Periodic trends are properties of the elements of the periodic table which are based on the structure of the atoms that make up each element. The periodic table was designed with these trends in mind, so it is always helpful to refer to the periodic table when comparing periodic trends of certain atoms.

Two important concepts to understand for periodic trends are Coulomb's Law and effective nuclear charge. Coulomb's Law states that two charged particles apply a force on each other, and this force is greatest when the charges are close together. In an atom, this means that the positively charged nucleus has less of an attraction on electrons that are farther away, and a greater attraction on electrons that are closer. Effective nuclear charge states that internal electrons essentially "block" the charge of the nucleus from reaching the electrons, so an atom with more internal electrons will have valence electrons that experience a lesser effective nuclear charge. An increase in the number of protons in an atom without changing the number of internal electrons (like moving to the right across a period) increases the effective nuclear charge on the valence electrons.

Listed below are examples of periodic trends, how they change throughout the periodic table, and why the structure of the atom results in this trend.

Atomic size: the measure of how large an individual atom is

• Increases down a group as more electron shells are added which become farther away from the nucleus

• Increases to the left across a period since the increased number of protons in the nucleus moving to the right results in a greater effective nuclear charge on the electrons which pulls them closer to the nucleus

Ionic radius: the size of an ion of an element

• Cations (positive charge) become smaller since they lose an electron (or more) to ionize

• Anions (negative charge) become larger since they gain an electron (or more) to ionize.

Electronegativity: the pull an atom applies on a shared electron in a compound

• Increases up a group and to the right across a period (except for the noble gasses) as Fluorine is the most electronegative element

Electron affinity: the change in energy when an atom accepts another electron

• Higher electron affinity means an atom wants an electron, lower electron affinity means an atom does not want an electron.

• Increases to the right as elements are closer to forming an octet or filling their valence shell of electrons

• No clear pattern is seen with electron affinity going down a group as the value varies throughout the periodic table

Ionization energy: the force that needs to be exerted in order to remove an electron

• Increases to the right across periods as more protons in the nucleus result in a higher effective nuclear charge on the electrons that attracts them with a stronger force and makes the electrons more difficult to remove

• Increases up groups as electrons get farther away in higher energy levels and Coulomb's Law says that the attractive force is reduced. Atoms higher in their groups have fewer energy levels, so the electrons are closer to the nucleus and are harder to remove

• Exceptions include Oxygen and Boron as half-filled or fully-filled orbitals are more stable than atoms where there is a single extra election present

• Both first ionization energy and subsequent ionization energies can be measured, and increased levels of ionization energy are always larger. As more electrons are removed, there is less repulsion between the electrons so they are harder to take away from the atom.

UV/Vis spectroscopy