Mendeleev understood the ‘Periodic Law’ which states: When arranged by increasing atomic number, the chemical elements display a regular and repeating pattern of chemical and physical properties.

Atoms with similar properties appear in groups or families (vertical columns) on the periodic table. They are similar because they all have the same number of valence (outer shell) electrons, which governs their chemical behavior.

There are two aspects about how atomic structure changes from atom to atom that impact/cause these trends: 1. effective nuclear charge 2. shielding

What keeps electrons from simply flying off into space? Effective nuclear charge is the pull that an electron “feels” from the nucleus. The closer an electron is to the nucleus, the more pull it feels. As effective nuclear charge increases, the electron cloud is pulled in tighter.

As more energy levels are added to atoms, the inner layers of electrons shield the outer electrons from the nucleus. The effective nuclear charge on those outer electrons is less, and so the outer electrons are less tightly held.

Radius decreases across a period Increased effective nuclear charge due to addition of protons Radius increases down a group Each row adds another “shell” or energy level to the atom so valence electrons are that much further from nucleus

Tends to increase across a period As radius decreases across a period, the electron you are removing is closer to the nucleus and harder to remove. Tends to decrease down a group Outer electrons are farther from the nucleus and easier to remove.

If an electron is given enough energy (in the form of a photon) to overcome the effective nuclear charge holding the electron in the cloud, it can leave the atom completely. The atom has been “ionized” or charged. The number of protons and electrons is no longer equal. The larger the atom is, the easier its electrons are to remove. Ionization energy and atomic radius are inversely proportional.

What does the word ‘affinity’ mean? Electron affinity is the energy change that occurs when an atom gains an electron (also measured in kJ). This is directly related to Ionization Energy; the more energy it takes to remove an electron, the more the atom would rather gain an electron.

Electronegativity tends to increase across a period As radius decreases, electrons get closer to the bonding atom’s nucleus so the nuclear charge exerts more influence Electronegativity tends to decrease down a group or remain the same As radius increases, electrons are farther from the bonding atom’s nucleus so its nuclear charge has less influence

It is an arbitrary scale that ranges from 0 to 4. The units of electronegativity are Paulings. Generally, metals are electron givers and have low electronegativities. Nonmetals are electron takers and have high electronegativities.

Positively charged ions form when an atom of a metal loses one or more electrons Smaller than the corresponding atom Negatively charged ions form when an atom of a nonmetal gains one or more electrons Larger than the corresponding atom Cations Anions

Graphic courtesy Wikimedia Commons user Popnose

This ties all the previous trends together in one package. However, we must treat metals and nonmetals separately. The most reactive metals are the largest since they are the best electron givers. The most reactive nonmetals are the smallest ones, the best electron takers.

Shielding is constant Atomic Radius decreases Ionization energy increases Electronegativity increases Nuclear charge increases Shielding increasesAtomic radius increasesIonic size increasesIonization energy decreasesElectronegativity decreases Nonmetallic character increases Metallic character increases