Periodic trends and effective nuclear charge

• henry3369
In summary, the book discusses the relationship between atomic radius and effective nuclear charge on the periodic table. As you move towards the right side, the effective nuclear charge increases, resulting in a smaller radius. However, an increase in valence electrons can also decrease the effective nuclear charge, as the shielding effect increases. This means that when the proton number Z increases by 1, the new electron's shielding effect is less than one elementary charge.
henry3369
My book is trying to explain why atomic radii decreases as you move toward the right side of the periodic table because the effective nuclear charge increases. I understand why an increase in effective nuclear charge results in a smaller radius, but I don't know why the effective nuclear charge would decrease with more valence electrons. If Zeff= Z-S, then an increase in valence electrons would increase shielding; thus, Zeff would be smaller with an increase in S.

You increase the proton number Z by 1, but the new electron won't be completely inside so its shielding effect is less than one elementary charge.

1. What are periodic trends in the periodic table?

Periodic trends refer to the patterns or trends that can be observed in the properties of elements as you move across a period or down a group in the periodic table. These trends include atomic radius, ionization energy, electron affinity, and electronegativity, among others.

2. What is effective nuclear charge?

Effective nuclear charge is the net positive charge experienced by an electron in an atom. It takes into account the attraction between the positively charged nucleus and the negatively charged electron, as well as the shielding effect of inner shell electrons. The higher the effective nuclear charge, the stronger the attraction between the nucleus and outer electrons, resulting in smaller atomic radius and higher ionization energy.

3. How does atomic radius change across a period and down a group?

Across a period, atomic radius decreases as effective nuclear charge increases. This is due to the increasing number of protons in the nucleus, which pulls the electrons closer to the center, resulting in a smaller atomic radius. Down a group, atomic radius increases due to the addition of new energy levels, which increases the distance between the nucleus and outermost electrons.

4. What is the trend for ionization energy in the periodic table?

Ionization energy generally increases across a period and decreases down a group. This is because as you move across a period, the effective nuclear charge increases, making it more difficult to remove an electron from an atom. On the other hand, down a group, the outermost electrons are farther from the nucleus and experience less attraction, making it easier to remove them.

5. How does electron affinity change across a period and down a group?

Across a period, electron affinity generally increases as effective nuclear charge increases. This is because as the effective nuclear charge increases, the attraction between the nucleus and electrons also increases, making it easier for the atom to gain an electron. Down a group, electron affinity decreases due to the increasing atomic radius, which results in less attraction between the nucleus and electrons.

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