Lesson Explainer: Periodicity Chemistry

In this explainer, we will learn how to describe the trends in properties across periods of the periodic table.

The periodic table provides us with a lot of additional information that is not necessarily obvious at first inspection. The groups and periods of the periodic table reveal what are known as periodic trends, and this periodicity gives us valuable insights into the properties of the different chemical elements.

Definition: Periodic Trend

A periodic trend is a specific pattern in the physical or chemical properties of the elements.

Periodic trends may occur across a period, up or down a group, or from one corner of the periodic table to another corner. In this explainer, we will investigate periodic trends as they relate to the atomic radius, ionic radius, melting point, and conductivity.

We will focus primarily on the period 3 elements, as these provide us with a good selection of elements, a few anomalies, and an understanding that we can apply to other rows of the periodic table.

The first periodic trend we will investigate is the trend of atomic radius values. However, before we begin to examine the trend itself, is important to have a clear understanding of exactly what we mean by atomic radius and how it is measured.

The atomic radius is essentially a quantity that defines the size of an atom. However, it can be challenging to define the outer edge of an atom as a fixed line, because electron density is spread out over space. This means that it is not possible to measure the size of an individual atom directly, and chemists have had to devise different ways to measure the atomic radius indirectly.

Most chemists have used covalent atomic radius values to indirectly determine the size of an individual atom. The covalent radius values are determined from covalent bond lengths. The covalent bond lengths are expressed in units of either picometres (pm) or angstroms (̊A).

Definition: Covalent Atomic Radius

The covalent atomic radius is one-half the distance between the nuclei of two atoms covalently bonded together.

We make the assumption that the covalent bond length is equivalent to the sum of two covalent atomic radius values. For example, depending on what data booklet is used, we may discover that the bond length between the two atoms in a molecule of hydrogen is 0.56̊A. From this value, we would then know that the covalent atomic radius of each hydrogen atom would be 0.28̊A.

It is clear to see, provided we know the bond length, that determining the covalent atomic radius in homogeneous diatomic molecules is relatively simple. However, determining the covalent atomic radius in heterogeneous diatomic molecules requires an extra step. Imagine, for example, that we were interested in determining the covalent atomic radius of a chlorine atom in a molecule of hydrogen chloride. We may be able to measure the bond length, but it would be impossible for us to determine which percentage of the bond length was attributable to the hydrogen atom and to the chlorine atom.

In this case, the solution is to use previous data from homogeneous diatomic molecules and simple subtraction equations to determine values that we do not know. For example, we may have measured the bond length of a hydrogen molecule to be 0.56̊A.

We then measure the bond length of the hydrogen chloride molecule to be 1.27̊A.

By halving the bond length of the hydrogen molecule, we know that the atomic radius of hydrogen is 0.28̊A, and we can take this value away from 1.27̊A: =1.27̊0.28̊=0.99̊.AAA

We have determined that the covalent atomic radius of chlorine is 0.99̊A using this data.

Now that we understand what atomic radius is and how we calculate it, we will discuss how it varies across the periodic table. The atomic radius tends to increase as we move from the right-hand side of the periodic table to the left-hand side of the periodic table. The atomic radius also tends to increase as we move down each column or group of the periodic table. These two trends are displayed on a blank periodic table for clarity.

The trend in atomic radius values can be understood by considering the different numbers of protons and electrons in the atoms of chemical elements. The atomic radius represents the distance from the nucleus to the outer electrons of an atom.

Outer-shell electrons are subjected to a greater nuclear charge if the atomic number is high. The electrons are pulled closer toward the nucleus, and the size of the atom is reduced. This explains why the atomic radius tends to decrease as we move across any one row or period of the periodic table. The nuclear charge increases, and the electrons are pulled closer to the atomic nucleus. It is true that the number of electrons increases as well, but they are filling up the same quantum shell, and, consequently, they do not offset the increasing nuclear charge.

