Periodic Table JEE Mains Questions

Explanation: The question asks which element among the given options shows the least tendency to accept an additional electron, meaning it releases the smallest amount of energy when gaining one.

Electron affinity measures how much energy is released when an Atom gains an electron in the gaseous state. Typically, non-Metals have high electron affinity, while noble gases have very low values because their outer shells are already complete. Across a period, electron affinity generally increases, while it decreases down a group due to increased atomic size and shielding.

To determine the lowest electron affinity, consider the nature of the elements provided. Halogens such as fluorine, chlorine, and bromine are highly reactive and readily gain an electron to complete their valence shell, resulting in high electron affinity. In contrast, noble gases already possess a stable electronic configuration with a full outer shell. Adding another electron would disturb this stability, making them highly resistant to electron gain. Therefore, the element with a completely filled valence shell will exhibit the lowest electron affinity among the given options.

Think of atoms like seats in a bus. If a seat is empty, it is easy for someone to occupy it. But if the bus is already full, no one can easily get in.

Elements with completely filled outer shells resist gaining electrons, resulting in very low electron affinity compared to reactive non-Metals.

Explanation: This question asks for the term that describes how strongly an Atom attracts electrons when they are shared between atoms in a chemical bond.

In Chemical Bonding, atoms often share electrons to achieve stability. The strength with which an Atom pulls these shared electrons depends on factors such as its size and nuclear charge. Smaller atoms with higher effective nuclear charge attract electrons more strongly. This property is important in determining bond polarity and Molecular characteristics.

To solve this, it is important to distinguish between similar concepts. Electron affinity refers to energy change when an isolated Atom gains an electron, and ionization energy refers to removing an electron. However, the question specifically involves atoms within a compound, meaning the electrons are shared in a bond. Therefore, the required concept must describe the attraction exerted on shared electron pairs. This property varies periodically, increasing across a period and decreasing down a group, reflecting changes in atomic size and nuclear charge.

Imagine two people holding a rope. The stronger person pulls it closer, just like an Atom with stronger attraction pulls shared electrons toward itself.

This concept explains how strongly an Atom attracts shared electrons in a bond, influencing the nature and polarity of chemical compounds.

Explanation: The question asks which property shows an opposite trend compared to electronegativity, meaning as one increases, the other decreases.

Electronegativity is the ability of an Atom to attract shared electrons in a bond. It generally increases across a period and decreases down a group. Elements with high electronegativity are usually non-Metals, while those with low electronegativity are Metals. Metallic character is associated with the tendency to lose electrons rather than attract them.

To identify the inversely related property, observe how different characteristics behave across the Periodic Table. As electronegativity increases from left to right, metallic character decreases. Conversely, as we move down a group, electronegativity decreases while metallic character increases. This opposite behavior indicates an inverse relationship. Therefore, the property that reflects a tendency opposite to electron attraction—such as ease of losing electrons—will show an inverse trend relative to electronegativity.

Think of it like a tug-of-war: if one side pulls strongly (high electronegativity), the other side’s ability to give up easily (metallic nature) weakens.

Properties related to losing electrons increase when electronegativity decreases, showing a clear inverse relationship across Periodic trends.

Explanation: This question asks about the general position of metallic elements in the Periodic Table based on their chemical properties.

Metals are elements that tend to lose electrons easily and form positive ions. They have characteristics such as good conductivity, malleability, and ductility. In the Periodic Table, elements are arranged based on atomic number, and their properties change systematically across periods and groups.

To determine their position, observe Periodic trends. As we move from left to right across a period, metallic character decreases and non-metallic character increases. This is because atoms on the left side lose electrons more easily, while those on the right tend to gain electrons. Therefore, Metals are predominantly found on one side of the table, while non-Metals occupy the opposite side. The transition Metals also lie in between but still exhibit metallic behavior.

Imagine a gradient where one side represents strong electron donors and the other represents strong electron acceptors. Metals occupy the region where atoms readily give away electrons.

Thus, metallic elements are concentrated in a specific region of the Periodic Table where electron loss is most favorable.

Explanation: This question examines how the reactivity of non-metals changes as we move down a group in the Periodic Table.

Non-metals react mainly by gaining electrons to achieve a stable electronic configuration. Their reactivity depends on how easily they can attract and gain electrons. Factors such as atomic size, nuclear charge, and shielding effect influence this behavior.

As we move down a group, the atomic size increases due to the addition of new electron shells. This increased distance between the nucleus and the valence electrons reduces the effective nuclear attraction for incoming electrons. As a result, the ability of atoms to gain electrons decreases. Since reactivity in non-metals is closely tied to their electron-gaining tendency, a reduction in this tendency leads to lower reactivity down the group.

Think of trying to pull an object from a distance—the farther it is, the weaker the pull becomes. Similarly, larger atoms attract electrons less effectively.

Thus, the trend in reactivity for non-metals is governed by decreasing attraction for electrons as atomic size increases down the group.

Explanation: The question asks to determine the correct sequence of elements arranged according to their reactivity.

Reactivity of metals depends on their tendency to lose electrons. Metals that lose electrons easily are more reactive. This tendency is influenced by ionization energy, atomic size, and shielding effect. Larger atoms with lower ionization energy tend to lose electrons more readily.

To arrange elements in order of reactivity, compare their positions in the periodic table. Alkali metals are highly reactive because they have a single valence electron that is easily lost. As we move down a group, reactivity increases due to increased atomic size and reduced nuclear attraction. Transition metals and less reactive metals have higher ionization energies, making them less reactive.

