The long form of the periodic table, also known as the modern periodic table, is better than Mendeleev's periodic table in several ways, primarily due to its ability to accurately represent the relationships between the properties of elements and their atomic structure. Here are some key reasons why the modern periodic table is superior to Mendeleev's periodic table:

1. Organization by Atomic Number: The modern periodic table is arranged based on the increasing atomic number of elements. This arrangement reflects the number of protons in the nucleus of each atom, which determines the chemical properties of the element. In contrast, Mendeleev's periodic table was organized primarily by atomic mass, which sometimes led to anomalies where elements did not fit into the predicted patterns.

2. Accurate Prediction of Element Properties: The modern periodic table accurately predicts the properties of elements based on their electronic configurations and atomic structure. The arrangement of elements in periods and groups reflects their similarities in chemical behavior and allows for easy identification of trends in properties such as electronegativity, ionization energy, and atomic radius. Mendeleev's periodic table, while groundbreaking for its time, did not always accurately predict the properties of elements or account for inconsistencies in their behavior.

3. Inclusion of Noble Gases: The modern periodic table includes noble gases as a separate group, which was not present in Mendeleev's original periodic table. Noble gases have unique properties, including their inertness and full valence electron shells, which distinguish them from other elements. The inclusion of noble gases in the modern periodic table provides a more comprehensive representation of the elements and their properties.

4. Transition Metals and Inner Transition Metals: The modern periodic table effectively organizes transition metals and inner transition metals into specific regions within the table. Transition metals are placed in the d-block, while inner transition metals are located in the f-block. This organization reflects the unique electronic configurations and properties of these elements, which was not explicitly addressed in Mendeleev's periodic table.

5. Greater Detail and Information: The modern periodic table provides more detailed information about each element, including its atomic number, atomic mass, electron configuration, and chemical symbol. This comprehensive information allows scientists to study and understand the properties and behavior of elements in greater depth, facilitating advancements in chemistry and other scientific disciplines.

In summary, the long form of the periodic table, with its organization based on atomic number and accurate representation of element properties, provides a more systematic and comprehensive framework for understanding the relationships between elements and their behavior compared to Mendeleev's periodic table.

Mendeleev's periodic table, while groundbreaking and influential in the development of modern chemistry, had several drawbacks that eventually led to its modification and the development of the modern periodic table. Some of the key drawbacks include:

1. Use of Atomic Mass as the Organizing Principle: Mendeleev organized elements primarily by increasing atomic mass. While this arrangement generally worked well and allowed him to predict the properties of unknown elements, it led to inconsistencies and anomalies, such as the inversion of the order of certain elements to maintain similar chemical properties. Additionally, the discovery of isotopes and the inability to accurately measure atomic masses at the time caused some discrepancies in the arrangement.

2. Omission of Noble Gases: Mendeleev's periodic table did not include noble gases because they were not yet discovered at the time of its creation. As a result, the periodic table did not account for these elements with unique properties, such as their inertness and full valence electron shells.

3. Positioning of Hydrogen: Mendeleev faced challenges in determining the appropriate placement of hydrogen in the periodic table due to its unique properties. He ultimately placed hydrogen in Group 1 alongside alkali metals, even though hydrogen shares some similarities with Group 17 elements (halogens) as well.

4. Incompleteness and Gaps: The periodic table developed by Mendeleev contained gaps for elements that had not yet been discovered. While Mendeleev successfully predicted the properties of some of these missing elements, such as gallium and germanium, the periodic table was not fully comprehensive or systematic.

5. Limited Explanation of Periodicity: Mendeleev's periodic table provided a systematic arrangement of elements but did not offer a comprehensive explanation for the observed periodic trends, such as variations in atomic radius, electronegativity, and ionization energy across periods and groups.

6. Failure to Account for Electron Configuration: Mendeleev's periodic table did not consider the electronic configuration of atoms, which plays a crucial role in determining the chemical properties of elements. The modern periodic table organizes elements based on increasing atomic number and electron configuration, providing a more accurate representation of their properties and behavior.

Overall, while Mendeleev's periodic table laid the foundation for understanding the relationships between elements, its limitations and drawbacks necessitated the development of the modern periodic table, which is based on atomic number and provides a more systematic and comprehensive framework for organizing and studying the elements.

The alkali metals are located in Group 1 of the periodic table. The outermost electronic configuration of alkali metals can be represented as ns^1, where 'n' represents the principal quantum number of the valence shell.

Here's the outermost electronic configuration of alkali metals:

• Lithium (Li): 1s^2 2s^1
• Sodium (Na): 1s^2 2s^2 2p^6 3s^1
• Potassium (K): 1s^2 2s^2 2p^6 3s^2 3p^6 4s^1
• Rubidium (Rb): 1s^2 2s^2 2p^6 3s^2 3p^6 3d^10 4s^2 4p^6 5s^1
• Cesium (Cs): 1s^2 2s^2 2p^6 3s^2 3p^6 3d^10 4s^2 4p^6 4d^10 5s^2 5p^6 6s^1
• Francium (Fr): 1s^2 2s^2 2p^6 3s^2 3p^6 3d^10 4s^2 4p^6 4d^10 4f^14 5s^2 5p^6 5d^10 6s^2 6p^6 7s^1

Justification for their placement in Group 1 of the periodic table:

1. Similar Outermost Electron Configuration: All alkali metals have one electron in their outermost shell, specifically in the s orbital. This makes them highly reactive as they tend to lose this single electron to achieve a stable noble gas configuration with a complete outer shell.

2. Similar Chemical Properties: Alkali metals share similar chemical properties due to their similar outermost electronic configuration. They readily lose their outermost electron to form positively charged ions (cations) with a charge of +1. This common behavior is a defining characteristic of Group 1 elements.

3. Reactivity Trend: Reactivity increases as you move down Group 1. This trend can be attributed to the decreasing ionization energy down the group, which makes it easier for the outermost electron to be removed. This leads to more vigorous reactions with water and other substances.

4. Similar Physical Properties: Alkali metals exhibit similar physical properties such as softness, low melting points, and low densities. These properties arise from the weak metallic bonding and the tendency for the atoms to form metallic lattices.

In summary, the placement of alkali metals in Group 1 of the periodic table is justified by their similar outermost electronic configuration, resulting in similar chemical properties, reactivity trends, and physical properties, all of which are characteristic of Group 1 elements.