The atomic radius increases as we move down any one group of the periodic table because elements have more electron shells if they are closer to the bottom of the periodic table. The outer shell of electrons ends up being farther from the atomic nucleus as we move down any one column of the periodic table. As we descend the groups, the nuclear charge also increases, creating a greater attraction to the center of the atom. However, the increase in quantum number and shielding from additional shells of electrons outweigh the increase in nuclear charge.

Example 1: Selecting the Element with the Largest Atomic Radius

Which of the following elements has the largest atomic radius across the periodic table?

  1. Copper
  2. Cesium
  3. Lithium
  4. Neon
  5. Magnesium

Answer

The periodic trend of atomic radius operates in two directions: across the periodic table and down the periodic table. As we go down the groups of the periodic table, the atomic radius increases due to additional electrons in increasing quantum shells. As we go across the periods of the periodic table, the atomic radius decreases due to increasing nuclear charge generated through the addition of extra protons in the nuclei of the atoms as the atomic number increases from left to right.

If we combine these two concepts, we would expect to find the element with the largest atoms to be at the bottom of a group and in the farthest left group of the periodic table. The farthest left group on the periodic table is the alkali metals, of which we have two choices: cesium and lithium. Cesium is at the very bottom of the group, discounting the synthetic element francium. Cesium is lower down the group than lithium and so will have the largest atomic radius across the periodic table, and as such, the answer is B.

Ionic radius is defined somewhat similarly to the covalent atomic radius. An ionic bond length is the distance between the nucleus of one positive ion and the nucleus of an adjacent negative ion. This distance between these two ions is the sum of their individual radii.

Definition: Ionic Radius

The ionic radius is the radius of a single atomic ion in an ionic lattice.

It is not possible to describe how ionic radius values change across the whole periodic table, because only some elements can form ionic compounds. Ionic compounds are formed when metal elements on one side of the periodic table exchange electrons with nonmetal elements on the other side of the periodic table. We cannot describe the trend in ionic radius values for all the chemical elements, but we can describe the trend in ionic radius values for the metal cations and nonmetal anions that do form ionic lattices.

Metal atoms lose electrons when they form ionic compounds, but they do not lose any protons. This means that sodium atoms form Na+ ions that have eight electrons. It also means that magnesium and aluminum atoms form Mg2+ and Al3+ ions that have eight electrons as well. The nuclear charge values of these ions are different, but the ions all have the same total number of electrons. The increasing nuclear charge is exerted on the same number of electrons, and the ionic radius has to decrease as we move from group one ions like Na+ through to group three ions like Al3+.

The negative ions of ionic compounds are formed when electrons are added to the valence shell of nonmetal atoms. Phosphorus is a group 15 element, and phosphorus atoms form P3 ions when they gain three electrons. Sulfur and chlorine are group 16 and 17 elements, and their atoms form the S2 and Cl ions when they gain two or one electrons respectively. All of these ions are isoelectronic, because they contain the same number of electrons and they have the same electronic configuration.

The three anions are isoelectronic, but they do not all have the same number of protons. The number of protons increases with the atomic number. The nuclear charge is higher in the ions that have more protons, and this higher concentration of protons makes the outer electrons draw closer to the atomic nucleus. The chlorine anion (Cl) has a larger ionic radius than the sulfur anion (S2), and the sulfur anion has a larger ionic radius than the phosphorus anion (P3). We can extend this line of reasoning to state that the ionic radius tends to decrease as we move from group 15 anions through to group 17 anions.

Example 2: Explaining Why an Oxygen Ion Has a Larger Radius than a Sodium Ion

Which of the following statements best explains why an oxygen ion (O2) has a larger ionic radius than a sodium ion (Na+)?

  1. An oxygen ion gains electrons and forms a negatively charged ion.
  2. There are more subatomic particles in the nucleus of an oxygen ion than in that of a sodium ion.
  3. Metal ions are always smaller than nonmetal ions.
  4. An oxygen ion has fewer protons in its nucleus than a sodium ion.
  5. A sodium ion only has a single charge, but an oxygen ion has a double negative charge.

