Following statements regarding the periodic trends of chemical reactivity of the alkali metals and t

Q: Why does reactivity mean different things for Group 1 and Group 17?

Reactivity is simply a measure of how easily an element achieves a stable noble gas configuration. Group 1 elements have 1 valence electron, so the easiest path is to lose it (oxidation). Group 17 elements have 7 valence electrons, so their easiest path is to gain one (reduction). Thus, the mechanism of reactivity is fundamentally opposite.

Q: Doesn't a larger size make it easier for halogens to accommodate an extra electron?

While a larger electron cloud has less inter-electronic repulsion, the primary driving force for gaining an electron is the attractive pull from the positively charged nucleus. As size increases, the incoming electron is placed further from the nucleus and is shielded by more inner electrons, making the attractive force significantly weaker.

Q: Is Fluorine really more reactive than Chlorine? I thought Chlorine had a higher electron gain enthalpy.

Yes, Fluorine is the most reactive halogen and the strongest oxidizing agent. Although Chlorine releases more energy upon gaining an electron in the gaseous state (due to Fluorine's small size causing electron-electron repulsions), Fluorine's extremely low bond dissociation energy and high hydration enthalpy make it far more reactive overall in chemical reactions.

Q: Why do we only consider the outermost electron for alkali metals?

Alkali metals have a general electronic configuration of $ns^1$. Removing this single $s$ electron requires relatively low energy (first ionization enthalpy) and results in a highly stable noble gas core. Removing a second electron would require breaking into that stable core, which takes an enormous amount of energy, so they strictly react by losing just one.

Q: Do these trends hold true for all groups in the periodic table?

These are general trends that are most pronounced in the extreme left (Group 1/2) and extreme right (Group 16/17) of the periodic table. Transition metals (d-block) and p-block metalloids have more complex reactivity patterns due to variable oxidation states, d-orbital shielding, and the inert pair effect.

Question

Following statements regarding the periodic trends of chemical reactivity of the alkali metals and the halogens are given. Which of these statements give the correct picture?

Options:

(A) The reactivity decreases in the alkali metals but increases in the halogens with increase in atomic number down the group

(B) In both the alkali metals and the halogens the chemical reactivity decreases with increase in atomic number down the group

(C) Chemical reactivity increases with increase in atomic number down the group in both the alkali metals and halogens

(D) In alkali metals, the reactivity increases but in the halogens it decreases with increase in atomic number down the group

Sigma Percentile · Accepted Answer

### The Dual Nature of Chemical Reactivity

When we talk about "chemical reactivity" in the periodic table, we are not talking about a single, universal mechanism. Reactivity is a measure of how desperately an element wants to achieve a stable, noble gas electronic configuration. However, the path to that stability is fundamentally different depending on whether the element is a metal or a non-metal.

Metals, found on the left side of the periodic table, are electropositive. They achieve stability by **losing** electrons. Non-metals, found on the right side, are electronegative. They achieve stability by **gaining** electrons. This fundamental dichotomy is the key to understanding why reactivity trends behave the way they do as we move down a group.

### Alkali Metals: The Eager Electron Donors

Let's focus on Group 1, the alkali metals: Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), and Cesium (Cs). These elements all have a single valence electron in their outermost $s$-orbital ($ns^1$). Their sole chemical goal in life is to get rid of this electron to reveal the stable noble gas core underneath.

The energy required to rip this electron away is called the **Ionization Enthalpy (IE)**. As we move down Group 1, the principal quantum number ($n$) increases, meaning new electron shells are added. The atomic radius becomes significantly larger.

Because the outermost electron is now much further away from the positively charged nucleus, and because it is shielded by an increasing number of inner core electrons, the effective nuclear charge it feels is drastically reduced. Consequently, the ionization enthalpy drops.

$$IE \propto \frac{1}{	ext{Atomic Radius}}$$

Since it requires less and less energy to remove the valence electron as we go down the group, the metals become increasingly eager to donate it. Therefore, **the chemical reactivity of alkali metals increases down the group**.

### Halogens: The Greedy Electron Acceptors

Now, let's cross over to the other side of the periodic table to Group 17, the halogens: Fluorine (F), Chlorine (Cl), Bromine (Br), and Iodine (I). These elements have seven valence electrons ($ns^2 np^5$). They are agonizingly close to a full octet and are highly motivated to snatch just one more electron from any available source.

The tendency of an isolated gaseous atom to accept an electron is measured by its **Electron Gain Enthalpy ($\Delta_{eg}H$)** or, more broadly, its electronegativity.

As we move down Group 17, the atomic radius increases just as it did for the alkali metals. However, for halogens, a larger size is a severe disadvantage. When a halogen atom tries to capture a new electron, that incoming electron must be attracted by the nucleus. If the atom is large, the incoming electron is placed in a shell that is far from the nucleus and heavily shielded by inner electrons.

The attractive force from the nucleus is weak, meaning the atom's ability to pull in and hold onto an extra electron diminishes.

$$	ext{Electron Pull} \propto \frac{1}{	ext{Atomic Radius}}$$

Because their ability to gain an electron weakens as they get larger, **the chemical reactivity of halogens decreases down the group**.

### The Grand Conclusion

By comparing these two extreme groups, we see a beautiful symmetry in the periodic table. The reactivity of metals (driven by electron loss) is enhanced by a larger atomic size, while the reactivity of non-metals (driven by electron gain) is hindered by a larger atomic size.

Therefore, as atomic number increases down the respective groups, the reactivity of alkali metals increases, but the reactivity of halogens decreases. This perfectly aligns with the statement provided in option (d).

Animated Solution for Chemistry - Periodicity in Properties: Following statements regarding the periodic trends of chemical reactivity of the alkali metals and the halogens are given. Which of these statements give the correct picture?

Select Answer:

Visualized Solution

$G r o u p 1 & G r o u p 17$

$M \to M^{+} + e^{-}$

$I E ↓ ⟹ R e a c t i v i t y ↑$

$X + e^{-} \to X^{-}$

$E A ↓ ⟹ R e a c t i v i t y ↓$

$O pt i o n (d)$

The Sigma Insight: Periodic Table and Periodic Properties

The Dual Nature of Chemical Reactivity

Alkali Metals

Halogens

The Grand Conclusion

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