The question asks about Dobereiner's Law of Triads, a historical attempt to classify elements. We need to recall the specific relationship Dobereiner proposed for the atomic weights of elements within a triad.
C) Average of the other two — This correctly describes Dobereiner's Law of Triads, where the atomic weight of the middle element was approximately the average of the atomic weights of the other two elements in the triad.
The question asks about the anomalous/dual behavior of hydrogen in the periodic table. This anomaly arises because hydrogen exhibits properties similar to both alkali metals (Group 1) and halogens (Group 17) due to its electronic configuration.
C) Both alkali metals and halogens — Hydrogen can lose one electron to form \(\text{H}^+\) like alkali metals, and it can gain one electron to form \(\text{H}^-\) like halogens. It also shares other properties with both groups, such as forming diatomic molecules (like halogens) and combining with electronegative elements (like alkali metals).
The question asks to identify the phenomenon where elements from different groups, like Lithium and Magnesium, exhibit similar chemical properties due to comparable charge/radius ratios. This is a specific concept in periodic trends.
Correct Option: D) Diagonal relationship — The diagonal relationship is a phenomenon in the periodic table where elements of the second period show similarities in properties with elements of the third period that are diagonally opposite to them. This similarity arises due to comparable ionic sizes and charge/radius ratios (polarizing power). Lithium and Magnesium are a classic example of this relationship.
To determine the trend of atomic radius across a period, we need to consider the changes in nuclear charge and the number of electron shells as we move from left to right in the periodic table.
A) Decreases. As explained above, moving from left to right across a period, the effective nuclear charge increases while the number of electron shells remains the same. This stronger attraction pulls the valence electrons closer to the nucleus, causing the atomic radius to decrease.
The question asks about the general trend of atomic radius as one moves down a group in the periodic table. To answer this, we need to understand the factors that influence atomic radius, particularly the number of electron shells and nuclear charge.
A) Increases — As explained above, moving down a group adds new electron shells, which are further from the nucleus and experience increased shielding, causing the atomic radius to increase.
The question asks about the general trend of ionization enthalpy (energy) across a period from left to right in the periodic table. Ionization enthalpy is the energy required to remove an electron from a gaseous atom. Understanding the factors that influence this energy, such as nuclear charge, atomic radius, and shielding effect, is crucial to determining the trend.
Definition of Ionization Enthalpy: Ionization enthalpy is the minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state to form a positive ion. For example:
\[ \text{X(g)} + \text{Energy} \to \text{X}^+\text{(g)} + \text{e}^- \]
Factors Affecting Ionization Enthalpy:
Trend Across a Period (Left to Right):
Conclusion on Ionization Enthalpy: Because the effective nuclear charge increases and the atomic radius decreases across a period, more energy is required to remove an electron from the outermost shell. Therefore, ionization enthalpy generally increases from left to right across a period.
B) Increases — As explained above, moving from left to right across a period, the effective nuclear charge increases and the atomic radius decreases, making it harder to remove an electron. Thus, ionization enthalpy generally increases.