Question

Arrange the following species in order of decreasing radius. (a) \(\mathrm{C}, \mathrm{Mg}, \mathrm{Ca}, \mathrm{Si}\) (b) \(\mathrm{Sr}, \mathrm{Cl}, \mathrm{Br}, \mathrm{I}\)

Short answer

Arrange the given species in the order of decreasing atomic radius: (a) C, Mg, Ca, Si (b) Sr, Cl, Br, I Answer: (a) Ca > Mg > Si > C (b) Sr > I > Br > Cl

Step-by-step solution

Identify the positions of elements in the periodic table

Find the positions of the given elements in the periodic table to understand their relationship in terms of periods and groups. (a) C, Mg, Ca, Si: C and Si are in Group 14 (Carbon group) Mg and Ca are in Group 2 (Alkaline Earth Metals) Ca is right below Mg in the periodic table. (b) Sr, Cl, Br, I: Sr is in Group 2 (Alkaline Earth Metals). Cl, Br, and I are in Group 17 (Halogens). Br is below Cl, and I is below Br in the periodic table.

Arrange elements based on group and period trends

Based on the trends in atomic radius, arrange the elements in each set. (a) C, Mg, Ca, Si: 1. Moving down a group (from Mg to Ca): Mg has a smaller atomic radius than Ca. 2. Moving right across a period (from C to Si): C has a larger atomic radius than Si. 3. Comparing groups (C/Si and Mg/Ca): Group 14 elements have smaller atomic radii than Group 2 elements. So, the order is: Ca > Mg > Si > C. (b) Sr, Cl, Br, I: 1. Moving down a group (for Cl, Br, and I): Cl > Br > I. 2. Comparing groups (Sr and Cl/Br/I): Sr has a larger atomic radius than Group 17 elements. So, the order is: Sr > I > Br > Cl.

Final Answer

The species are arranged in the order of decreasing radius as follows: (a) \(\mathrm{Ca} > \mathrm{Mg} > \mathrm{Si} > \mathrm{C}\) (b) \(\mathrm{Sr} > \mathrm{I} > \mathrm{Br} > \mathrm{Cl}\)

Key concepts

Periodic Table

The periodic table is a comprehensive chart that organizes the chemical elements based on their atomic number, electron configurations, and recurring chemical properties. Elements are listed in order of increasing atomic number in rows called periods. Columns, known as groups or families, contain elements with similar chemical behaviors. For example, Group 1 elements are known as alkali metals and share characteristics like high reactivity and one valence electron.

The table is structured such that elements with similar properties fall into the same column, while the rows roughly align with the addition of electron shells. As you move from left to right across a period, the atomic number increases, and elements gain more protons and electrons. In groups, as you move down, the number of electron shells increases. This structure is crucial in predicting an element's properties and its behavior in chemical reactions.

Atomic Radius

Understanding Atomic Radius

The atomic radius refers to the size of an atom, usually measured from the center of the nucleus to the boundary of the surrounding cloud of electrons. Since the exact position of the electron cloud is challenging to define, the atomic radius is often estimated based on how closely atoms are able to pack together.

The atomic radius matters in context because it affects how elements bond with each other and the stability of the compounds they form. This can be visually illustrated by comparing the size of different atoms on a graph or model, indicating that atomic size varies noticeably across different elements.

Group Trends

Variations Within Groups

The size of an atom generally increases as you move down a group in the periodic table. This trend occurs because each succeeding element has an additional electron shell, making the atom larger. As in our exercise, calcium (Ca) is larger than magnesium (Mg) because Ca is below Mg in Group 2, indicating it has more electron shells and hence a larger radius.The increase in atomic radius down a group also changes how the atoms interact with one another, influencing the chemical reactivity and the nature of the bonds they form. These trends are predictable and provide a basis for understanding the behavior of the elements in various chemical reactions.

Period Trends

Changes Across Periods

As one moves from left to right across a period in the periodic table, the atomic radius tends to decrease. This trend is attributed to the increase in the number of protons in the nucleus, which results in a stronger attraction between the positively charged nucleus and negatively charged electrons. This increased nuclear charge pulls the electron cloud closer to the nucleus, effectively decreasing the size of the atom.

For instance, comparing carbon (C) and silicon (Si), both in Group 14, Si is to the right of C in the same period; hence, Si has a smaller atomic radius due to a stronger nuclear charge. These periodic trends in the atomic radius are crucial for predicting and explaining various physical and chemical properties of elements, such as ionization energy, electronegativity, and metallic character.