Question

Arrange these elements in order of increasing atomic size: \(\mathrm{Ca}, \mathrm{Rb}, \mathrm{S}, \mathrm{Si}, \mathrm{Ge}, \mathrm{F}\).

Short answer

The order of increasing atomic size is F < S < Si < Ge < Ca < Rb.

Step-by-step solution

Understand Periodic Trends

The atomic size generally increases from top to bottom within a group (column) in the periodic table, because elements at the bottom have more electron shells compared to those at the top. Also, atomic size decreases from left to right across a period (row) because of the increase in effective nuclear charge, which pulls the electron cloud closer to the nucleus.

Locate Elements on the Periodic Table

Find the position of each given element on the periodic table. The elements are as follows: Calcium (Ca) is in the 2nd group and 4th period, Rubidium (Rb) is in the 1st group and 5th period, Sulfur (S) is in the 16th group and 3rd period,Silicon (Si) is in the 14th group and 3rd period,Germanium (Ge) is in the 14th group and 4th period,Fluorine (F) is in the 17th group and 2nd period.This will help in comparing their sizes.

Compare Atomic Sizes within Groups

Looking at elements within the same group, remember that the size increases as you move down the group. Thus, among Ca and Rb that are part of the same group (alkali earth metals), Rb is larger than Ca.

Compare Atomic Sizes within Periods

Now compare the sizes of elements in the same period. Remember that size decreases as you move from left to right across a period. This means that Si is larger than Ge (though they are in the same group, Ge is in the next period to the right), which is larger than S, as it is farther to the right.

Organize the Elements

With the above comparisons in mind, order the elements from the smallest to largest atomic size. Since we have elements that are in the same period, we start with them from the rightmost element (F), and then we continue with Si, Ge, and S within the same period, followed by Ca and Rb which are within the same group.

Key concepts

Understanding Periodic Trends

The concept of periodic trends refers to patterns in properties of elements across the periodic table. One such property is atomic size. As we observe the elements in the periodic table, atomic size tends to increase when moving down a group. This is because as you move down, elements have more electron shells, adding a layer of electrons and effectively increasing the distance from the nucleus.

Conversely, atomic size generally decreases from left to right across a period. The addition of protons to the nucleus increases the effective nuclear charge, which attracts the electrons more strongly and pulls them closer to the nucleus, thereby reducing the size of the atom.

To grasp these concepts better, visualize placing elements on a staircase where moving down a step increases size, and moving right a step decreases size - this will help simplify the periodic trends for atomic size!

Decoding the Periodic Table Arrangement

The periodic table is meticulously arranged to showcase a variety of chemical and physical properties. In terms of atomic size, the table is divided into rows called periods and columns known as groups. Groups share chemical characteristics, and in the context of size, elements within a group increase in size as you move down because each successive element has an additional electron shell.

Locating Elements

For example, Calcium (Ca) located in group 2 period 4, has fewer electron shells than Rubidium (Rb), which is also in group 2 but in period 5. Periods help us understand that within the same shell level, elements gain protons and electrons as you move from left to right, increasing effective nuclear charge and often leading to a decrease in size.

The Role of Electron Shells in Atomic Size

Electrons reside in regions around the nucleus called electron shells or energy levels. With each additional shell, electrons are further away from the nucleus, resulting in an increase in atomic size. Think of this as an expanding universe: just as stars and galaxies become more spread out, so do electron shells as you descend the periodic table.

However, the number of shells is not the sole factor in determining size; it's important to consider other influences such as effective nuclear charge, which can pull electrons closer and compact the atom's size even if the number of shells increases.

Effective Nuclear Charge

The concept of effective nuclear charge (ENC) can be thought of as the net positive charge experienced by an electron in a multi-electron atom. The more protons in the nucleus, the more an electron is pulled towards it. However, not all electrons feel this full charge due to shielding by other electrons.

Shielding Effect

Inner electrons can shield outer electrons from the full force of the nuclear charge. This results in outer electrons feeling a lesser pull towards the nucleus which is referred to as a lower effective nuclear charge. This phenomenon plays a crucial role in understanding why atoms do not simply become infinitely compact as you move across a period and why the trends in size across the periodic table are not strictly linear.