ICH 101 Assignment Solutions To Be Written in 20 Leaves
Number 1
Diagonal relationship in the periodic table refers to the similarity in properties between certain elements that are located diagonally to each other. These elements are usually found in different periods and groups but exhibit comparable characteristics due to their similar electronic configurations. This phenomenon occurs because elements that are diagonally positioned share similar effective nuclear charges and electron configurations, resulting in analogous chemical behaviors.
Let's explore two examples of diagonal relationships in the periodic table:
Lithium (Li) and Magnesium (Mg):
Lithium is located in Group 1, Period 2, while Magnesium is found in Group 2, Period 3. Although they belong to different groups and periods, they share some similarities due to their diagonal relationship. Both elements have a +2 oxidation state, and their compounds display similar chemical properties. For instance, lithium oxide (Li2O) and magnesium oxide (MgO) are both basic in nature. Additionally, lithium hydroxide (LiOH) and magnesium hydroxide (Mg(OH)2) are both sparingly soluble in water.
Beryllium (Be) and Aluminium (Al):
Beryllium is found in Group 2, Period 2, whereas Aluminium is located in Group 13, Period 3. Despite their distinct positions in the periodic table, they exhibit some resemblances due to their diagonal relationship. Both elements have a tendency to form covalent compounds. Beryllium forms the covalent compound beryllium chloride (BeCl2), whereas aluminium forms the covalent compound aluminium chloride (AlCl3). Additionally, both elements have a relatively small ionic radius and exhibit amphoteric behavior, meaning they can act as both acids and bases.
Number 2
Electron affinity refers to the energy change that occurs when an atom or ion in the gas phase accepts an electron to form a negatively charged ion (anion). It is a measure of the attraction between the incoming electron and the nucleus of the atom.
The electron affinity of an element can vary across the periodic table due to several factors, including the effective nuclear charge, atomic radius, and electron configuration. The electron affinity tends to increase from left to right across a period and decrease from top to bottom within a group.
Here are some key variables that influence electron affinity across the periodic table:
Effective Nuclear Charge (Zeff): Electron affinity generally increases as the effective nuclear charge experienced by the valence electrons increases. As you move from left to right across a period, the number of protons in the nucleus increases, resulting in a higher effective nuclear charge. This increased nuclear charge attracts the incoming electron more strongly, leading to a higher electron affinity.
Atomic Radius: The size of an atom also affects its electron affinity. As the atomic radius decreases from left to right across a period, the valence electrons are closer to the nucleus, leading to a stronger attraction and higher electron affinity. Conversely, as you move down a group, the atomic radius increases, causing the valence electrons to be farther from the nucleus, resulting in a weaker attraction and lower electron affinity.
Electron Configuration: The stability of the resulting anion can impact the electron affinity. Elements with nearly filled or fully filled orbitals have higher electron affinities since gaining an electron allows them to achieve a more stable electronic configuration. For example, the halogens (Group 17) have high electron affinities as they require only one additional electron to achieve a stable noble gas configuration.
It's important to note that while these trends generally hold across the periodic table, there may be exceptions or anomalies due to other factors at play, such as electron-electron repulsion or electron pairing effects.
In summary, electron affinity is the energy change associated with the addition of an electron to an atom or ion. Its variation across the periodic table is influenced by factors like effective nuclear charge, atomic radius, and electron configuration.
Number 3
The periodic table can be represented in two main forms: the short form and the long form. Here's a comparison of these two forms:
Short Form of the Periodic Table:
The short form of the periodic table, also known as the standard form, is the most commonly used representation. It is a compact version that condenses the elements into a concise table. In this form, the elements are arranged in rows (periods) and columns (groups) based on their atomic numbers and chemical properties.
Key features of the short form of the periodic table include:
Elements are represented by their atomic symbols (one or two-letter abbreviations).
The table consists of seven horizontal rows called periods, labeled from 1 to 7.
The table is divided into 18 vertical columns called groups or families, numbered from 1 to 18. Some groups have specific names, such as the alkali metals (Group 1) and the halogens (Group 17).
Elements within the same group share similar chemical properties due to having the same number of valence electrons.
Transition metals are located in the middle of the table, spanning groups 3 to 12.
Lanthanides and actinides, known as the inner transition metals, are often placed below the main body of the table.
Long Form of the Periodic Table:
The long form of the periodic table, also called the extended form, provides additional information and details about the elements. It includes extra rows, known as f-block, to accommodate the lanthanide and actinide series.
Key features of the long form of the periodic table include:
It includes additional rows, often placed below the main body of the table, to incorporate the lanthanides (elements 57 to 71) and the actinides (elements 89 to 103).
The long form extends the periods to include all the elements, including the lanthanide and actinide series.
The numbering of the periods extends beyond 7 to accommodate the additional rows.
The f-block elements are sometimes presented separately at the bottom of the table, placed outside the main body.
The long form provides a more comprehensive view of the periodic table, allowing for the inclusion of all known elements and a better understanding of the relationships between them. However, due to its larger size and complexity, the short form is more commonly used for day-to-day reference and study purposes.
