Chem: Trends Of The Periodic Table

Effective nuclear charge

The net positive charge from the nucleus that an electron can “feel” attractions from. The core electrons are said to shield the valence electrons from the full attractive forces of the protons in the nucleus. 

Shielding

 core (nonvalence) electrons shield the valence electrons from the full attractive forces of the protons in the nucleus. 

Electron-electron repulsions

due to their like charges, electron pairs orient themselves as far away as possible from each other, causing the electron cloud to expand (justifies trends across a period.

Atomic radius

Atomic radius is the distance from the atom’s nucleus to the outer edge of the electron cloud. 

In general, atomic radius decreases across a period and increases down a group. 

Atomic radius across the period

Across a period, effective nuclear charge increases as electron shielding remains constant. A higher effective nuclear charge causes greater attractions to the electrons, pulling the electron cloud closer to the nucleus which results in a smaller atomic radius. 

Atomic Radius: Down the group

Down a group, the number of energy levels increases, so there is a greater distance between the nucleus and the outermost orbital. This results in a larger atomic radius

Ionic radius

The distance from the nucleus to the outer edge of the electron cloud of an ion. 

The same trend of atomic radius applies once you divide the table into metal and nonmetal sections. 

Ionic Radius Cation

A cation has a smaller radius than its neutral atom because it loses valence electrons. The “new” valence shell is held closer to the nucleus, resulting in a smaller radius for the cation.

Ionic Radius Anion

An anion has a larger radius than the neutral atom because it gains valence electrons. There are added electron/electron repulsions in the valence shell that expand the size of the electron cloud, which results in a larger radius for the anion.

Ionization energy

(IE) is the energy required to remove the highest-energy electron from a neutral atom. 

In general, ionization energy increases across a period and decreases down a group. 

Ionization energy across the period

Across a period, effective nuclear charge increases as electron shielding remains constant. This pulls the electron cloud closer to the nucleus, strengthening the nuclear attraction to the outermost electron, and is more difficult to remove (requires more energy).

Ionization energy down the group

Down a group, the number of energy levels increase and the distance is greater between  the nucleus and highest-energy electron. The increased distance weakens the nuclear attraction to the outermost electron, and is easier to remove (requires less energy).

Electronegativity

Electronegativity is the measure of the ability of an atom in a bond to attract electrons to itself. 

Electronegativity increases across a period and decreases down a group.

Electronegativity on left and right side

Towards the left of the table, valence shells are less than half full, so these atoms (metals) tend to lose electrons and have low electronegativity. Towards the right of the table, valence shells are more than half full, so these atoms (nonmetals) tend to gain electrons and have high electronegativity. 

Electronegativity down the group

Down a group, the number of energy levels (n) increases, and so does the distance between  the nucleus and the outermost orbital. The increased distance and the increased  shielding weaken the nuclear attraction, and so an atom can’t attract electrons as strongly. Fluorine is the most electronegative element, whereas francium is the least  electronegative element. 

Electron affinity

Electron affinity is the energy change that occurs when an atom in the gas phase gains an electron to form a negatively charged ion.

Electron affinity increases across a period and decreases down a group

Electron affinity across the period

As you move across a period, atoms have more protons in the nucleus, increasing the effective nuclear charge, which attracts electrons more strongly, so energy is released when an electron is added. Additionally, atoms are closer to filling their valence shells, so gaining an electron is more energetically favorable.

Electron affinity down the group

As you move down a group, the added electron is farther from the nucleus due to larger atomic size, and electron shielding reduces the attraction, so less energy is released when an electron is gained. Also, the increased distance and repulsion from inner electrons make it harder for the atom to accept an extra electron.