Answer

An atom of sodium has 11 protons and 11 electrons. When it forms a positive ion one of these electrons is lost, leaving 10. Comparatively, an atom of oxygen has eight protons and eight electrons, however, when it forms a negative ion, it gains two electrons, leaving a total of 10 as well. We can use the term isoelectronic to describe a sodium ion and an oxygen ion. This word means that they are identical in terms of electronic configuration.

When comparing two ions with identical electronic configurations, the number of protons in the nucleus of the respective ions will determine the ionic radius. Electrostatic attraction will mean that the ion with the greater number of protons in the nucleus will create a greater force of attraction, resulting in a smaller ionic radius.

Using this information, we can evaluate the answer options. Answer A is certainly true but does not explain why the oxygen ion has a larger ionic radius. Sodium has more subatomic particles in the nucleus than oxygen, and so answer B is also incorrect.

Answer C is far too general a statement to be true, given the vast range of different metals and nonmetals on the periodic table. Answer E is incorrect, as the amount of charge does not result in a pattern in terms of ionic radius.

Going back to answer D, an oxygen ion does have fewer protons, and as such, it will mean that the ionic radius is larger than sodium, as described above. Answer D is the correct answer.

Carbon and silicon in group 14, as well as the noble gases in group 18, do not readily form ions, and as such, they are not considered when discussing the periodic trend of the ionic radius. A graphical representation of the atomic radii combined with the ionic radii (measured in picometres) for periods 1 to 4 and groups 1, 2, 13, 16, and 17 can be seen in the following diagram.

Due to the complexity of accurately measuring atomic and ionic radii described above, you may see slightly different values than those in this explainer depending on what source you use. However, the important thing here is the trend that these values display.

Next, we will look for any periodic trends in the melting points of the period 3 elements. The following bar chart displays the melting points of the group 1 and 2 elements and the group 13 to group 18 elements. The orange bars represent period 2 elements, and the blue and gray bars represent period 3 and period 4 elements.

Carbon does not have a recorded melting point because it transforms from a solid phase to a gaseous phase when it is heated to a very high temperature.

The bar chart shows that period 2 and period 3 melting points increase as we move from group 1 elements to group 13 elements. This is due to an increase in the strength of metallic bonding. The number of valence electrons increases from one to three as we move from the group 1 elements to the group 13 elements. The metallic bonding becomes stronger when there are more delocalized electrons that can interact with the lattice of positively charged metal ions. Metal elements have stronger metallic bonding and higher melting points if they have more delocalized electrons.

The bar chart also shows that the period 2 and period 3 group 16, 17, and 18 nonmetal elements all have low melting points. The melting points are low because the nonmetal molecules are linked together with nothing more than weak van der Waal forces. It does not take a lot of thermal energy to break the molecules apart from each other.

Example 3: Explaining Why the Melting Point of the Period Three Metals Increases from Group 1 to Group 13

Which of the following statements does not in part explain why the melting point of the period 3 elements increases from Na to Al?

  1. From Na to Al, the charge on the metal ion decreases from +3 to +1.
  2. The number of delocalized electrons increases from Na to Al
  3. The strength of the metallic bonding increases.
  4. All three elements are metals, so they exhibit metallic bonding.

Answer

This question is asking us to identify which of these four statements does not contribute to an explanation of the periodic trend related to melting points of the period three metals.

Answer A states that as we move from sodium to aluminum, the charge on the metal ion decreases. However, group one metals only have one valence electron and typically form 1+ ions, whereas group 13 metals have three valence electrons and typically form 3+ ions. And so, it appears at first inspection that this answer is factually incorrect and likely the right answer to the question. We will examine the other options to be sure.

Answer B states that the number of delocalized electrons increases from sodium to aluminum, and this is in fact true. The increased number of delocalized electrons results in stronger metallic bonds, as mentioned in answer C, which leads to a higher melting point. And so, these answers cannot be correct.

