Question

Which of the following statements about effective nuclear charge for the outermost valence electron of an atom is incorrect? (i) The effective nuclear charge can be thought of as the true nuclear charge minus a screening constant due to the other electrons in the atom. (ii) Effective nuclear charge increases going left to right across a row of the periodic table. (iii) Valence electrons screen the nuclear charge more effectively than do core electrons. (iv) The effective nuclear charge shows a sudden decrease when we go from the end of one row to the beginning of the next row of the periodic table. (v) The change in effective nuclear charge going down a column of the periodic table is generally less than that going across a row of the periodic table.

Short answer

Statement (iii) is incorrect, as core electrons, not valence electrons, screen the nuclear charge more effectively due to their closer proximity to the nucleus.

Step-by-step solution

Statement (i)

The effective nuclear charge can be thought of as the true nuclear charge minus a screening constant due to the other electrons in the atom. This statement is correct, as the effective nuclear charge is the net positive charge experienced by the valence electrons from the nucleus, after accounting for the shielding/screening effect of the inner core electrons.

Statement (ii)

Effective nuclear charge increases going left to right across a row of the periodic table. This statement is also correct. As we move left to right across a period, the number of protons in the nucleus increases, which in turn increases the nuclear charge. The increase in nuclear charge outweighs the shielding effect of increasing electrons, leading to a higher effective nuclear charge on the valence electrons.

Statement (iii)

Valence electrons screen the nuclear charge more effectively than do core electrons. This statement is incorrect. Core electrons (those closer to the nucleus) are generally more effective at shielding the nuclear charge, as they are located between the valence electrons and the nucleus. Valence electrons do not shield the nuclear charge as effectively as the core electrons because they are further from the nucleus and don't fully block the attraction between the nucleus and the valence electrons.

Statement (iv)

The effective nuclear charge shows a sudden decrease when we go from the end of one row to the beginning of the next row of the periodic table. This statement is correct. Going from the end of one period to the beginning of the next one, there is an addition of a new energy level (shell). This causes a noticeable shielding effect, indicating a sudden decrease in effective nuclear charge experienced by the valence electrons even though the number of protons has increased.

Statement (v)

The change in effective nuclear charge going down a column of the periodic table is generally less than that going across a row of the periodic table. This statement is correct. As we move down a column (group), the number of energy levels (shells) increases, which adds more shielding electrons. This increase in shielding outweighs the increase in nuclear charge, so there is a smaller change in effective nuclear charge when going down a column than when going across a row in the periodic table.

Key concepts

Shielding Effect

The shielding effect is a fundamental concept in atomic physics, referring to the reduction in the effective nuclear charge on the electron cloud due to the repulsion from electrons in lower energy levels. Imagine the nucleus of an atom as a powerful magnet and each electron as a mini shield, with core electrons acting as stronger shields than valence electrons.

Core electrons, being closer to the nucleus, shield valence electrons from the full effect of the nuclear charge. As you might picture, the inner-layered electrons intercept some of the 'magnetic pull' from the nucleus before it reaches the outermost electrons. It's like trying to hear someone across a crowded room; the more people (core electrons) between you, the less clearly you hear (less nuclear pull felt).

When considering the shielding effect, it's essential to remember that valence electrons, which are in the highest energy level, do not significantly shield each other from the nuclear charge.

Periodic Trends

Periodic trends are patterns observed within elements on the periodic table, which help predict an element's properties. Among these is atomic radius, ionization energy, electronegativity, and the topic at hand, effective nuclear charge.

As we move across a period (left to right) on the periodic table, the effective nuclear charge experienced by valence electrons increases. This is because while additional electrons are being added to the same energy level, the nucleus gains more protons, enhancing its pull on the electrons. Even as electrons 'shield' each other, the growing number of protons has a more robust ability to draw the outer electrons inwards, decreasing the atomic radius and increasing attraction in the atom.

Going Down a Group

Conversely, when descending down a group (top to bottom), each element acquires a new electron shell. This additional shell increases the distance from the nucleus and adds more shielding electrons, which almost dampens the effect of adding more protons. The result? There is less change in effective nuclear charge down a column compared to across a row. This helps explain why atoms generally get larger as you go down a group on the table.

Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom. They are extremely important because they dictate how an atom will interact with others to form chemical bonds. If atomic behavior at parties were to be considered, valence electrons would be the social butterflies—they make the connections.

These electrons are also subjected to the effective nuclear charge, which influences their energy and reactivity. The fewer valence electrons compared to the available spaces in an atom's outer shell, the more reactive the atom tends to be.

Role in Chemical Bonding

For example, atoms strive for a full valence shell, often leading to the transfer or sharing of valence electrons to achieve stability, thus forming ionic or covalent bonds respectively. Understanding both the shielding effect and how it influences the effective nuclear charge helps provide insight into the atomic radius, ionization energy, and overall reactivity of an element based on its position on the periodic table.