Question

Draw the Lewis structures for each of the following ions or molecules. Identify those in which the octet rule is not obeyed; state which atom in each compound does not follow the octet rule; and state, for those atoms, how many electrons surround these atoms: (a) \(\mathrm{PH}_{3}\), (b) \(\mathrm{AlH}_{3}\), (c) \(\mathrm{N}_{3}^{-}\), (d) \(\mathrm{CH}_{2} \mathrm{Cl}_{2}\), (e) \(\mathrm{SnF}_{6-}\)

Short answer

In summary, for the given ions and molecules: (a) PH3 - Phosphorus doesn't follow the octet rule with 6 valence electrons. (b) AlH3 - Aluminum doesn't follow the octet rule with 6 valence electrons. (c) N3- - All atoms follow the octet rule. (d) CH2Cl2 - All atoms follow the octet rule. (e) SnF6^2- - Tin doesn't follow the octet rule with 12 valence electrons.

Step-by-step solution

(a) Drawing Lewis structure for PH3

For PH3, we first count the total number of valence electrons: P has 5 valence electrons and each H has 1 valence electron, so there are a total of 8 valence electrons. Arrange the atoms with P in the center and connect H atoms to it. Distribute the remaining electrons to fulfill octet rule: P:H:H:H Phosphorus has 6 electrons in its valence shell.

(b) Drawing Lewis structure for AlH3

For AlH3, we first count the total number of valence electrons: Al has 3 valence electrons and each H has 1 valence electron, so there are a total of 6 valence electrons. Arrange the atoms with Al in the center and connect H atoms to it. There are no additional electrons to distribute: Al-H-H-H Aluminum has only 6 electrons in its valence shell.

(c) Drawing Lewis structure for N3-

For N3-, we first count the total number of valence electrons: Each N has 5 valence electrons, and there is an extra electron from the negative charge, so there are a total of 16 valence electrons. In this case, we use two N atoms to form a triple bond with other. The negative charge will reside on the lone N atom: [N---N=N]⁻ All nitrogen atoms follow the octet rule in this ion.

(d) Drawing Lewis structure for CH2Cl2

For CH2Cl2, we first count the total number of valence electrons: C has 4 valence electrons, each H has 1 valence electron, and each Cl has 7 valence electrons, so there are a total of 20 valence electrons. Arrange the atoms with C in the center and connect H and Cl atoms to it: H-C-H | | Cl-Cl All atoms follow the octet rule in this molecule.

(e) Drawing Lewis structure for SnF6^2-

For SnF6^2-, we first count the total number of valence electrons: Sn has 4 valence electrons, each F has 7 valence electrons, and there are two extra electrons from the negative charge, so there are a total of 46 valence electrons. Arrange the atoms with Sn in the center and connect Sn to each F. Use the remaining electrons to satisfy octet rules for F: F-F | | F-Sn-F | | F-F Tin has just 12 valence electrons.

Identifying non-octet atoms and their electron count

Based on the Lewis structures above, the following atoms don't follow the octet rule: (a) Phosphorus in PH3 has 6 valence electrons. (b) Aluminum in AlH3 has 6 valence electrons. (e) Tin in SnF6^2- has 12 valence electrons.

Key concepts

Octet Rule

The octet rule is a fundamental concept in chemistry that explains the tendency of atoms to prefer having eight electrons in their valence shell. This rule is a guiding principle for the formation of stable molecules and ions. Atoms achieve a full octet by either sharing, gaining, or losing electrons.

• Atoms like carbon, nitrogen, and oxygen strictly follow the octet rule.
• Some elements, such as phosphorus and sulfur, can have expanded octets.
• Elements like hydrogen and helium are exceptions, preferring a duo instead of an octet.

In the exercise, aluminum in AlH₃ and phosphorus in PH₃ do not fulfill the octet rule, as they have fewer than eight electrons; tin in SnF₆²⁻ exceeds it, having twelve.

Valence Electrons

Valence electrons are the electrons found in the outermost shell of an atom. These electrons play a crucial role in chemical bonding and the formation of compounds. They are easily accessible for interactions with other atoms due to their position.

• Varying across the periodic table, elements in the same group usually have the same number of valence electrons.
• For example, Group 1 elements like hydrogen have only one valence electron.
• Group 15 elements like nitrogen have five valence electrons.

In the given exercise, identifying the total number of valence electrons helps in drawing accurate Lewis structures and ensuring that atoms like nitrogen in N₃⁻ have a full octet.

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. It is determined by the types and numbers of bonds, as well as the presence of lone pairs of electrons.

• VSEPR theory (Valence Shell Electron Pair Repulsion) helps in predicting molecular shapes.
• Bond angles and the presence of lone pairs significantly influence the geometry.
• Common shapes include linear, trigonal planar, tetrahedral, and bent.

Regarding the molecules in the exercise, CH₂Cl₂ has a tetrahedral geometry, while PH₃ has a trigonal pyramidal shape, influenced by the lone pair on the phosphorus atom.

Ions

Ions are charged particles that form when atoms gain or lose electrons. The charge indicates whether electrons were lost (positive ion) or gained (negative ion). This conversion often occurs to achieve a stable electron configuration.

• Cations are positive ions, usually formed by metals losing electrons.
• Anions are negative ions, typically formed by nonmetals gaining electrons.
• Polyatomic ions, like N₃⁻, consist of more than one element bonded together, carrying a net charge.

In the exercise, N₃⁻ and SnF₆²⁻ are examples of polyatomic ions, with N₃⁻ taking an extra electron to complete its octet, resulting in a negative charge.

Covalent Bonding

Covalent bonding occurs when atoms share pairs of electrons to achieve a full outer electron shell, thereby reaching more stable electronic configurations. This type of bonding is common in molecules composed of nonmetals.

• Single, double, or triple covalent bonds can form based on the number of shared electron pairs.
• In single covalent bonds, one pair of electrons is shared, like in CH₂Cl₂.
• Triple covalent bonds, observed in N₃⁻, involve three pairs of shared electrons.

The exercise demonstrates how molecules and ions like PH₃ and N₃⁻ utilize covalent bonds to connect atoms and fill their valence shells, ensuring molecular stability.