Our final option, answer D, states that all three metals exhibit metallic bonding. Again, this statement is factually correct and so is an incorrect answer for this question. As we initially thought, answer A is the correct answer.

The final periodic trend we will examine in this explainer is the trend of electrical conductivity values. The number of free or delocalized electrons increases as we move from group 1 to group 13 elements. The electrical conductivity values increase at the same time, because delocalized electrons are highly mobile particles that can act as charge carriers. The following table shows how electrical conductivity values change across the third row of the periodic table. The electrical conductivity values are displayed with the siemens per metre (S/m) unit. The electrical conductivity values clearly increase from sodium to aluminum.

Period 3 ElementNaMgAlSiPSClAr
Electrical Conductivity (S/m)0.2180.2240.3822×101×101×10

As nonmetals, chlorine and argon have no free electrons to carry charge and are not considered to conduct electricity at all. In addition, sulfur and phosphorus conduct only a very small amount of electricity. The table shows that silicon can conduct electricity somewhat effectively, but it should be remembered that silicon is a semiconductor. Semiconductors conduct electricity differently at low and high temperatures. Semiconductors can conduct electricity effectively at high temperatures, but they act as insulators at low temperatures.

Definition: Semiconductor

Semiconductors are materials that act as insulators at low temperatures and conductors at higher temperatures.

Example 4: Matching the Periodic Trends of Period Three Elements to a Sketch Graph

A student is learning about periodicity trends of period 3 elements. She draws a sketch graph to show a trend of a property across the period but forgets to label the 𝑦-axis.

What label best fits the 𝑦-axis?

  1. Melting point
  2. Atomic radii
  3. Ionic radii
  4. Boiling point
  5. Electrical conductivity

Answer

In this question, we need an understanding of the different periodic trends that affect the elements of period 3. As we move across period 3, the melting point increases; however, this periodic trend continues up to and including silicon, and the graph in the question dips at silicon, so the answer cannot be A or by extension D.

In terms of atomic radius, atomic radius decreases across each period of the periodic table due to increasing nuclear charge, and so we would expect a graph of that periodic trend to have a linear, downward, or negative slope. Consequently, answer B is incorrect.

The ionic radius of the period 3 metals decreases due to the formation of positive ions with aluminum losing one more electron than magnesium which loses one more electron than sodium. As these atoms lose electrons, nuclear charge increases from sodium to aluminum, and these two effects combined result in the ionic radius decreasing. Just from examining the first three elements in this period, we can eliminate answer C.

By this stage, we have already identified our correct answer as answer E through a process of elimination, however, it is always good practice to check. As we move from left to right, initially extra electrons are delocalized in the metals, which allows greater current to flow and increases conductivity as we move from sodium to aluminum. In group 14, silicon is a semiconductor, and whilst it shows some conductivity, it is much less than the metals. From group 15 through to group 18, we are in the territory of the nonmetals: well-known insulators that cannot conduct electricity in any meaningful way. All of this is displayed on the sketch graph above, and together with what we already knew, answer E is the correct answer.

Key Points

  • Covalent atomic radius is defined as one-half the distance between the nuclei of two atoms covalently bonded together.
  • As we move across period 3, from group 1 to group 18, the atomic radius decreases.
  • The atomic radius increases as we move down each group of the periodic table.
  • As we move from group one to group 13 in period 3, the ionic radius decreases.
  • As we move from group 15 to group 17 in period 3, the ionic radius decreases.
  • Periodic trends in melting points are difficult to comprehensively classify, however, minor periodic trends in melting points are present.
  • Generally, melting points increase from group 1 to group 13 in period 2 and period 3 due to an increasing strength of metallic bonding.
  • The melting points of nonmetals are low because there are only weak van der Waal forces between the nonmetal molecules.
  • Conductivity increases from group 1 to group 13 in period 2 and period 3 due to an increased number of valence electrons.

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