Question

Draw Lewis structures for the following ligands: (a) \(\mathrm{H}_{2} \mathrm{O} ;\) (b) \(\mathrm{CH}_{3} \mathrm{NH}_{2} ;\) (c) \(\mathrm{ONO}^{-} ;\) (d) SCN \(^{-}\).

Short answer

The Lewis structures for these ligands are: (a) H:O:H for H2O, (b) H3C-NH2 for CH3NH2, (c) O=N-O: for ONO-, and (d) S=C=N: for SCN-.

Step-by-step solution

H2O Lewis Structure

Begin with the water molecule. Oxygen atom has 6 valence electrons and each Hydrogen atom has 1 valence electron, for a total of 8 valence electrons. Arrange the atoms as H—O—H, then start using the valence electrons to form bonds and fill the outer shells. The Lewis structure for H2O is: \n H:O:H \n ..

CH3NH2 Lewis Structure

Next, draw the structure for CH3NH2. Recall that Carbon has 4, Nitrogen has 5, and Hydrogen has 1 valence electron. Since there are three Hydrogens bound to one Carbon and two Hydrogens bound to the Nitrogen, the total valence electrons would be equal to 14. One way to draw the structure would be: H3C-NH2, where each dash represents a pair of shared electrons.

ONO- Lewis Structure

In the ONO- molecule, there is 1 extra electron due to the negative charge. So Oxygen has 6, Nitrogen has 5, and plus the extra electron, gives a total of 18 electrons. The structure can be drawn like this: O=N-O:, where the last Oxygen atom has three pairs of lone electrons.

SCN- Lewis Structure

Finally, in the SCN- molecule, Carbon has 4, Sulfur has 6, and Nitrogen has 5 valence electrons; and again, there is 1 extra electron due to the negative charge. Hence, the total number of valence electrons is 16. The structure is best represented as: S=C=N:, where both Sulfur and Nitrogen atoms have two pairs of lone electrons and Carbon has no lone pairs.

Key concepts

Valence Electrons

Valence electrons are the electrons that reside in the outermost shell of an atom. These electrons play a crucial role in chemical bonding as they are involved in forming bonds between atoms. Each type of atom has a specific number of valence electrons:

• Hydrogen has 1 valence electron.
• Oxygen has 6 valence electrons.
• Carbon has 4 valence electrons.
• Nitrogen has 5 valence electrons.
• Sulfur has 6 valence electrons.

When drawing Lewis structures, it's essential to count the total number of valence electrons available, with adjustments for any charges on the molecule.
For example, the ONO^- ion has an extra electron due to the negative charge, which means we add one more to the total count.

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms in a molecule. It is determined by the electron pairs surrounding a central atom. The water molecule (H₂O) is a classic example; it has a bent or angular shape due to the presence of two lone pairs on oxygen.
Different shapes arise from different combinations of bonded and lone pairs:

• Linear geometry occurs in molecules with no lone pairs and only two bonded atoms, like SCN^-.
• Tetrahedral geometry is common when there are four bonding pairs surrounding a central atom. CH₃NH₂ showcases this arrangement, considering hybridization effects.
• Trigonal planar or bent shapes can occur depending on any lone pairs present.

Understanding molecular geometry is important because it affects physical properties and reactivity.

Bonding Pairs

Bonding pairs of electrons are pairs of valence electrons that are shared between atoms and form a covalent bond. Each bond involves two electrons. When drawing Lewis structures, lines between atoms represent these bonding pairs.
In the molecule H₂O, each H-O bond results from one bonding pair. For CH₃NH₂, the bonds between carbon and hydrogen (C-H) and carbon and nitrogen (C-N) represent shared pairs of electrons. In ONO^-, the bonding between nitrogen and the two oxygens also involves shared pairs.
Identifying bonding pairs is crucial because they determine the connectivity and ultimate layout of the molecule.

Lone Pairs

Lone pairs, also called non-bonding pairs, are valence electrons that are not involved in bonding. They belong exclusively to one atom. Lone pairs affect the shape of the molecule and can change its reactivity.

• In H₂O, the oxygen atom has two lone pairs. These pairs repel each other, resulting in the bent shape.
• In ONO^-, the terminal oxygen has three lone pairs, impacting its resonance structures and molecular stability.
• In SCN^-, both sulfur and nitrogen possess two lone pairs each, which influence bond angles and polarity.

Lone pairs are visually represented in Lewis structures as dots around the atom they belong to. Understanding lone pairs provides insight into many chemical properties and behaviors.