Problem 33 - Chemistry and Chemical Reactivity

Question

Explain each answer briefly. (a) Place the following elements in order of increasing ionization energy: \(F, O,\) and \(S\) (b) Which has the largest ionization energy: \(\mathbf{O},\) S, or Se? (c) Which has the most negative electron affinity: \(\mathrm{Se}, \mathrm{Cl}\) or Br? (d) Which has the largest radius: \(\mathbf{O}^{2-}, \mathbf{F}^{-},\) or \(\mathrm{F} ?\)

Step-by-Step Solution

Answer

(a) S < O < F; (b) O; (c) Cl; (d) O^{2-}.

Step 1: Understanding Ionization Energy Order

Ionization energy generally increases across a period and decreases down a group in the periodic table. Therefore, within the elements F, O, and S, F is highest because it's furthest right in the same period as O; S is lower than both because it is further down in the periodic table. Hence, in increasing order: S < O < F.

Step 2: Analyzing Highest Ionization Energy

Comparing O, S, and Se within the periodic table, O is topmost, with the highest ionization energy, as ionization energy decreases down a group. Therefore, O has the largest ionization energy among the three.

Step 3: Identifying Most Negative Electron Affinity

Electron affinity becomes more negative from left to right across a period and bottom to top up a group. Among Se, Cl, and Br, Cl has the most negative electron affinity as it is the farthest right and highest in the group among them. Thus, Cl has the most negative electron affinity.

Step 4: Determining the Largest Radius

Atomic radius increases down a group and decreases across a period. For ions: adding electrons increases radius due to additional electron-electron repulsion. Therefore, O^{2-} > F^{-} > F. So, the largest radius is O^{2-}.

Key Concepts

Periodic Table Trends

The periodic table is an incredible tool for predicting various properties of elements, such as ionization energy, electron affinity, and atomic radius. These properties are not random; they follow certain predictable trends across and down the table.
For instance, **ionization energy**, which is the energy required to remove an electron from an atom, generally increases as you move from left to right across a period and decreases as you move down a group. This happens because elements on the right have more protons, increasing the nuclear charge and pulling the electrons closer, making them harder to remove.
**Electron affinity**, the energy change when an electron is added to an atom, becomes more negative as you move across a period. This is due to atoms wanting to achieve a noble gas electron configuration.
**Atomic radius**, or the size of an atom, decreases across a period as the added protons increase nuclear attraction and squeeze electrons closer to the nucleus. Conversely, atomic radius increases down a group as more electron shells are added.

Electron Affinity

Electron affinity refers to the energy change that occurs when an electron is added to a neutral atom. It measures the atom's ability to gain an electron.
Most nonmetals, particularly those towards the top right of the periodic table, have more negative electron affinities. This means they release more energy when gaining an electron, indicating a stronger attraction for an added electron.
Elements like chlorine have a highly negative electron affinity because they are close to achieving a full valence shell, and gaining an additional electron provides significant stability. While it generally trends towards increasing negativity across periods, the group trend is less pronounced due to the impact of atomic size and shielding.

Atomic Radius

The atomic radius is a measure of the size of an atom. It becomes crucial when determining properties such as the reactivity and bonding patterns of elements.
As you move across a period from left to right, the atomic radius decreases. This is because added electrons enter the same shell while protons increase, leading to more nuclear attraction pulling the electron cloud inward.
However, as you move down a group, the atomic radius increases. Each new row adds an additional electron shell, spreading electrons out further from the nucleus, which results in a larger atomic size.
For ions, adding electrons creates electron-electron repulsion, which increases the size; therefore, anions are typically larger than their neutral atoms.

Chemical Elements

Chemical elements are the building blocks of all matter. Each element is defined by its atomic number, which is the number of protons in its nucleus, determining the element's identity.
The periodic table organizes elements according to atomic number and groups them by similar properties.
Some elements are metals, characterized by their ability to conduct electricity and heat. Others are nonmetals, known for non-conductive properties. Between metals and nonmetals are metalloids, which possess a mix of the two types' characteristics.

• **Metals**: Typically solid, shiny, good conductors.
• **Nonmetals**: Mostly gases or brittle solids, poor conductors.
• **Metalloids**: Have properties of both metals and nonmetals, useful in electronics.

Each group and period on the periodic table exhibits predictable chemical behavior.