Problem 98 - A Complete Resource Book in Chemistry for JEE Main

Question

Among the elements \(\mathrm{W}, \mathrm{X}, \mathrm{Y}\) and \(\mathrm{Z}\) having atomic numbers \(9,10,11\) and 12 respectively, the correct order of ionization energies is (a) \(\mathrm{W}>\mathrm{Y}>\mathrm{X}>\mathrm{Z}\) (b) \(\mathrm{X}>\mathrm{W}>\mathrm{Z}>\mathrm{Y}\) (c) \(X>Z>Y>W\) (d) \(\mathrm{Z}>\mathrm{Y}>\mathrm{X}>\mathrm{W}\)

Step-by-Step Solution

Answer

(b) \( \mathrm{X} > \mathrm{W} > \mathrm{Z} > \mathrm{Y} \).

Step 1: Identify the elements

The atomic numbers given are 9, 10, 11, and 12. These correspond to the elements Fluorine (F), Neon (Ne), Sodium (Na), and Magnesium (Mg) respectively. Therefore, \( \mathrm{W} \) is F, \( \mathrm{X} \) is Ne, \( \mathrm{Y} \) is Na, and \( \mathrm{Z} \) is Mg.

Step 2: Recall periodic trends

Ionization energy increases across a period (from left to right) and decreases down a group (from top to bottom). Among the elements identified, Neon (Ne) should have the highest ionization energy as it is a noble gas.

Step 3: Apply ionization energy trend

The elements ordered by increasing atomic number are F (W), Ne (X), Na (Y), and Mg (Z). Based on their positions in the periodic table: - Neon has the highest ionization energy among these because it is a noble gas. - Fluorine follows Neon because it is a highly electronegative halogen. - Magnesium follows next because it's in the same period but to the right of Sodium. - Sodium, being an alkali metal, has the lowest ionization energy.

Step 4: Determine the correct order

Based on the ionization energy trend, the order should be: F (\( \mathrm{W} \)) > Ne (\( \mathrm{X} \)) > Mg (\( \mathrm{Z} \)) > Na (\( \mathrm{Y} \)). This matches none of the given options, suggesting a re-evaluation of placement, factoring Ne being much higher than others. Reconsideration verifies: \( \mathrm{X} > \mathrm{W} > \mathrm{Z} > \mathrm{Y} \), which matches option (b).

Key Concepts

Periodic Table Trends

The periodic table is like a roadmap for understanding the elements and their properties. One of the key trends you’ll often hear about is how properties change as you move across periods (rows) and down groups (columns). Ionization energy, the energy required to remove an electron from an atom, is one of these trends. - Across a period from left to right, ionization energy generally increases. This is because more protons in the nucleus attract the electrons more strongly, making them harder to remove. - Down a group, the ionization energy decreases as the electrons are further from the nucleus and feel less attractive force from the protons. Understanding these trends is essential because they help predict how an element might react chemically.

Atomic Numbers

The concept of atomic numbers is fundamental in chemistry. Every element on the periodic table has a unique atomic number which represents the number of protons present in the nucleus of an atom for that element. - The atomic number is found above the element symbol in the periodic table. - Atomic numbers help to determine the position of elements within the table, hence predicting their properties and behaviors. For example, in our exercise, the atomic numbers 9, 10, 11, and 12 represent Fluorine (F), Neon (Ne), Sodium (Na), and Magnesium (Mg), respectively. This sequential order tells us about their chemical relations and helps us infer their ionization energies based on periodic trends.

Electronegativity

Electronegativity is an element's ability to attract and hold on to electrons. Although it is different from ionization energy, understanding electronegativity can give us insights into an element's chemical behavior. - Elements with high electronegativity, such as Fluorine, attract electrons strongly. - Electronegativity tends to increase across a period and decrease down a group, similar to ionization energy. Fluorine is one of the most electronegative elements and often gains electrons easily to form stable compounds, which is why it is placed after Neon in order of ionization energy despite its distinct electronegativity levels.

Noble Gases

Noble gases compose group 18 of the periodic table and are known for their lack of reactivity due to having full valence electron shells. This completion of their outer shell makes them very stable and gives them high ionization energies. - Noble gases include elements like Neon, which require a lot of energy to remove an electron because of this stability. - Because their valence shells are full, they do not seek electrons from other elements, nor do they easily lose their own. In our exercise context, Neon, being a noble gas, naturally has the highest ionization energy among the compared elements due to these unique characteristics.