3.7.3. Periodic Trends Chemical and Reactivity

We have observed the periodic trends in certain fundamental properties such as atomic and ionic radii, ionization enthalpy, electron gain enthalpy and valence. We know by now that the periodicity is related to electronic configuration.

Periodic trends refer to the recurring patterns observed in the properties of elements as we move through the periodic table. For example, atomic and ionic radii refer to the size of atoms and ions, while ionization enthalpy is the energy required to remove an electron from an atom. These properties are influenced by the arrangement of electrons in an atom. As we study these properties, we can make connections between the structure of the atoms and their chemical behavior, leading to a deeper understanding of why elements behave the way they do in reactions.

Consider the periodic table like a family where each group consists of relatives sharing common traits. Just like some family members might have similar interests or characteristics based on their upbringing, elements in the same group have similar electron configurations, resulting in similar chemical properties. This helps chemists predict how different elements will react with each other.

The atomic and ionic radii, as we know, generally decrease in a period from left to right. As a consequence, the ionization enthalpies generally increase (with some exceptions as outlined in section 3.7.1(a)) and electron gain enthalpies become more negative across a period.

As we move across a period in the periodic table, the size of atoms typically decreases. This is because the number of protons in the nucleus increases, pulling the electrons closer to the nucleus and resulting in a smaller atomic size. Consequently, it becomes harder to remove an electron (higher ionization enthalpy) and easier to gain an electron (lower or more negative electron gain enthalpy) as we move from left to right. This trend directly relates to how reactive an element can be; elements on the far left are highly reactive metals, while those on the far right are highly reactive nonmetals.

Imagine a game of tug-of-war. When the people on one side pull the rope (electrons) closer to their side (the nucleus), it becomes much harder for someone from the other side to snatch the rope. Similarly, in an atom, as more protons pull on the electrons, those electrons become harder to remove, increasing ionization enthalpy and decreasing their likelihood to react by gaining more electrons.

This results in high chemical reactivity at the two extremes and the lowest in the centre. Thus, the maximum chemical reactivity at the extreme left (among alkali metals) is exhibited by the loss of an electron leading to the formation of a cation and at the extreme right (among halogens) shown by the gain of an electron forming an anion.

The reactivity of elements varies significantly across the periodic table. On the left side, the alkali metals (Group 1) readily lose an electron to form positive ions (cations), making them highly reactive. Conversely, the halogens (Group 17) are very reactive nonmetals that tend to gain electrons to form negative ions (anions). The elements in the center typically have moderate reactivity. This behavior can be attributed to their ionization energies and their ability to gain or lose electrons based on their atomic structure.

Think of a crowded party where people have different personalities. Some individuals (like sodium) are eager to share their snacks (electrons) with others and leave the party, while others (like chlorine) are keen on grabbing snacks from everyone else and joining together to make a great dish (an anion). The more people crowding the space (electrons), the more chaotic things become, reflecting the different levels of reactivity among elements.

The oxides formed by the elements on extreme left are basic and those formed by the elements on extreme right are acidic in nature, whereas the oxides of elements in the center are amphoteric (e.g., Al2O3, As2O3) or neutral (e.g., CO, NO, N2O).

The properties of oxides formed by elements in the periodic table vary based on their position. Basic oxides, often formed by metals, react with acids and can neutralize them. Acidic oxides, commonly formed by nonmetals, react with bases. Amphoteric oxides have the ability to react with both acids and bases. This classification helps in predicting the nature of oxides formed by different elements based on their group and period in the periodic table.

Imagine a school where students can take on different roles. Some students (basic oxides) are great at helping others learn when they struggle with difficult subjects (neutralizing acids), whereas others (acidic oxides) can teach others but are firm and don’t take assistance (reacting with bases). Some students can adapt their teaching methods (amphoteric oxides) depending on whether they are working with a classmate or needing help themselves.

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