Question

Rank these ionization energies (IE) from the smallest to the largest value. Briefly explain your answer. (a) First \(I E\) of \(B e\) (b) First IE of \(\mathrm{Li}\) (c) Second IE of \(\mathrm{Be}\) (d) Second IE of Na (e) First IE of \(\mathrm{K}\)

Short answer

First IE of K < First IE of Li < First IE of Be < Second IE of Na < Second IE of Be.

Step-by-step solution

Understand Ionization Energy Concept

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. The first ionization energy refers to the energy needed to remove the first electron, while the second ionization energy is for removing the second electron. Generally, as you move across a period from left to right, ionization energy increases, and it decreases as you move down a group.

Identify Ionization Energies

List out the given ionization energies: - (a) First IE of Be - (b) First IE of Li - (c) Second IE of Be - (d) Second IE of Na - (e) First IE of K First ionization energy removes one electron, while second ionization energy removes the next electron, which is harder to remove due to increased nuclear charge for the remaining electrons.

Compare Trends Across Period and Group

Lithium (Li) and beryllium (Be) are in the same period. Beryllium's first IE is higher than lithium's because it has a smaller atomic radius. Sodium (Na) and potassium (K) are in the same group; sodium's outer electron is closer to the nucleus compared to potassium, thus Na's second IE will be higher than K's first IE.

Evaluate Successive Ionization Energies

The second ionization energy of an element is higher than the first since it involves removing an electron from a positively charged ion. Therefore, the second IE of Be is greater than the first IE of Be, and the second IE of Na will be higher than its first IE.

Rank the Ionization Energies

Considering the above points: 1. First IE of K 2. First IE of Li 3. First IE of Be 4. Second IE of Na 5. Second IE of Be The second ionization energies are generally larger than the first ionization energies of the same element because it is more difficult to remove electrons once an atom is positively charged.

Key concepts

Periodic Trends

Periodic trends are patterns in the periodic table that illustrate how certain properties of elements change. One of the key periodic trends is ionization energy, which demonstrates these patterns effectively. Generally, ionization energy increases as you move across a period from left to right. This is because the number of protons (and hence, the nuclear charge) increases, pulling the electrons closer to the nucleus and making them harder to remove.
On the other hand, as you move down a group, ionization energy decreases. This is due to the increase in atomic radius as electrons are added to new shells further from the nucleus. The increased distance from the nucleus means the outer electrons experience less electrostatic pull from the nucleus.
Understanding these trends helps predict the reactivity and properties of elements, making the study of periodic trends an essential part of chemistry.

Atomic Structure

Atomic structure refers to the arrangement of protons, neutrons, and electrons within an atom. The core idea is that an atom consists of a nucleus, containing positively charged protons and neutral neutrons, surrounded by negatively charged electrons in orbitals.
The number of protons in the nucleus determines the atomic number and defines the element itself. Electrons, which are negatively charged, are arranged in increasing energy levels (or shells) around the nucleus. The energy required to remove one of these electrons from an atom is its ionization energy.
Each electron shell can hold a specific maximum number of electrons. As the energy level increases, the distance from the nucleus also increases, which means electrons in outer shells are more easily removed. This forms a basis for the changes seen in ionization energy across different elements.

Electronegativity

Electronegativity is the measure of an atom’s ability to attract and hold onto electrons within a chemical bond. Unlike ionization energy, which focuses on the removal of electrons, electronegativity deals with the attraction of electrons.
In general, electronegativity increases from left to right across a period, because atoms have a stronger attraction for electrons as they near completing their valence shell. Conversely, electronegativity decreases down a group. This is because as atoms get larger, the increased distance between the nucleus and valence shell reduces the nucleus's ability to attract bonding electrons.
Factors such as atomic number and distance between electrons and the nuclear charge influence electronegativity. Understanding electronegativity is crucial for predicting the nature of bonds and understanding molecular structures.