Chemistry · Ch 3 — Classification of Elements and Periodicity in Properties
Periodic Trends and Chemical Reactivity
Periodic Trends and Chemical Reactivity
3.7.3 Periodic Trends and Chemical Reactivity
The periodic table is not just an arrangement of elements — it is a map of chemical behaviour. Every physical and chemical property of an element ultimately stems from its electronic configuration. The trends we have studied in atomic radii, ionization enthalpy, electron gain enthalpy, and valence are not isolated facts; they work together to determine how an element reacts. This section connects those fundamental properties to chemical reactivity, showing why elements at the two ends of a period are the most reactive, and why reactivity changes as we move down a group.
Reactivity Across a Period: The Two Extremes
As we move from left to right across a period, atomic and ionic radii generally decrease. This shrinking size has two direct consequences:
- Ionization enthalpy generally increases (with some exceptions, as discussed in Section 3.7.1(c)).
- Electron gain enthalpy becomes more negative — meaning the element releases more energy when it gains an electron.
Noble gases are an exception: they have positive electron gain enthalpy values because their completely filled shells make them reluctant to accept an extra electron.
The result is striking: chemical reactivity is highest at the two extremes of a period and lowest in the centre.
Why? At the extreme left (group 1, alkali metals), the ionization enthalpy is the lowest in the period. These elements can lose an electron very easily, forming a cation. At the extreme right (group 17, halogens), the electron gain enthalpy is the most negative (highest magnitude). These elements readily gain an electron, forming an anion. In the centre of the period, elements have intermediate ionization enthalpies and electron gain enthalpies — they neither lose nor gain electrons easily, so their reactivity is lower.
This behaviour directly relates to two fundamental chemical properties:
- Reducing behaviour — the tendency to lose electrons (strongest at the left extreme).
- Oxidizing behaviour — the tendency to gain electrons (strongest at the right extreme).
You will study these in detail later, but for now, the key link is to metallic and non-metallic character.
Metallic character is highest at the extreme left of a period and decreases as we move right. Non-metallic character increases from left to right. The most metallic element in a period is the alkali metal; the most non-metallic is the halogen.
Chemical Reactivity Shown Through Reactions with Oxygen
The textbook uses reactions with oxygen as a clear demonstration of periodic trends in chemical reactivity. Elements at both extremes of a period combine readily with oxygen to form oxides. But the nature of these oxides is completely different.
Oxides of Elements at the Left Extreme
The normal oxide formed by the element on the extreme left is basic. For example, sodium (group 1, period 3) forms sodium oxide, .
Oxides of Elements at the Right Extreme
The normal oxide formed by the element on the extreme right is acidic. For example, chlorine (group 17, period 3) forms dichlorine heptoxide, .
Oxides of Elements in the Centre
Elements in the centre of a period form oxides that are either amphoteric or neutral.
- Amphoteric oxides behave as acids with bases and as bases with acids. Examples: (aluminium oxide), (arsenic trioxide).
- Neutral oxides have neither acidic nor basic properties. Examples: (carbon monoxide), (nitric oxide), (dinitrogen oxide).
The term "normal oxide" here means the oxide in which the element exhibits its most common oxidation state. For sodium, that is (oxidation state +1). For chlorine, it is (oxidation state +7).
Problem 3.10 in this section's Problems group below shows these acid–base characters through the reactions of and with water.
This pattern — basic oxides on the left, acidic oxides on the right, amphoteric/neutral in the centre — is a powerful tool for predicting the chemical behaviour of an element based solely on its position in the periodic table.
Trends Among Transition and Inner-Transition Metals
The discussion so far has focused on representative elements (s- and p-block). Transition metals and inner-transition metals show different behaviour.
Transition Metals (3d Series)
Across a period in the transition series, the change in atomic radii is much smaller than in representative elements. This is because the added electrons go into inner d-orbitals, which do not expand the atom as much. The ionization enthalpies of transition metals are intermediate between those of s-block and p-block elements.
Because their ionization enthalpies are higher than those of group 1 and 2 metals, transition metals are less electropositive than alkali and alkaline earth metals. They do not lose electrons as readily.
Inner-Transition Metals (4f Series)
The change in atomic radii is even smaller among the lanthanoids (4f series). This is the phenomenon of lanthanoid contraction — the gradual decrease in atomic size across the series is very slight because the 4f electrons are poorly shielding.
Reactivity Trends Down a Group
Moving down a group, the pattern reverses compared to moving across a period.
For Main Group Elements
As atomic number increases down a group:
- Atomic and ionic radii increase.
- Ionization enthalpy gradually decreases. …