Question

Arrange these elements in order of increasing atomic size: \(\mathrm{Cs}, \mathrm{Sb}, \mathrm{S}, \mathrm{Pb}\), Se.

Short answer

In order of increasing atomic size: S < Se < Sb < Pb < Cs.

Step-by-step solution

Understand Atomic Size Trends on the Periodic Table

Atomic size generally increases as you move down along a group in the periodic table because each row adds a new electron shell. Atomic size decreases across a period from left to right due to the increase in the nuclear charge which attracts the electrons more strongly and thus decreases the atomic radius.

Locate the Elements on the Periodic Table

Find the position of each element in the periodic table to understand the trend in their atomic sizes. Cesium (Cs) is in group 1, period 6. Antimony (Sb) is in group 15, period 5. Sulfur (S) is in group 16, period 3. Lead (Pb) is in group 14, period 6. Selenium (Se) is in group 16, period 4.

Compare Atomic Size within the Same Group

For elements in the same group, compare their periods. The higher the period number, the larger the atomic size. Thus, within the same group, Cs (period 6) is larger than Sb (period 5), and Se (period 4) is larger than S (period 3).

Compare Atomic Size within the Same Period

For elements in the same period, compare their group numbers. The higher the group number, the smaller the atomic size. Thus, within period 3, S is larger than Se and within period 5, Sb is larger than Pb.

Arrange the Elements According to the Observed Trends

Based on the trends observed, arrange the elements from smallest to largest atomic size. S is smaller than Se, and both are smaller than Sb, which is smaller than Pb. Cs, having the largest atomic radius, goes last. So the order from smallest to largest is: S < Se < Sb < Pb < Cs.

Key concepts

Periodic Table Trends

Understanding the periodic table trends is a fundamental step in mastering the principles of chemistry. The periodic table is systematically organized to display patterns, which include atomic size trends, reactivity, electronegativity, and more.

Specifically, when discussing atomic size, we focus on two main trends: the variation of atomic size down a group, and the change of atomic size across a period. As you move down a group, from top to bottom, atomic size increases. This happens because each successive element has an additional electron shell, thereby increasing the distance between the outermost electrons and the nucleus. On the contrary, as you move from left to right across a period, atomic size decreases. This is due to the increasing nuclear charge with each successive element; as protons are added to the nucleus, electrons are pulled closer to the nucleus, resulting in a smaller atomic radius.

To visualize these concepts better, picture the periodic table like a matrix with rows (periods) and columns (groups). Elements within the same column share similar properties and the number of electron shells, while elements in rows have the same number of electron layers but vary in nuclear charge and size.

Atomic Radius

Delving into the concept of atomic radius helps us quantify the size of atoms. The atomic radius is defined as the distance from the center of the nucleus to the boundary of the surrounding electron cloud. Due to the nature of electron orbitals, this measurement isn't definite but rather an average or an estimated value.

Various factors affect the atomic radius, notably the number of electron shells and the effective nuclear charge. As the number of electron shells increases, the radius of an atom expands, which means atoms become larger as you move down a group in the periodic table. Meanwhile, a higher effective nuclear charge, which is the net positive charge experienced by valence electrons, can cause the electron cloud to be pulled tighter around the nucleus, leading to a smaller atomic radius.

For example, in the given exercise, cesium (Cs) has the largest atomic radius due to its position in group 1 and period 6, while sulfur (S), located in group 16 and period 3, has a smaller atomic radius because it has fewer electron shells and a smaller nuclear charge compared to Cs.

Electron Shells

Understanding electron shells is essential for grasping why atoms have different sizes. An electron shell is a grouping of electrons surrounding the nucleus of an atom that are bound at approximately the same energy level. The further the shell is from the nucleus, the higher the energy level. Following the principle of quantum mechanics, electrons fill these shells in a specific order, starting from the shell closest to the nucleus and moving outward.

Each electron shell can hold a maximum number of electrons: the first shell can hold up to 2 electrons, while the second and third can hold up to 8 and 18, respectively. As elements increase in atomic number, more electron shells are needed to accommodate the electrons. Therefore, when you compare the elements within the same group, such as sulfur (S) and selenium (Se) from the original exercise, the principle shows that Se, being in a lower period, will have more electron shells than S, and thus, a larger atomic size.

To illustrate, imagine the nucleus as the sun and the electron shells as orbits around it. Planets (electrons) on closer orbits (inner shells) are held more tightly by the sun's gravity (nuclear charge), while those on distant orbits (outer shells) are further away and less tightly bound.