Problem 29 - Chemistry and Chemical Reactivity

Question

Which of the following groups of elements is arranged correctly in order of increasing ionization energy? (a) \(\mathrm{C} < \mathrm{Si} < \mathrm{Li} < \mathrm{Ne}\) (c) \(\mathrm{Li} < \mathrm{Si} < \mathrm{C} < \mathrm{Ne}\) (b) \(\mathrm{Ne} < \mathrm{Si} < \mathrm{C} < \mathrm{Li}\) (d) \(\mathrm{Ne} < \mathrm{C} < \mathrm{Si} < \mathrm{Li}\)

Step-by-Step Solution

Answer

Option (c) \\( \text{Li} < \text{Si} < \text{C} < \text{Ne}\\) is correct.

Step 1: Understand Ionization Energy

Ionization energy is the amount of energy required to remove an electron from an atom in its gaseous state. Typically, ionization energy increases across a period from left to right and decreases down a group in the periodic table due to increasing atomic size.

Step 2: Analyze Ionization Trends

Elements within the same period generally have increasing ionization energy as you move from left to right. Elements within the same group generally have decreasing ionization energy as you move from top to bottom.

Step 3: Consider Periodic Table Position

Recall the periodic table arrangement: **Li** (Group 1, Period 2), **C** (Group 14, Period 2), **Si** (Group 14, Period 3), **Ne** (Group 18, Period 2). **Li** has the lowest ionization energy and **Ne** has the highest in this set.

Step 4: Apply Ionization Energy Rules

Arranging in order of increasing ionization energy based on their periodic positions: **Li** (lowest), **Si**, **C**, **Ne** (highest). This is valid because Li (period 2) has lower ionization energy compared to Si (below C in group 14), and Ne, being a noble gas, has the highest.

Step 5: Match with Options

The correct order as per increasing ionization energy from options provided is (c) **Li < Si < C < Ne**.

Key Concepts

Periodic Table Trends

The periodic table is a systematic arrangement of elements organized based on their atomic number, electron configurations, and recurring chemical properties. One of the main trends observed in the periodic table is the behavior of ionization energy. As you move across a period from left to right, the ionization energy typically increases. This increase is due to the addition of protons in the nucleus, leading to a stronger attractive force on the electrons. Moving down a group decreases ionization energy because electrons are farther away from the nucleus, experiencing less electrostatic pull.

• Increasing ionization energy across a period
• Decreasing ionization energy down a group

These trends help predict how elements will behave in reactions and allow us to understand the arrangement in terms of ionization energy, which is crucial for solving exercises like the one presented.

Electron Removal

Ionization energy is directly related to electron removal, referring to the energy needed to remove an electron from an atom. The process of electron removal is significant because it influences an element's chemical properties and reactivity. When an electron is removed, an atom becomes a positively charged ion, or cation. More precisely, the first ionization energy refers to removing the first electron, and it's typically easier than removing subsequent electrons. Several factors affect this:

• The atomic radius: Larger atomic radius means electrons are further from the nucleus, requiring less energy for removal.
• Nuclear charge: More protons lead to stronger attraction to electrons, increasing ionization energy.
• Shielding: Inner shell electrons repel outer electrons, making them easier to remove, lowering ionization energy.

Understanding electron removal helps clarify why some elements have higher ionization energy and are less reactive, as they hold onto their electrons tightly.

Atomic Structure

The atomic structure is the framework that explains many periodic trends, including ionization energy. It comprises the nucleus (protons and neutrons) and electrons arranged in shells around the nucleus. The arrangement of electrons, based on atomic structure, dictates how tightly electrons are held, influencing ionization energy. For instance, electrons in outer shells are held less tightly due to shielding from inner electrons, affecting ionization energy. Key points about atomic structure:

• Electrons closest to the nucleus are held more tightly due to less shielding.
• As effective nuclear charge increases across a period, electrons experience a stronger pull towards the nucleus.

Understanding atomic structure provides insight into why certain elements have a harder time losing electrons and thus have higher ionization energies. This is a fundamental reason why Ne has a higher ionization energy compared to Li, Si, and C.

Element Groups

Elements are classified into groups based on their chemical properties and electron configurations. These groups provide a framework to predict elements' behavior, including trends in ionization energy. For instance, Group 1 elements (alkali metals) have lower ionization energies, as losing an electron achieves a stable noble gas configuration. Noble gases (Group 18), however, have full outer electron shells, meaning they have high ionization energies since electron removal disrupts stability. Some important attributes of element groups:

• Alkali metals: Lowest ionization energies within a period.
• Noble gases: Highest ionization energies.
• Transition from metallic to non-metallic behavior across periods influences ionization energy trends, increasing from metals to non-metals.

Recognizing these group characteristics helps in determining a sequence of elements by ionization energy, like arranging in a given problem: Li, Si, C, Ne.