Question

Choose the element with the larger atoms from each pair. (a) Sn or Si (b) Br or Ga (c) \(S n\) or \(B i\) (d) Se or Sn

Short answer

Larger atoms for each pair are (a) Sn (b) Br (c) Bi (d) Se.

Step-by-step solution

Understanding Atomic Size

The size of an atom generally increases as one moves down a group in the periodic table because more electron shells are added. The atomic size decreases as one moves from left to right across a period because the number of protons increases, pulling the electron clouds closer.

Comparing Atoms in the Same Group

For elements in the same group of the periodic table, choose the element that is located further down, since it will have more electron shells and therefore larger atoms.

Comparing Sn and Si

Sn (Tin) and Si (Silicon) are in the same group of the periodic table (group 14). Since Sn is lower down the group than Si, Sn has larger atoms.

Comparing Br and Ga

Br (Bromine) and Ga (Gallium) are in different groups. Br is in group 17 and Ga is in group 13. Despite being in different periods, since Br is farther to the right and down the periodic table, previous trends suggest Br has a larger atomic size than Ga.

Comparing Sn and Bi

Sn (Tin) and Bi (Bismuth) are in the same group of the periodic table (group 15). Bi is located further down the group than Sn, meaning Bi has a larger atomic size.

Comparing Se and Sn

Se (Selenium) and Sn (Tin) are in the same period but different groups, with Se in group 16 and Sn in group 14. Sn is farther to the right of Se, indicating that Sn has more protons and a smaller atomic radius.

Key concepts

Periodic Table Trends

When comparing atomic sizes, it's crucial to understand the trends present in the periodic table. The periodic table is systematically arranged to showcase patterns in the properties of elements, one of which is the variation in atomic sizes.

As we move from top to bottom within a group (vertical columns), we observe an increase in the number of electron shells. This addition of shells results in larger atoms further down a group. On the other hand, as we go from left to right across a period (horizontal rows), the atomic number, which represents the number of protons in the nucleus, increases. This increase leads to a stronger attraction between the positively charged nucleus and the negatively charged electrons, which pulls the electron clouds closer and results in a decrease in atomic size.

• Moving down a group increases atomic size due to additional electron shells.
• Moving across a period decreases atomic size due to greater nuclear attraction.

By understanding these trends, determining which of two elements has the larger atom becomes a systematic process based on their positions on the periodic table.

Electron Shells

Electron shells are the layers of electrons that surround the nucleus of an atom. Each shell can hold a certain maximum number of electrons and is filled in order of increasing energy.

The innermost shell is the first to be filled and can hold up to two electrons. Subsequent shells can hold more, with the general formula for the maximum number of electrons being given by the equation \( 2n^2 \), where \( n \) is the principal quantum number or the shell number. The presence of more electron shells in an atom means it is larger because the outermost electrons are further from the nucleus.

Principal Quantum Number (Shell Number)

The principal quantum number, starting at 1 and increasing for outer shells, determines the size and energy level of the shell. As 'n' increases, the shell can accommodate more electrons, and its average distance from the nucleus also increases. This results in larger atomic sizes for elements with more electron shells.

• More electron shells mean a larger electron cloud and therefore a bigger atom.
• The number of electrons a shell can hold increases with its distance from the nucleus.

Atomic Radius

The atomic radius is a measure of the size of an atom, specifically the distance from the center of the nucleus to the outer boundary of the electron cloud. Since the electron cloud is not a hard shell, the atomic radius is often described in terms of where there is a high probability of finding the outermost electrons.

Factors that affect the atomic radius include the number of electron shells, as previously mentioned, and the effective nuclear charge—the net positive charge experienced by the outermost electrons. As the effective nuclear charge increases due to additional protons in the nucleus, the electrons are pulled closer to the center, leading to a smaller atomic radius.

Comparing Atomic Radii

To accurately compare atomic radii, you must consider both the number of electron shells and the effective nuclear charge.

• An atom with more electron shells will generally have a larger radius.
• An atom with a higher effective nuclear charge (more protons) will generally have a smaller radius, due to electrons being pulled in tighter.

For elements within the same period, the increasing number of protons leads to a decrease in atomic radius from left to right. For elements within the same group, the addition of electron shells leads to an increase in atomic radius from top to bottom.