values are always positive. In general, larger...

ionic radii

How Ionic Charge Changes Ion Size

Ionic radius describes the size of an ion. It changes when an atom gains or loses electrons. Cations are positive ions with fewer electrons than the parent atom, and they are smaller than the atoms they come from. Aluminum shows this pattern clearly. A neutral aluminum atom has a covalent radius of 118 pm, while Al3+ has an ionic radius of 68 pm. When electrons are removed from the outer valence shell, the remaining core electrons feel a stronger effective nuclear charge, or Zeff. That pull...

electron affinity

Electron Affinity Trends on the Periodic Table

Electron affinity is the energy change that happens when a gaseous atom gains an electron and becomes a negative ion, or anion. For many elements, this process releases energy, so the electron affinity value is negative. For some elements, energy must be added instead, so the value is positive. Electron affinity also applies when more than one electron is added. The second electron affinity is the energy change for adding an electron to an anion to make a 2− ion. Later electron affinities...

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Alkali metals are Group 1 elements with one valence electron. They are soft, shiny metallic solids that are malleable, ductile, and good conductors of heat and electricity. Their melting points are unusually low for metals and drop as you move down the group. Density generally increases down the group, except for potassium. The trend in atomic size helps explain their behavior. Moving from lithium to cesium, each element has a higher principal quantum number, so the outer electron is farther...

halogens

Halogens: Trends and Reactions

Halogens are Group 17 nonmetals that show clear trends in structure and reactivity. At room temperature, fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid. Astatine is highly unstable and radioactive, so most of its properties are still unknown because it has a very short half-life. Tennessine is a synthetic element that is also predicted to belong to this group. These elements do not usually exist as single atoms. Instead, they form diatomic molecules, which means...

noble gases

Noble Gases: Sources, Uses, and Reactions

Noble gases are the group 18 elements helium, neon, argon, krypton, xenon, and radon. They were called “noble” because they were once thought to be nonreactive. That idea changed in 1962, when Dr. Neil Bartlett at the University of British Columbia showed that this assumption was not fully correct. These gases occur in very small amounts in the atmosphere. Some natural gas contains 1–2% helium by mass. Helium is separated from natural gas by removing the condensable components, which leaves...

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Learn why atomic radius changes across periods and groups, and how effective nuclear charge and shielding shape the size of atoms in the periodic table.

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Periodic Trends in Atomic Size

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