Periodic Trends
Coulomb’s law of attraction states that like charges repel and opposite charges attract. Although it was created for electrostatics, it is still used here. It is as follows:
where F is the force of attraction, k is a constant roughly equal to 9 x 10⁹, q is the strength of the charge, and r is the distance between the charges. The force of attraction on an atomic level is called the effective nuclear charge (called the Z effective). The electrons are pulled towards the nucleus with the Z effective. The further the electrons are from the nucleus, the less Z effective they will have. As you move horizontally across the periodic table from left to right, the Z effective will increase due to an increased number of electrons. As you move down the periodic table, the Z effective will decrease due to a change in energy levels.
As you move horizontally from left to right on the periodic table, you are moving across what is called a period. When you move across a period, the amount of protons (charge: +1) and electrons (charge: -1) increases, creating a greater difference in charge. Each vertical column in the periodic table is called a family.
Electrons surround the nucleus in concentric circles called energy levels. The 1st energy level is closest to the nucleus and holds the least amount of electrons. The 2nd holds more electrons and is outside the 1st level. This pattern continues for 7 energy levels in total.
Atomic size (radius) is calculated using half the distance between the nuclei of two bonded atoms. So to determine the atomic radius of oxide, O, you would use half the distance between the nuclei of the oxides in O₂. It simply shows how large an atom is. It increases as you move down each family, due to an extra energy level for electrons to orbit. The extra energy level causes the electrons to repel each other further away, decreasing the Z effective, thereby increasing the radius. The atomic size also decreases as you move across each period, as the increase in electrons causes them to be pulled into the nucleus with a greater nuclear charge.
Ionic size relates the size of ions of metal atoms and ions of nonmetal atoms. When looking at any positive ion, it is called a cation, and any negative ion is called an anion. Cations coming from metal atoms are smaller than their neutral atom, whereas anions coming from nonmetal ions are larger than the neutral atom. When the atom becomes ionized, it either loses or gains an electron which accounts for the size difference. Since electrons are negatively charged, decreasing the charge by one adds one electron.
For two elements to be isoelectronic, they have to have the same amount of electrons. For example, C⁻¹ and N are isoelectronic with each other because they have the same number of electrons. This can be very useful because when two elements are isoelectronic, they can share properties and reactivity.
Ionization energy is the energy required to remove the most loosely bound electron when the atom is in its gaseous state. It can only remove a valence electron at a time (electron in the outermost energy level), as it would be impossible to remove an electron from an inner energy level with electrons still orbiting around it. As you move down a family, ionization energy decreases because the Z effective decreases, allowing electrons to be removed with less energy. Ionization energy increases as you move along the period, as the Z effective increases with more electrons.
Most elements have paired electrons. The ionization energy largely increases when trying to remove a paired electron, or an electron in a full energy level. The equation for ionization energy is:
X⁰(g) + energy -> X⁺¹(g) + e⁻
with X as any element and e⁻ as an electron.
Electron affinity is the exact opposite of the ionization energy. Rather than the energy required to remove an electron, electron affinity is the energy released when a gaseous atom gains an electron. However, it still follows the same pattern across the periodic table for the same reasons. The equation for electron affinity is
X⁰(g) + e⁻ -> X⁻¹(g) + energy
Electronegativity is the ability of an atom to pull electrons into a chemical bond. The pattern of electronegativity follows that of the Z effective, decreasing down a family and increasing across a period. The higher the electronegativity, the more an atom pulls electrons towards it, creating a bond.
