Nuclear Charge, Shielding Effect & Effective Nuclear Charge (Singapore A Level H2 Chemistry)

Summary

This video explores key concepts in atomic structure, focusing on nuclear charge, shielding effect, and effective nuclear charge. It explains how the number of protons in an atom affects the nuclear charge, which in turn impacts the atomic radius. The shielding effect, caused by inner electrons repelling outer electrons, works against nuclear charge, leading to an overall effective nuclear charge that determines the strength of attraction between the nucleus and valence electrons. Understanding these principles helps explain trends in atomic properties like ionization energies and atomic size across periods and groups in the periodic table.

Takeaways

• 😀 Nuclear charge refers to the positive charge in the nucleus, which is determined by the number of protons present in an atom.
• 😀 The nuclear charge creates an attractive force between the nucleus and the valence electrons, leading to a contractionary effect on the atomic size.
• 😀 As the number of protons in an atom increases, the nuclear charge also increases, causing the atomic radius to become smaller due to stronger attraction between nucleus and valence electrons.
• 😀 The shielding effect occurs when electrons in inner shells repel the outer shell electrons, reducing the effective nuclear charge felt by the valence electrons.
• 😀 The shielding effect has an expansionary impact, causing the atomic radius to increase as the valence electrons are less attracted to the nucleus.
• 😀 Nuclear charge and shielding effect act together to determine the overall size and behavior of the atom, which results in the concept of effective nuclear charge (Z_eff).
• 😀 Effective nuclear charge (Z_eff) is calculated by subtracting the shielding effect (S) from the nuclear charge (Z): Z_eff = Z - S.
• 😀 When nuclear charge outweighs the shielding effect, the effective nuclear charge increases, leading to a stronger attraction between the nucleus and the valence electrons.
• 😀 A higher effective nuclear charge results in a smaller atomic radius, as the stronger pull from the nucleus contracts the atom.
• 😀 Both nuclear charge and shielding effect play a key role in explaining periodic trends, including atomic radius, ionic radius, ionization energies, and successive ionization energies.

Q & A

What is nuclear charge, and how does it affect atomic radius?

Nuclear charge refers to the positive charge in the nucleus due to protons. As the number of protons increases, the nuclear charge becomes stronger, attracting the valence electrons more tightly, which causes the atomic radius to shrink.

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

The shielding effect occurs when inner-shell electrons repel outer-shell electrons, reducing the full effect of the nuclear charge on the valence electrons. This effect causes the atomic radius to expand as the outer electrons experience less attraction from the nucleus.

How is effective nuclear charge (Z_eff) defined?

Effective nuclear charge (Z_eff) is the net charge felt by the valence electrons after accounting for the shielding effect. It is calculated by subtracting the shielding effect from the nuclear charge.

What is the formula for calculating effective nuclear charge?

The formula for effective nuclear charge is: Effective Nuclear Charge (Z_eff) = Nuclear Charge - Shielding Effect.

How does the nuclear charge affect ionization energies?

A higher nuclear charge increases the electrostatic attraction between the nucleus and the valence electrons, making it harder to remove electrons. This results in higher ionization energies across a period.

Why does the atomic radius generally decrease across a period?

The atomic radius decreases across a period because as the nuclear charge increases (more protons), the effective nuclear charge also increases, pulling the electrons closer to the nucleus and reducing the atomic radius.

What trend occurs in atomic radius as you move down a group in the periodic table?

As you move down a group, the atomic radius increases. This happens because new electron shells are added, increasing the distance between the nucleus and the valence electrons, while the shielding effect becomes stronger.

How does shielding affect the effective nuclear charge?

Shielding reduces the effect of the nuclear charge on outer electrons. The more inner-shell electrons there are, the greater the shielding effect, which lowers the effective nuclear charge experienced by the valence electrons.

What happens to the effective nuclear charge when nuclear charge exceeds shielding?

When the nuclear charge is greater than the shielding effect, the effective nuclear charge increases, leading to a stronger attraction between the nucleus and the valence electrons.

Why do successive ionization energies increase?

Successive ionization energies increase because as electrons are removed, the remaining electrons experience a stronger effective nuclear charge, making them harder to remove. This increase in attraction between the nucleus and remaining electrons results in higher ionization energies.