For each of the following pairs, indicate which element has the smaller first ionization energy: (a) Cs, Cl (b) Fe, Zn (c) I, Cl (d) Se, Sn.

Ionization energy is the energy required to remove an electron from an atom in its gaseous state. It is a key factor in determining an element's reactivity and is influenced by the atomic size and the effective nuclear charge. Generally, ionization energy increases across a period and decreases down a group in the periodic table.

Periodic trends refer to the predictable patterns observed in the properties of elements as you move across or down the periodic table. For ionization energy, elements on the right side of the table tend to have higher ionization energies than those on the left, while elements higher up in a group have higher ionization energies than those lower down due to increased electron shielding and distance from the nucleus.

Electron configuration describes the distribution of electrons in an atom's orbitals. The arrangement of electrons affects an atom's ionization energy; for example, elements with a stable electron configuration (like noble gases) have higher ionization energies. Understanding the electron configuration of the elements in the pairs helps predict which will have a lower ionization energy based on their stability and reactivity.