Question

What is the formal charge of the indicated atom in each of the following structures? (a) the central \(\mathrm{O}\) atom in \(\mathrm{O}_{3}\) (b) \(\mathrm{Al}\) in \(\mathrm{AlH}_{4}^{-}\) (c) \(\mathrm{Cl}\) in \(\mathrm{ClO}_{3}^{-}\) (d) \(\sin \sin ^{2} \theta^{-}\) (e) \(\mathrm{Cl}\) in \(\mathrm{Cl} \mathrm{F}_{3}\)

Short answer

The formal charges for the atoms in the given molecules and ions are: (a) 0, (b) -1, (c) -1, (d) -2, (e) 0.

Step-by-step solution

Calculate the formal charge of O in O3

Oxygen normally has 6 valence electrons. For O3, each oxygen atom is bonded to another oxygen (that's 2 bonding electrons). Each outer oxygen has six non-bonding electrons and the central one has four. So for the central oxygen, the formal charge = 6 - 4 - 1/2*2 = 0.

Calculate the formal charge of Al in AlH4-

Aluminum normally has 3 valence electrons, but in AlH4-, it's bonded to 4 hydrogen atoms (4 bonding electrons) and there are no non-bonding electrons. So, the formal charge = 3 - 0 - 1/2*4 = -1

Calculate the formal charge of Cl in ClO3-

Chlorine normally has 7 valence electrons, but in ClO3-, it's bonded to 3 oxygen atoms (6 bonding electrons as each bond is a double bond) and has one pair of non-bonding electrons. So the formal charge = 7 - 2 - 1/2*6 = -1.

Calculate the formal charge of S in S2-

Sulfur normally has 6 valence electrons. In S^2-, these electrons are all non-bonding. So, the formal charge = 6 - 6 - 1/2*0 = 0. (We add 2 for the extra electrons that sulfur has gained, giving a formal charge of -2.)

Calculate the formal charge of Cl in ClF3

Chlorine normally has 7 valence electrons, but in ClF3, it's bonded to 3 fluorine atoms (6 bonding electrons) and has one pair of non-bonding electrons. So, the formal charge = 7 - 2 - 1/2*6 = 0.

Key concepts

Valence Electrons

Valence electrons are the outermost electrons in an atom. These electrons are essential because they participate in chemical bonding. The number of valence electrons can influence an element's reactivity and the type of bonds it can form. You can find valence electrons by looking at an element's group number on the periodic table. For example:

• Oxygen (O) is in group 16 and typically has 6 valence electrons.
• Aluminum (Al) is in group 13 and typically has 3 valence electrons.
• Chlorine (Cl) is in group 17 and typically has 7 valence electrons.

These electrons help us determine the formal charge, which is crucial for understanding molecule stability and reactivity.

Lewis Structures

Lewis Structures are diagrams that represent the bonding between atoms and the lone pairs of electrons in a molecule. They help visualize how atoms connect and provide a way to understand molecule geometry and resonance. When drawing Lewis Structures:

• Count the total number of valence electrons available.
• Draw single bonds between connected atoms first.
• Distribute the remaining electrons to satisfy the octet rule (8 electrons) or duplet rule (2 electrons for hydrogen).

For example, ozone (O₃) has a Lewis Structure where the central oxygen atom forms bonds with the other two oxygen atoms. This helps determine bond types and the central atom's formal charge.

Molecular Geometry

Molecular geometry describes the 3D arrangement of atoms in a molecule. This geometry is determined by the number of bonding pairs and lone pairs around a central atom. VSEPR (Valence Shell Electron Pair Repulsion) theory is commonly used to predict molecular geometry based on the repulsion between electron pairs:

• Linear: occurs with two bonding pairs, like in a molecule of O₂.
• Tetrahedral: occurs when there are four bonding pairs, like in AlH₄⁻.
• Trigonal pyramidal: occurs with three bonding pairs and one lone pair, like in NH₃.

Understanding molecular geometry helps in predicting the behavior and interaction of molecules, which is essential for advanced chemistry studies.