Periodic Trends

Every trend on the periodic table traces back to one fight: protons pulling electrons in versus inner electrons shielding that pull. Move left to right across a period and protons increase one at a time, but new electrons go into the same outer shell, which barely shields at all. So effective nuclear charge, the net positive pull an outer electron actually feels, climbs steadily. That stronger pull drags the electron cloud inward, so atoms shrink across a period. Move down a group instead, and each row adds a whole new shell farther from the nucleus, with inner shells shielding the added protons almost completely. Size wins out over charge, so atoms get bigger going down.

Ionization energy, the energy needed to rip off an outer electron, is just the flip side of that same pull. Tighter grip means higher ionization energy, so it rises across a period and falls down a group, mirroring radius exactly. But the climb across a period is not perfectly smooth. Boron dips below beryllium because boron's lone 2p electron sits in a higher-energy, less-shielded orbital than beryllium's full 2s pair. Oxygen dips below nitrogen because nitrogen has three 2p electrons, one in each orbital, unpaired and stable, while oxygen must pair a fourth electron into an already-occupied orbital, and that pairing repulsion makes it easier to remove.

Electronegativity, the pull an atom exerts on shared electrons in a bond, follows ionization energy's logic: it rises across a period and falls down a group, so fluorine sits at the top of the trend and francium near the bottom. These three trends set up everything that follows, including how atoms form ions and bonds in the next units.