## Periodic Table and Periodic Trends

Published: Dec 27th, 2020

The periodic table is among the most important tools you can possibly have in chemistry. It is a representation of all known elements. The elements are organized according to their properties into a shape you might find familiar. By knowing why elements are arranged as they are, we can predict the properties of an unknown element.

The diagram below is of the periodic table. We will refer back to this table often throughout this course, so get acquainted! Here is a direct link to a PDF of the table:

Some important elements of the periodic table include:

1. Hydrogen: an outlier! It has some metal-like properties and some non-metal properties
2. Alkali metals: These are soft, silver-grey metals that are solids are standard ambient room temperature (SATP; 25°C and 1 atm pressure) and exhibit metallic properties (e.g., conducting electricity). Alkali metals react with water explosively to form hydrogen gas. Alkali metals also oxidize in air, which is why they are often stored in oil.
3. Alkaline earth metals: These are light, reactive metals that also exhibit metallic properties. Like the alkali metals, these release hydrogen when they react with water. Alkaline earth metals are also solids at SATP.
4. Transition metals: These are the more typical metals. They are good electrical conductors, they are often multivalent (can lose electrons from more than one valence shell), and they exhibit a wide range of chemical and physical properties.
5. Halogens: May be solids, liquids, or gases at room temperature and are very reactive. They do not conduct electricity and are often dull in colour
6. Noble gases: These are the most unreactive/inert elements. Most of these elements are gases at SATP; they have low melting and boiling points.
7. You might also notice a staircase shaped line. The elements adjacent to this staircase are called the metalloids, as they exhibit both metal and non-metal properties.

In addition to these groups, the periodic table is organized into rows and columns, called periods and groups, respectively. There are 7 periods and 18 groups. The alkali metals, for example, are group 1 of the table.

In the last lesson, we discussed Bohr-Rutherford diagrams, a representation of atomic nuclei. Above is a diagram showing the Bohr-Rutherford configurations of the first 18 elements. Looking at each column/group, what do you notice about the number of electrons in the outermost shell / valence shell of each atom?

They all have the same number of electrons, called valence electrons. One factor that contributes to the similar reactivity among elements in a group is that they all share the same number of valence electrons. Lewis diagrams are a type of atomic representation that only show the number of valence electrons surrounding an element. For example, sodium would be represented as $$Na:$$, with 2 valence electrons surrounding the atom.

Now that we’ve become acquainted with the periodic table, let’s go over some periodic trends, shall we?

The radius of an atom is defined as the distance from the nucleus to the outermost electrons. If an atom is diatomic (e.g., nitrogen usually exists as diatomic N2), the radius is defined as the distance between the two nuclei, divided by two.

In general, as we move from left to right along the table, the atomic radius decreases. Why? Well, look back at the Bohr diagrams. In a given period, all the atoms have the same number of valence shells. However, as we move from left to right, the number of electrons in those valence shells increases. Recall that electrons are negatively charged particles. This is in contrast to the nucleus, which contains positively charged protons. Therefore, as we increase the number of electrons, there’s a greater negative force being pulled on by the protons, so the valence shell is brought closer to the nucleus, shrinking the atomic radius. The force the nucleus exerts on each electron in an atom is called the effective nuclear charge.

When an atom gains or loses electrons, it becomes an ion. For example, sodium in solution dissociates from its crystalline form and becomes Na+. The radius of the ion is called its ionic radius, and the convention for measurement is the same. A positive ion is called a cation, and it results from losing electrons. A negative ion is called an anion, and it results from gaining electrons (we can’t gain or lose protons because otherwise the atom would no longer be the same atom anymore!).

In general, an atom is smaller than its corresponding anion and larger than its corresponding cation. Why? It has to do with attraction and repulsion. When a chlorine atom gains an electron, it becomes an anion. Although it still has the same number of protons, it now has a greater number of electrons. Therefore, the amount of positive charge is constant, while the amount of negative charge has increased. This causes the electrons to repel each other, causing ionic radius to increase relative to the atom.

What about the cation? In this case, we’re holding the number of protons constant but decreasing the number of electrons. Therefore, there is more positive charge relative to the number of electrons, so the valence electrons are pulled on more tightly, decreasing the ionic radius. You might be wondering why the protons don’t repel like the electrons.

Ionization Energy

Ionization energy refers to how easily an atom is willing to give up an electron to become an ion. Moving from left to right along the table, the general trend is that the ionization energies increase. Moving down the table, the ionization energies decrease. Why?

Just as for ionic radius, moving left to right along the table causes an increase in the number of electrons while holding the protons constant. There is therefore a greater effective nuclear charge, so the electrons are being tightly held by the protons. Therefore, there is more energy required to pull off an electron, so there is a higher ionization energy

Down the table, the ionization energy decreases. This is because although the number of protons is increasing, the valence shells are getting farther and farther away from the nucleus. Therefore, the effective nuclear charge is lesser for the valence electrons. This makes it much easier to pull off an electron, which is reflected by the lower ionization energy.

Electron Affinity

Electron affinity is essentially the opposite of ionization energy; it is a measure of how much energy an atom is willing to “spend” to gain another electron. As you might have learned in past classes, the most stable configuration of electrons is that of the noble gases, since their valence shells are full. Therefore, all of the other atoms are “jealous,” and want a similar configuration for themselves.

Going from left to right across the table, the electron affinity increases (not including the noble gases, of course). Halogens, for example, only need to gain one more electron to get a noble configuration, so they’re willing to spend a LOT of energy to become stable. On the other hand, sodium would have to gain seven electrons to get the noble configuration, so it’s not willing to spend a lot of energy. It’s much easier for sodium to simply lose an electron (hence the low ionization energy) to get a noble configuration.

The following diagram summarizes these periodic trends across the periodic table:

Conclusion

In this lesson, we’ve discussed the periodic table and have observed some periodic trends. Be sure to try some practice problems!