Q.The increasing order of reactivity among group 1 elements is Li < Na < K < Rb < Cs whereas that among group 17 elements is F > Cl > Br > I. Explain.

Concept: Periodic trends in metallic and non-metallic character

The reactivity pattern reverses between groups 1 and 17 because their chemistry is fundamentally opposite—alkali metals lose electrons while halogens gain them.

Group 1 (Alkali Metals): Reactivity increases down the group because the valence electron becomes easier to remove. As atomic size increases from Li to Cs, the outermost electron is farther from the nucleus and more shielded by inner shells. The ionization energy drops sharply ($\text{IE}_\text{Li} = 520 \text{ kJ/mol}$ vs. $\text{IE}_\text{Cs} = 376 \text{ kJ/mol}$), so cesium loses its electron most readily and reacts most vigorously. …

The heart of this question lies in understanding what drives chemical reactivity for metals versus non-metals, and how atomic structure changes as you descend a group in the periodic table.

Group 1 elements (alkali metals) react by losing their single valence electron to form $\ce{M+}$ cations. The easier it is to remove that electron—the lower the ionization energy—the more reactive the metal. Group 17 elements (halogens) react by gaining an electron to form $\ce{X-}$ anions. The more readily they accept that electron—the more negative the electron affinity (or the stronger the attraction for the incoming electron)—the more reactive the non-metal.

As we move down any group, several periodic trends come into play:

1. Atomic size increases. Each successive element adds a new electron shell, pushing the valence electrons farther from the nucleus.

2. Shielding increases. Inner-shell electrons repel the outer electrons, reducing the effective nuclear charge ($Z_{\text{eff}}$) felt by the valence electrons.

3. Ionization energy decreases. The combination of larger radius and greater shielding means the outermost electron is held less tightly, making it easier to remove.

4. Electron affinity becomes less negative (weaker). A larger atom has its valence shell farther from the nucleus, so the attraction for an incoming electron diminishes. The added electron experiences more repulsion from inner shells and less pull from the nucleus.

Now let's apply these trends to each group.

Group 1: Li < Na < K < Rb < Cs (Increasing Reactivity)

For alkali metals, reactivity hinges on how easily the atom surrenders its lone $ns^1$ electron.

1. Lithium has the smallest atomic radius in the group. Its $2s^1$ electron is close to the nucleus and experiences a relatively high $Z_{\text{eff}}$ (minimal shielding from just the $1s^2$ core). Ionization energy is highest here, so Li is the least reactive.

2. Sodium adds a third shell. The $3s^1$ electron is farther out, shielded by the $1s^2 2s^2 2p^6$ core. Ionization energy drops, reactivity increases.

3. Potassium, Rubidium, Cesium continue the trend. Each additional shell increases the distance and shielding. By the time we reach Cs, the $6s^1$ electron is so loosely held that cesium reacts explosively with water—it's the most reactive stable alkali metal.

The metallic character—the tendency to lose electrons—increases down the group.

Group 17: F > Cl > Br > I (Decreasing Reactivity)

Halogens react by accepting an electron into their $np^5$ orbital to complete the octet.

1. Fluorine is the smallest halogen. The incoming electron enters the $2p$ subshell, very close to a highly charged nucleus ($Z = 9$) with minimal shielding (only $1s^2$ core). The electron affinity is extremely negative (though not the most negative in the group due to electron-electron repulsion in the tiny $2p$ orbital, this is a subtle exception). More importantly, F has the highest electronegativity and forms the strongest bonds—it's the most reactive non-metal. …