AT_Periodic Table Trends Lesson

Periodic Table Trends 

There is much more that the periodic table can tell us. Many chemical and physical properties vary in a systematic way and we can therefore make generalized statements about them based on their position on the periodic table. Recall where the metals and nonmetals are located on the table. We can discuss their activity based on their positions. The most reactive metals are found at the bottom, left of the table. This activity (also called metallic character) decreases as you move up and to the right. Conversely, the most reactive nonmetals are found at the top right, not including the noble gases. Because the noble gases are already stable and do not easily react, they are often excluded from the trends.

Effective Nuclear Charge

The first trend we will discuss in depth is effective nuclear charge. Understanding this trend will help you to explain many of the other trends. Effective nuclear charge (Zeff) is the positive charge "felt" by the valence electrons. This positive charge is always less than the full nuclear charge (with the exception of H that has only 1 electron) because the negative charge of the inner electrons partially cancels out the positive charge of the nucleus.

Effective Nuclear Charge of Lithium Start with the electron configuration for Lithium (Li) 1s2 2s¹ Diagram 1 The nucleus has a +3 charge because Li has 3 protons (shown in red). The inner, or core electrons (shown in green) have a charge of -2. (Diagram 1) The position of the inner and outer electrons creates an effect called shielding. (Diagram 2). We say that the valence electrons are shielded from the pull of the nucleus by the inner electrons. (shown in yellow) Diagram 2 If you consider the charges inside the shield, you have an overall +1 = 3 protons and 2 electrons. (Diagram 3). So the Zeff "felt" by the outer, or valence, electron (shown in blue) is +1. (Diagram 4)

Inner electrons effectively shield outer electrons from nuclear charge. However, it is important to realize that electrons in the same shell (or energy level) are not very effective at shielding each other since they are about the same average distance from the nucleus.

ZEFF EQUATION Effective Nuclear Charge -Zer=Z-0 eff Number of inner (core) electrons This is sometimes referred to as the shielding constant. Atomic Number Number of Protons

Atomic Radius

Atomic Radius =r An atom does not have a defined edge. This makes calculating its radius difficult. Instead, the radius is calculated as half the distance between the nuclei of two adjacent atoms.Remember that in many ways electrons behave like waves, not particles. We have to continually remind ourselves of this because it is so easy to think of an electron as a particle. Because of this wave nature of electrons, it is difficult to define atomic "size".  

Since there isn't a defined edge, we say that atomic radius is half the distance between the nuclei of two adjacent atoms.

The diagram below is a representation of atomic radii. As you study this diagram, look for a general trend of increasing or decreasing across a period or group. There is, of course, variation here, but still try to look for a general trend.

periodic table of atomic radii

The general trend is that atomic radius decreases as you move from left to right across a period and increases as you move down in a group. This is easy to remember if you start from the bottom left and move up and to the right, atomic radius decreases.

In addition to knowing the trends, you need to be able to explain the reason for the trends. As you move from left to right across a period, nuclear charge (# of protons) increases while the # of inner electrons stays the same, therefore the Zeff increases to the right. This means that the valence electrons "feel" more of a pull from the nucleus, causing them to move in closer. This results in a smaller radius.

Within a group, the Zeff is nearly constant, but electrons are added to higher energy levels that are farther away from the nucleus. So, the radius is increased.

Ionic Radius

When atoms become ions, we see big changes in size.   We refer to this as ionic radius rather than atomic radius. When an atom gains an electron (becoming an ANION), it joins the exiting valence electrons. These electrons repel each other, causing them to push apart and the radius therefore increases.

Negative ions are always larger than their atoms.

When an atom loses an electron (becoming a CATION), the repulsion among the remaining valence electrons decreases and they spread out causing the radius to increase.

Positive ions are always smaller than their atoms.

table of ionic radii with ions are colored red and blue; parent atoms brown. Radii are in picometers.

 

Ionization Energy

Ionization energy (I.E.) is defined as the energy required to remove an electron (from an isolated, gaseous atom or ion in its ground state).

As expected, it is easier to remove an electron from a metal than a nonmetal. After all, that's what metals do. They lose electrons. So the trend is that I.E. increases as you move up a group and to the right across a period.

Here is another way to think about this. The larger an atom, the easier it is to pull an electron from its outermost shell. Large atoms have low ionization energies. Electrons that are held closer and tighter to the nucleus are harder to remove, and therefore have greater ionization energies. Notice that this trend is the opposite of the atomic radius trend.

I.E. specifically refers to the energy required to remove the first electron. If more electrons are removed, those energies are called 2nd I.E., 3rd, I.E., etc.  

Remember to work on the module practice problems as you complete each section of content.

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