Think of reactivity like willingness to give away something valuable—the easier it is to give, the more reactive the element.

By comparing electron loss tendencies and periodic positions, a proper sequence of reactivity can be established among the given elements.

Explanation: This question asks which metal among the given options shows the highest reactivity based on its tendency to lose electrons.

Metal reactivity is determined by how easily an Atom can lose its valence electron. Factors such as atomic size and ionization energy play a key role. Larger atoms with lower ionization energy lose electrons more easily, making them more reactive.

In the periodic table, reactivity of metals increases down a group because the outer electron is farther from the nucleus and less tightly held. Alkali metals are among the most reactive elements due to having a single valence electron. As we move down this group, reactivity increases further.

Imagine holding an object loosely versus tightly—the loosely held object is easier to drop. Similarly, atoms with loosely held electrons react more readily.

Thus, identifying the most reactive metal involves recognizing which element has the weakest hold on its outermost electron.

Explanation: This question focuses on identifying which alkali metal has the highest melting point despite the general trend of low melting points in this group.

Alkali metals have one valence electron and relatively weak metallic Bonding compared to other metals. As atomic size increases down the group, the strength of metallic Bonding decreases because atoms are held together less tightly.

To determine the highest melting point, consider how Bonding strength varies. Smaller atoms have stronger attraction between nuclei and the delocalized electrons, resulting in stronger metallic bonds. As we move down the group, atomic size increases and Bonding becomes weaker, lowering melting points.

Think of tightly packed particles requiring more energy to separate compared to loosely packed ones.

Thus, the element with the smallest atomic size among alkali metals will exhibit the strongest Bonding and therefore the highest melting point.

Explanation: The question evaluates understanding of periodic trends by asking which statements correctly describe variations in ionization energy, electron affinity, and electronegativity.

Ionization energy generally increases across a period due to increasing nuclear charge. Electron affinity trends vary but often decrease down a group because of increasing atomic size and shielding. Electronegativity typically increases across a period and decreases down a group.

To analyze the statements, compare each with known periodic trends. Across a period, stronger nuclear attraction increases ionization energy rather than decreasing it. Down a group, increased distance and shielding reduce the tendency to gain electrons. Across a period, atoms attract electrons more strongly, leading to increasing electronegativity.

Think of the nucleus as a magnet: the stronger and closer it is, the more it pulls electrons.

Evaluating each statement against these trends helps identify which ones align with established periodic behavior.

Explanation: This question asks which pair of elements exhibits a similar ratio of ionic charge to ionic size, an important factor in determining chemical similarity.

The charge-to-size ratio, also known as ionic potential, influences properties such as polarization and Bonding behavior. Ions with similar ratios tend to show comparable chemical characteristics. This concept is especially useful in understanding diagonal relationships in the periodic table.

To solve this, consider both the ionic charge and the size of the ions formed by the elements. Smaller ions with higher charges have greater charge density. Sometimes, elements from different groups show similarities because their ionic sizes and charges balance out in a similar way. This is often observed between certain pairs of elements placed diagonally in the periodic table.

Think of it like comparing density—different objects can behave similarly if their Mass-to-volume ratio is alike.

By evaluating both charge and size together, it becomes possible to identify pairs with similar ionic behavior.

Explanation: The question asks which types of atomic radii accurately represent the overall size of an atom.

Atomic size can be measured in different ways depending on how atoms interact with each other. Common types include covalent radius, metallic radius, and van der Waals radius. Each type is defined based on specific Bonding or interaction conditions.

To determine which radii reflect an increase in atomic size, consider how each measurement relates to the physical extent of an atom. Covalent radius is based on bonded atoms, metallic radius applies to metals in a lattice, and van der Waals radius represents non-bonded interactions. All these radii provide a measure of atomic size under different conditions. However, some are more commonly used to reflect practical changes in atomic dimensions.

Imagine measuring a person’s size in different ways—height, arm span, or personal space—all give slightly different but related values.

Understanding how each radius is defined helps identify which ones effectively represent changes in atomic size.

Explanation: This question explores why two elements from different groups in the periodic table show similar chemical behavior.

Lithium and magnesium exhibit what is known as a diagonal relationship. This occurs when elements diagonally adjacent in the periodic table share similar properties due to a balance of factors like atomic size and charge density.

To understand this similarity, consider both ionic size and charge. Lithium forms a small ion with a certain charge, while magnesium forms a slightly larger ion with a higher charge. These differences balance out to produce a similar charge-to-size ratio. As a result, both elements show comparable polarizing power and similar chemical behavior in compounds.

Think of two different objects having the same density despite differences in size and Mass.

Thus, their similarity arises from a balance between ionic charge and size, leading to comparable chemical properties despite different group positions.

Explanation: The question asks how the size of an atom changes when it loses an electron in one case and gains an electron in another case.

Atomic size depends on the balance between nuclear attraction and electron-electron repulsion. When an atom loses or gains electrons, this balance changes significantly, affecting the overall size of the resulting species.

When an electron is removed, the number of electrons decreases while the nuclear charge remains the same. This increases the effective nuclear attraction on the remaining electrons, pulling them closer to the nucleus and reducing the size. In contrast, when an electron is added, electron-electron repulsion increases because more electrons are present in the same space. This repulsion pushes electrons farther apart, increasing the size of the atom. Additionally, the effective nuclear pull per electron decreases in this case.

Think of it like a group sharing space—fewer members allow tighter packing, while adding more members causes crowding and expansion.

Thus, removal of an electron leads to contraction, while addition of an electron leads to expansion of atomic size.