Question

Consider these elements: \(\mathrm{P}, \mathrm{Ca}, \mathrm{Si}, \mathrm{S}, \mathrm{Ga}\) a. Write the electron configuration for each element. b. Arrange the elements in order of decreasing atomic radius. c. Arrange the elements in order of increasing ionization energy. d. Use the electron configurations in part a to explain the differences between your answers to parts b and c.

Short answer

The electron configurations are: P: [Ne] 3s2 3p3, Ca: [Ar] 4s2, Si: [Ne] 3s2 3p2, S: [Ne] 3s2 3p4, Ga: [Ar] 3d10 4s2 4p1. Decreasing atomic radius order: Ca, Ga, Si, P, S. Increasing ionization energy order: Ca, Ga, Si, P, S. Ionization energy increases with a greater effective nuclear charge and fewer electron shells, while atomic radius decreases.

Step-by-step solution

Writing Electron Configurations

Write the electron configurations for the given elements. Use the periodic table to determine the number of electrons and their arrangement. The order of filling orbitals is based on increasing energy levels typically following the pattern '1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, etc.'

Ordering by Atomic Radius

Arrange the elements in order of decreasing atomic radius. Generally, atomic radius decreases across a period from left to right and increases down a group in the periodic table. The order will reflect the elements' positions on the periodic table.

Ordering by Ionization Energy

Arrange the elements in order of increasing ionization energy. Ionization energy typically increases across a period and decreases down a group in the periodic table. The arrangement will be based on the elements' ability to attract and hold electrons.

Explaining Differences

Explain the differences between atomic radius and ionization energy orders using the electron configurations obtained in Step 1. Consider the effect of the number of energy levels (shells) and effective nuclear charge on the atomic radius and ionization energy.

Key concepts

Atomic Radius

Understanding the concept of atomic radius is quintessential when diving into the realm of chemistry, particularly when exploring the periodic table. The atomic radius is a measure of the size of an atom, from the center of the nucleus to the outer boundary of the surrounding cloud of electrons. As we move across a period on the periodic table, the atomic radius usually decreases. This trend occurs because each successive element has an additional proton and electron, which means a greater effective nuclear charge attracts the electrons closer to the nucleus.

However, when we journey down a group, the atomic radius increases. This is due to each added electron shell which is further from the nucleus, thus increasing the overall size of the atom. Elements like Cesium (Cs), for example, which are located at the bottom of a group, will have a larger atomic radius compared to those at the top, such as Hydrogen (H).

An understanding of atomic radius can reveal much about an element's properties, such as reactivity and bonding behavior. For instance, elements with larger atomic radii tend to be more reactive in groups because the outer electrons are further away from the nucleus and more easily lost during reactions.

Ionization Energy

Ionization energy is a term that signifies the amount of energy required to remove an electron from an isolated gaseous atom or ion. It is an intrinsic property that reflects how tightly an atom holds on to its electrons. Generally, ionization energy increases as we move from left to right across a period. This trend is because elements on the right side of a period have more protons, thus exert a stronger attraction on their electrons, making it more difficult to remove one.

Conversely, as we descend down a group, ionization energy tends to decrease. The electrons being removed are found in orbitals further away from the nucleus and are also more shielded by the inner electrons from the pull of the nucleus. This results in these outer electrons being removed more easily, requiring less energy.

Key Factors Influencing Ionization Energy

• Effective Nuclear Charge: The net positive charge experienced by an electron in an atom. Higher effective nuclear charge increases ionization energy.
• Electron Shielding: As more electron shells are added, inner electrons can block the attractive force from the nucleus on the valence electrons, thus decreasing ionization energy.
• Electron Distance: Electrons further from the nucleus can be removed more easily, hence lower ionization energy.

Understanding these factors provides insight into the behavior of atoms during chemical reactions and why certain elements have higher or lower reactivities.

Periodic Table Trends

The periodic table is more than just a list of elements; it's a comprehensive map of chemical properties and trends that provide insights into the world of atoms. One of the fundamental trends observed is in atomic radii and ionization energies, as discussed previously. These trends are not arbitrary but are based on the effective nuclear charge, shielding effect, and the number of electron shells within atoms.

As we move across a period, effective nuclear charge increases due to more protons being added to the nucleus. This increased charge pulls the electrons closer, leading to a smaller atomic radius and higher ionization energy. On the flip side, down a group, additional electron shells are added, increasing the distance between the outer electrons and nucleus. This results in a larger atomic radius and lower ionization energy.

Additional Trends to Consider

• Electronegativity: Typically increases across a period and decreases down a group. It describes how strongly an atom can attract and hold onto bonding electrons.
• Metallic character: It generally decreases across a period and increases down a group. Metals tend to lose electrons easily, indicated by their low ionization energies.

These periodic table trends help predict and explain the chemical behavior of elements in various contexts, providing a foundation for understanding patterns in element properties and guiding scientists and students alike in anticipating how different elements will interact.