MATRICULATION CHEMISTRY SK015: 3.2 PERIODICITY (ATOMIC RADII, IONIC RADII & ISOELECTRONIC SPECIES)

Summary

This educational video delves into the concept of periodicity within the periodic table, focusing on atomic, ionic, and iso-electronic species radii. It explains how atomic radius decreases across a period due to increasing nuclear charge and strengthens electromagnetic attraction. Conversely, it increases down a group due to more electron shells and shielding effects. The video also discusses the trends in the first row of transition elements and the impact of electron configurations on ionic radii, illustrating the differences between cations and anions. Lastly, it touches on isoelectronic species and how their radii are influenced by nuclear charge and electron configurations across periods.

Takeaways

• 🌐 The script discusses the concept of periodicity in the periodic table, focusing on the trends of atomic and ionic radii.
• 🔬 Atomic radius decreases across a period due to increasing proton number and effective nuclear charge, leading to stronger electromagnetic attraction between the nucleus and valence electrons.
• 📉 The smallest atomic radius is found at the far right of a period due to the strongest nuclear-valence electron attraction.
• 📈 Atomic radius increases down a group due to the addition of electron shells and an increase in the shielding effect, which weakens the attraction between the nucleus and valence electrons.
• 🏙️ The largest atomic radius is found at the bottom of a period due to the weakest nuclear-valence electron attraction.
• 🔄 For the first row of transition elements in period 4, atomic radius decreases and then remains fairly constant, influenced by the 3d and 4s orbitals' electron configurations.
• 🔮 Ionic radius is affected by the gain or loss of electrons; anions expand due to electron addition and increased electron-electron repulsion, while cations contract due to electron loss and reduced repulsion.
• 📊 Ionic size decreases across a period for both cations and anions due to increasing nuclear charge and effective nuclear charge.
• 🧬 Isoelectronic species, which have the same electronic configurations, can be compared across periods to understand trends in ionic radii.
• 🔑 The trend in isoelectronic species can be segmented and explained by the number of protons, effective nuclear charge, and the resulting nuclear-electron attractions.
• 🚀 The script also mentions upcoming discussions on other periodic trends such as ionization energy and electronegativity.

Q & A

What is the primary factor affecting atomic radius trends across a period?

The primary factor affecting atomic radius trends across a period is the increasing proton number in the nucleus, which leads to an increase in effective nuclear charge, resulting in stronger electromagnetic attractions between the nucleus and valence electrons and causing the radius to shrink.

Why does the atomic radius decrease as we move from left to right across a period?

The atomic radius decreases across a period because the increasing number of protons in the nucleus enhances the effective nuclear charge, leading to stronger attractions between the nucleus and valence electrons, which in turn pulls the electrons closer to the nucleus and reduces the atomic radius.

How does the atomic radius trend when moving down a group in the periodic table?

When moving down a group, the atomic radius is expected to increase due to the addition of more electron shells. This results in an increased number of electrons and a stronger shielding effect, which weakens the attraction between the nucleus and valence electrons, leading to a larger atomic radius.

What is the shielding effect and how does it influence atomic radius?

The shielding effect is the phenomenon where inner electrons shield the nucleus from the valence electrons, reducing the effective attraction between them. As a result, the nucleus's attraction towards the valence electrons decreases, leading to an increase in the distance between the nucleus and electrons, and thus a larger atomic radius.

Why does the atomic radius of elements in the first row of transition metals show a noticeable decrease across the first five elements and then remain fairly constant?

The atomic radius of the first row of transition metals decreases due to the increasing nuclear charge across the period. However, after the first five elements, the presence of 3d inner orbitals filled with electrons shields the 4s orbital more effectively from the nucleus, even as the nuclear charge increases, making the size remain fairly unchanged.

How does the ionic radius differ from the atomic radius and why?

Ionic radius differs from atomic radius because electrons repel each other. When an atom gains electrons to form an anion, the electron cloud expands, increasing the ionic radius. Conversely, when an atom loses electrons to form a cation, the electron cloud shrinks, decreasing the ionic radius.

What happens to the atomic radius when an atom forms a cation by losing an electron?

When an atom forms a cation by losing an electron, the electron cloud shrinks due to a decrease in electron-electron repulsions. This results in stronger attractions between the nucleus and the remaining electrons, causing the atomic radius to be smaller than that of the neutral atom.

How does the ionic radius trend when comparing isoelectronic species across a period?

For isoelectronic species across a period, the ionic radius decreases as the atomic number increases due to the stronger effective nuclear charge. This results in stronger attractions between the nucleus and the electrons, leading to a smaller ionic radius.

What causes the significant jump in ionic radius between cations and anions formed in the same period?

The significant jump in ionic radius between cations and anions is caused by the difference in the number of electron shells. Cations have fewer shells and thus stronger nucleus-electron attractions, while anions have more shells, leading to increased electron-electron repulsions and a larger ionic radius.

How can the trend of isoelectronic species be segmented and explained?

The trend of isoelectronic species can be segmented based on the groups of elements with the same electronic configurations. By identifying the species involved and their electronic configurations, one can explain how the trend across a period affects the size of the ions, considering the increasing nuclear charge and the resulting changes in nucleus-electron attractions.