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 (iii) Valence going left to right across a row of the periodic table. 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

The incorrect statement is (iii): Valence electrons screen the nuclear charge more effectively than do core electrons. Core electrons actually screen the nuclear charge more effectively due to their closer proximity to the nucleus.

Step-by-step solution

Understanding the concept of effective nuclear charge

Effective nuclear charge is the net positive charge experienced by the electrons in the outermost shell of an atom, considering the shielding or screening effects of the inner shell electrons. It represents the attractive force experienced by the valence electrons in an atom.

Analyzing 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 it accurately describes the concept of effective nuclear charge. The true nuclear charge is the total positive charge of the protons in the nucleus, while the screening constant accounts for the repulsion experienced by the valence electrons due to the core (inner) electrons.

Analyzing 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 across a period (left to right), the atomic number increases, and the number of protons in the nucleus also increases. Since there is no significant increase in the shielding effect (the electron shielding remains relatively constant), the effective nuclear charge felt by the valence electrons will increase.

Analyzing statement (iii)

Valence electrons screen the nuclear charge more effectively than do core electrons. This statement is incorrect. Core (inner) electrons actually screen the nuclear charge more effectively than the valence electrons, as they are closer to the nucleus and create a more significant shielding effect. The valence electrons, which are further from the nucleus, experience less effective screening.

Analyzing 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. When we move from the end of one row (a noble gas) to the beginning of the next row (an alkali metal), there is a significant increase in the shielding effect due to the addition of a new electron shell. This results in a sudden decrease in the effective nuclear charge experienced by the valence electrons.

Analyzing 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 also correct. As we move down a column (or group) in the periodic table, the shielding effect increases due to the addition of new electron shells, but the number of protons in the nucleus also increases. These two effects tend to balance each other out, resulting in a smaller change in effective nuclear charge compared to going across a row. The incorrect statement is (iii): Valence electrons screen the nuclear charge more effectively than do core electrons. Core electrons actually screen the nuclear charge more effectively due to their closer proximity to the nucleus.

Key concepts

Shielding Effect

The shielding effect, also known as the screening effect, plays a crucial role in understanding how valence electrons experience the nuclear charge of an atom. Inner electrons, particularly the core electrons, provide a repelling effect on the valence electrons. This effect reduces the full positive charge of the nucleus that valence electrons would otherwise feel. Since these core electrons are positioned closer to the nucleus, they effectively 'shield' outer electrons from the full force of positive charge.

To better understand this concept:

• The true nuclear charge ($Z$) represents the total charge from all protons.
• Core electrons dull this charge by providing a shielding constant ($S$).
• Thus, the effective nuclear charge ($Z_{exteff}$) felt by the valence electrons is calculated as $Z-S$.

In essence, while core electrons significantly shield the nuclear charge, valence electrons do not contribute much to this effect due to their distance from the nucleus. This is why the effective nuclear charge is lesser than the actual nuclear charge, especially in heavy elements.

Valence Electrons

Valence electrons are the electrons located in the outermost shell of an atom. They are crucial for determining an atom's chemical properties, including its bonding behavior. These outermost electrons are the ones involved in chemical reactions and forming bonds with other atoms.

One important thing about valence electrons is how they interact with the nucleus of an atom:

• Being the farthest away from the nucleus, they experience less nuclear attraction than inner electrons.
• The extent of attraction they feel is greatly influenced by the effective nuclear charge.

Despite being crucial in many chemical processes, valence electrons alone are not strong screeners of nuclear charge. Instead, the core electrons are the key players in reducing the nuclear charge that valence electrons sense, leading to unique trends across the periodic table.

Periodic Table Trends

The periodic table is not just a collection of elements, but a tool that helps predict and understand an atom's properties and behavior. When studying elements across periods and groups, various trends linked to the effective nuclear charge reveal themselves.

As you move across a period from left to right:

• The atomic number ($Z$) increases, adding more protons to the nucleus.
• Although electrons are added as well, they don't significantly increase shielding, leading to a higher effective nuclear charge.
• This causes valence electrons to be pulled closer and affects properties like atomic size and ionization energy.

Conversely, when moving down a group:

• The number of electron shells increases, enhancing the shielding effect considerably.
• This tends to offset the increase in nuclear charge due to additional protons, making the change in effective nuclear charge less pronounced compared to across a period.
• As a result, atoms generally become larger and ionization energies decrease down a group.

These insights highlight how the effective nuclear charge influences elemental characteristics across the periodic table. Understanding these trends helps predict how different elements might react chemically, aiding in grasping the broader principles of chemistry.