Question

Which of the following statements is false? (1) Electronegativity values of \(\mathrm{Si}\), Ge, Sn and \(\mathrm{Pb}\) are equal. (2) Electronegativity of carbon is greater than silicon. (3) Density of silicon is more than that of carbon (diamond). (4) The atomic radius of \(\mathrm{Si}\) is greater than that of carbon.

Short answer

Statement (1) is false.

Step-by-step solution

Analyze Statement (1)

Check the electronegativity values of Si, Ge, Sn, and Pb. Electronegativity generally decreases from Si to Pb in the periodic table. Therefore, the electronegativity values of Si, Ge, Sn, and Pb are not equal.

Analyze Statement (2)

Compare the electronegativity of carbon and silicon. Carbon has an electronegativity of 2.55, while silicon has an electronegativity of 1.90. Hence, the electronegativity of carbon is indeed greater than that of silicon.

Analyze Statement (3)

Investigate the densities of silicon and carbon (diamond). Diamond has a density of about 3.51 g/cm³, while silicon's density is approximately 2.33 g/cm³. Thus, the density of silicon is less than that of carbon (diamond).

Analyze Statement (4)

Check the atomic radii of silicon and carbon. The atomic radius of silicon (around 110 pm) is greater than that of carbon (around 70 pm).

Identify the False Statement

Based on the analysis: - Statement (1) is false because the electronegativity values of Si, Ge, Sn, and Pb are not equal. - Statements (2), (3), and (4) are true.

Key concepts

Electronegativity Trend

Electronegativity is a measure of an atom's ability to attract and hold onto electrons. In the periodic table, electronegativity generally increases as you move from left to right across a period. This means elements on the right side of the periodic table, like fluorine, tend to have high electronegativity values.
In contrast, as you move down a group (column), electronegativity tends to decrease. This decrease happens because the additional electron shells added as you go down a group increase the distance between the nucleus and the outer electrons, making it harder for the nucleus to attract additional electrons.
Therefore, the electronegativity values of Si, Ge, Sn, and Pb (all in group 14) are not the same and decrease from Si to Pb.

Periodic Table Trends

The periodic table displays several key trends that help in understanding the properties and behaviors of elements.

• Atomic Radius: The atomic radius generally decreases across a period from left to right due to the increase in positive charge in the nucleus, which pulls the electrons closer.
• Density: Density tends to increase as you move down a group. This is because additional protons and neutrons are added to the nucleus, making the atoms heavier while the volume doesn't increase as much.
• Electronegativity: As explained before, electronegativity increases across a period and decreases down a group.

By understanding these trends, it becomes easier to predict the properties of elements and their reactions.

Atomic Radius

Atomic radius refers to the size of an atom. It is the distance from the center of the nucleus to the outermost shell of electrons. Several factors affect the atomic radius:

• Nuclear Charge: Higher nuclear charge pulls electrons closer, reducing the atomic radius.
• Electron Shielding: Inner shell electrons can shield outer electrons from the nucleus, increasing the atomic radius.

Moving left to right across a period, the atomic radius decreases due to increasing nuclear charge without additional electron shielding.
As you move down a group, the atomic radius increases because of the additional electron shells added, which outweighs the increase in nuclear charge. This explains why silicon (Si) has a larger atomic radius (around 110 pm) compared to carbon (C), which has an atomic radius of around 70 pm.

Density Comparison

Density is defined as mass per unit volume and can vary greatly among different elements. Factors influencing density include atomic mass and atomic radius.
When comparing densities, for instance, silicon and carbon (diamond), we find:

• Carbon (diamond): Diamond is very densely packed due to its crystal structure, giving it a density of about 3.51 g/cm³.
• Silicon: Silicon has a less dense packing structure, resulting in a lower density of about 2.33 g/cm³.

This means that diamond, an allotrope of carbon, is denser than silicon, which is less densely packed despite being in the same group on the periodic table. By understanding these structural and positional influences on elements, one can make informed predictions about their physical properties.