Periodic trends explain why ionic bonds form the way they do.
Why is K at the top? Down group 1 the atoms get bigger (Li 152 → Na 186 → K 231 pm), so the single outer electron sits farther from the nucleus and is held less tightly. First ionization enthalpy decreases down the group, so the electron comes off more easily and the metal is more reactive: K is above Na.
Across period 3 it runs the other way. Na (496 kJ/mol) loses its electron more easily than Mg (737 kJ/mol), so Na sits above Mg. Metals on the left and lower down lose electrons most easily, and that is the job a metal does in an ionic bond.
Series from NCERT Class 10 Chemistry. Dotted chips are in the table above; tap one to select it.
The energy needed to pull off the outermost electron rises across the period (Na 496 → Si 786). Al dips below Mg, which is a known exception: Al's outer electron is in a higher-energy p orbital. Al still has to lose three electrons to become Al3+, so it sits below Mg in the series.
kJ/mol, NCERT Class 11 (Problem 3.6). Group 1: decreases down the group, shown qualitatively only.
One continuous spectrum, not three boxes. No real bond is 100% ionic or 100% covalent: the bigger the electronegativity difference (ΔEN), the more ionic the bond. The dashed lines at ΔEN ≈ 0.5 and ≈ 1.7 are a general-chemistry rule of thumb used here for placing the marker. They are not numbers from NCERT, which describes the change qualitatively.
Reasoning: NCERT Class 11 Chemistry, Ch 4, section 4.3.6 (Polarity of bonds). Electronegativity: Pauling scale.
Rules and examples: NCERT Class 9 Chemistry, Ch 9, section 9.4.2.