Question

Which of the following statements best explains why the atomic radius decreases from left to right across a period? A. The number of neutrons in an atom increases from left to right across a period, and they pull the valence electrons inward with greater force. B.The number of electrons decreases from left to right, so the electron cloud occupies a smaller volume. C The charge of the nucleus of an atom increases from left to right across a period, so the valence electrons are pulled inward with greater force. D. The number of energy levels decreases from left to right, so the electron cloud occupies a smaller volume.

Answer 1

The atomic radius decreases across a period primarily due to the increase in the positive charge of the nucleus attracting the valence electrons more strongly, which is option C. As the nuclear charge increases, the effective nuclear charge also increases, causing the electrons to be pulled closer to the nucleus, leading to a smaller atomic radius.

Step-by-step explanation:

The correct explanation for why the atomic radius decreases from left to right across a period is: C. The charge of the nucleus of an atom increases from left to right across a period, so the valence electrons are pulled inward with greater force.

As the number of protons in the nucleus increases within a period, the positive charge of the nucleus also increases. Consequently, the effective nuclear charge (Zeff), which is the net charge felt by valence electrons after accounting for shielding by inner electrons, increases. This stronger attraction pulls the valence electrons closer to the nucleus, which results in a smaller atomic radius.

The notions of electrons being added to the same energy level and an increasing nuclear charge are crucial. The inner electron shells shield the valence electrons to an extent, but as Zeff escalates, it overcomes this shielding, leading to a reduction in atomic size across a period.

Answer 2

Option C correctly explains this phenomenon by emphasizing the increase in the charge of the nucleus from left to right, which leads to the stronger attraction and the reduction in atomic radius.

The correct statement explaining why the atomic radius decreases from left to right across a period is:

C. The charge of the nucleus of an atom increases from left to right across a period, so the valence electrons are pulled inward with greater force.

Here's a step-by-step explanation:

1. Atomic radius is defined as the distance between the nucleus of an atom and its outermost electron shell (valence shell). As you move across a period from left to right on the periodic table, you are examining elements with increasing atomic number.

2. The atomic number represents the number of protons in the nucleus of an atom. Since each proton carries a positive charge, an increase in atomic number means an increase in the positive charge in the nucleus.

3. According to Coulomb's law, which governs the electrostatic attraction between charged particles, an increase in the positive charge (protons) in the nucleus will result in a stronger attraction between the positively charged nucleus and the negatively charged electrons in the electron cloud.

4. This increased attractive force causes the valence electrons (those in the outermost electron shell) to be pulled closer to the nucleus. As a result, the atomic radius decreases because the electrons are held more tightly to the nucleus.