ASP - Periodic Trends (Lesson)

Periodic Trends

When comparing certain chemical properties of elements it has long been noticed that some of these properties exhibit certain trends across the periodic table.  In this lesson trends in the following properties will be investigated: effective nuclear charge, electronegativity, ionization energy, electron affinity, atomic radius, melting point, and metallic character.

Periodic trends, arising from the arrangement of the periodic table, provide chemists with an invaluable tool to quickly predict an element's properties. These trends exist because of the similar atomic structure of the elements within their respective group families or periods, and because of the periodic nature of the elements. This is helpful in making predictions about the behavior and properties of an element when there is no data available.

Effective Nuclear Charge (Z_eff)

Effective nuclear charge can best be described as the force of attraction between the nucleus of an atom and its valence shell electrons.  To illustrate this, consider an atom of sodium that has an abbreviated electron configuration of [Ne]3s¹.  In this case, the effective nuclear charge would be the force of attraction between the positively charged nucleus and the single valence electron in the 3s orbital.  The reason that the 3s electron does not feel the full +11 charge of the sodium nucleus is because sitting in between the valence shell and the nucleus are the Ne core shell electrons.  The valence electron is simultaneously attracted to the nucleus and repelled by this core shell.  The core shell is described as shielding the valence electrons from this positive charge.  The net result is that the valence shell does not feel the full charge of the nucleus but rather an effective nuclear charge (Z_eff).

As we move from left to right on the periodic table the valence shell remains the same (e.g. as we move from sodium to magnesium the extra electron being added is being added to the same valence electron shell); however, the number of protons in the nucleus increases.  At the same time, the core shell remains exactly the same.  The final result is that the shielding electrons remain the same, but the charge of the nucleus increases.  For this reason, the effective nuclear charge increases from left to right on the periodic table.

When groups (columns) of elements are considered it can be seen that we not only adding protons to the nucleus, but also additional core shells of electrons.  The trend in effective nuclear charge can be difficult to predict, but it has been shown that the extra core electrons do a poor job of shielding the nucleus.  Therefore, the effective nuclear charge increases SLIGHTLY from the top of a group to the bottom.

Atomic Radius Trends

There are two ways to think of and measure the atomic radius of an atom.  The first is defined as half the distance between the nuclei of two atoms as shown below and is known as the bonding radius.  The second way is to simply think of the atom as being spherical in which case the atomic radius is simply defined as the distance between the nucleus and the edge of the valence shell.  This would then be the non-bonding radius.

Atomic size gradually decreases from left to right across a period of elements. This is due to the increase in the effective nuclear charge as described above.  Moving down a group it can be noticed that the atomic radius increases.  The reason for this cannot be attributed to effective nuclear charge, however, due to the fact that effective nuclear charge actually slightly increases down a group.  The reason for the top-to-bottom trend can be accounted for by considering the addition of electron core shells.  As more and more core shells of electrons are added, it pushes the valence shell further and further away from the nucleus resulting in increased atomic radii. 

Ionic Radius

• The ionic radius is defined as the distance between the nucleus and the electron in the outermost shell of an ion.
• When an atom loses electrons to form cations, the entire valence shell is lost.  This loss of valence shell causes the resulting cation to have a smaller radius.
• When an electron is added to an atom, forming an anion, the added electron increases the electron-electron repulsive forces, resulting in an increase in the size of the anion.
• To summarize:  cations are always smaller than their neutral counterpart whereas anions are always larger than their neutral counterpart.  

 

Ionization Energy

Ionization energy is the energy required to remove an electron from a neutral atom in its gaseous phase. A lower ionization value indicates that it is easier to remove an electron from an atom of that particular element and vice versa.  Generally, elements on the right side of the periodic table have higher IE values, measured in kilojoules per mole (kJ/mol) of atoms. This is because the electron arrangement in those atoms has the electron levels close to being filled and held closer to the positively charged nucleus than other atoms in the same row, called a period on the periodic table. 

This can be expressed by an equation:

X(g) + energy ⟶ X⁺(g) + e⁻

Where X is the neutral atom in gas phase and X⁺ the cation formed when the energy is added.  

IE Trends:

• The ionization energy of the elements within a period generally increases from left to right. This is due to an increase in effective nuclear charge making it more difficult to remove an electron due to the greater attractive force between the nucleus and the valence shell.
• The ionization energy of the elements within a group generally decreases from top to bottom and is due to an increase in atomic radius.  It is easier to remove electrons that are farther away from the nucleus.

What Do Multiple IE Values Mean?

It is possible to remove more than one electron from an atom resulting in several ionization energies; these varying energies are referred to as the first ionization energy, the second ionization energy, third ionization energy, etc. The first ionization energy is the energy required to remove the outermost, or highest, energy electron, the second ionization energy is the energy required to remove any subsequent high-energy electron from a gaseous cation, etc.

Electron Affinity

Electron affinity is the ability of an atom to accept an electron. Electron affinity is a quantitative measurement of the energy change that occurs when an electron is added to a neutral gas atom. The more negative the electron affinity value, the higher an atom's affinity for electrons.

Electron affinity generally decreases down a group of elements because each atom is larger than the atom above it (see atomic radius trend discussed above). This means that an added electron is further away from the atom's nucleus compared with its position in the smaller atom. With a larger distance between the negatively-charged electron and the positively-charged nucleus, the force of attraction is relatively weaker. Therefore, electron affinity decreases. Moving from left to right across a period, atoms become smaller as the forces of attraction become stronger. This causes the electron to move closer to the nucleus, thus increasing the electron affinity from left to right across a period.

EN Trends:

• Electron affinity increases from left to right within a period caused by the decrease in atomic radius.
• Electron affinity decreases from top to bottom within a group caused by the increase in atomic radius.

Electronegativity

Electronegativity can be understood as a chemical property describing an atom's ability to attract and bind with electrons. Electron Affinity is a quantitative property - energy can be measured. Electronegativity, on the other hand, is a qualitative property, so there is no standardized method for calculating electronegativity. 

Electronegativity measures an atom's tendency to attract and form bonds with electrons. This property exists due to the electronic configuration of atoms. Most atoms follow the octet rule (having the valence, or outer shell comprised of 8 electrons). Because elements on the left side of the periodic table have less than a half-full valence level, the energy required to gain electrons is significantly higher compared with the energy required to lose electrons. As a result, the elements on the left side of the periodic table generally lose electrons when forming bonds. Conversely, elements on the right side of the periodic table are more energy-efficient in gaining electrons to create a complete valence level of 8 electrons. Thinking about this in terms of electronegativity: the more inclined an atom is to gain electrons, the more likely that atom will pull electrons toward itself.

EN Trends:

• From left to right across a period of elements, electronegativity increases. If the valence shell of an atom is less than half full, it requires less energy to lose an electron than to gain one. Conversely, if the valence shell is more than half full, it is easier to pull an electron into the valence shell than to donate one.
• From top to bottom down a group, electronegativity decreases. This is because atomic number increases down a group, and thus there is an increased distance between the valence electrons and nucleus, or a greater atomic radius.
• Important exceptions to the above rules include the noble gases, lanthanides, and actinides. The noble gases possess a complete valence shell and do not usually attract electrons. The lanthanides and actinides possess more complicated chemistry that does not generally follow any trends. Therefore, noble gases, lanthanides, and actinides do not have electronegativity values.
• As for the transition metals, although they have electronegativity values, there is little variance among them across the period and up and down a group. This is because their metallic properties affect their ability to attract electrons as easily as the other elements.

According to these two general trends, the most electronegative element is fluorine, with 3.98 Pauling units.

 

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