Question

Write a Lewis structure for each of the following simple molecules. Show all bonding valence electron pairs as lines and all nonbonding valence electron pairs as dots. a. ${\mathrm{N}}_2{\mathrm{H}}_4$ b. ${\mathrm{C}}_2{\mathrm{H}}_6$ c. ${\mathrm{NCl}}_3$ d. ${\mathrm{SiCl}}_4$

Short answer

The Lewis structures for the given molecules are: a. ${\mathrm{N}}_2{\mathrm{H}}_4$: H - N: N - H | | H H b. ${\mathrm{C}}_2{\mathrm{H}}_6$ (ethane): H - C - C - H | | H H c. ${\mathrm{NCl}}_3$: Cl | N - Cl | Cl d. ${\mathrm{SiCl}}_4$: Cl | Cl - Si - Cl | Cl

Step-by-step solution

Determine the number of valence electrons

Nitrogen (N) belongs to Group 15 and has 5 valence electrons, and Hydrogen (H) belongs to Group 1 and has 1 valence electron. Since there are 2 Nitrogen and 4 Hydrogen atoms, the total valence electrons are (2*5) + (4*1) = 14.

Arrange the atoms

In ${\mathrm{N}}_2{\mathrm{H}}_4$, there are two Nitrogen atoms and four Hydrogen atoms. We need to arrange the atoms, placing the less electronegative atom at the center. So, N₂H₄ can be represented as H-N-N-H with 2 Hydrogen atoms bonded to each Nitrogen atom.

Form the bonds

We can form single covalent bonds between the Nitrogen and Hydrogen atoms and also between the two Nitrogen atoms. Each bond has two shared electrons, so we are left with 14 - (3*2) = 8 nonbonding electrons.

Distribute the remaining nonbonding electrons

The remaining 8 nonbonding electrons should be distributed around the Nitrogen atoms to satisfy the octet rule, which states that each atom (except Hydrogen) should have 8 electrons in its valence shell. Distributing 8 electrons equally between the 2 Nitrogen atoms, we have: H - N: N - H | | H H b. ${\mathrm{C}}_2{\mathrm{H}}_6$

Determine the number of valence electrons

Carbon (C) belongs to Group 14 and has 4 valence electrons, and Hydrogen (H) has 1 valence electron. The total valence electrons for ${\mathrm{C}}_2{\mathrm{H}}_6$ are (2*4) + (6*1) = 14.

Arrange the atoms

In ${\mathrm{C}}_2{\mathrm{H}}_6$, there are two Carbon atoms and six Hydrogen atoms. Arrange the atoms with Carbon as the central atom and the Hydrogen atoms around them: H-C-C-H.

Form the bonds and distribute remaining electrons

Form single covalent bonds between the Carbon and Hydrogen atoms and between the two Carbon atoms. We have used 8 electrons so far; the remaining 6 electrons will be used to complete the octet for each Carbon atom. The final Lewis structure for ${\mathrm{C}}_2{\mathrm{H}}_6$ (ethane) is: H - C - C - H | | H H c. ${\mathrm{NCl}}_3$

Determine the number of valence electrons

Nitrogen (N) has 5 valence electrons, and Chlorine (Cl) belongs to Group 17 and has 7 valence electrons. The total valence electrons for ${\mathrm{NCl}}_3$ are 5 + (3*7) = 26.

Arrange the atoms

In ${\mathrm{NCl}}_3$, there is one Nitrogen atom and three Chlorine atoms. Arrange the atoms with Nitrogen as the central atom surrounded by the Chlorine atoms: N(Cl)₃.

Form the bonds and distribute remaining electrons

Form single covalent bonds between the Nitrogen and Chlorine atoms. We have used 6 electrons so far, and the remaining 20 electrons should be distributed to satisfy the octet rule for all atoms. The final Lewis structure for ${\mathrm{NCl}}_3$ is: Cl | N - Cl | Cl d. ${\mathrm{SiCl}}_4$

Determine the number of valence electrons

Silicon (Si) belongs to Group 14 and has 4 valence electrons, and Chlorine (Cl) has 7 valence electrons. The total valence electrons for ${\mathrm{SiCl}}_4$ are 4 + (4*7) = 32.

Arrange the atoms

In ${\mathrm{SiCl}}_4$, there is one Silicon atom and four Chlorine atoms. Arrange the atoms with Silicon as the central atom surrounded by the Chlorine atoms: Si(Cl)₄.

Form the bonds and distribute remaining electrons

Form single covalent bonds between the Silicon and Chlorine atoms. We have used 8 electrons so far, and the remaining 24 electrons should be distributed to satisfy the octet rule for all atoms. The final Lewis structure for ${\mathrm{SiCl}}_4$ is: Cl | Cl - Si - Cl | Cl

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom. They play a crucial role in determining how an atom will interact or bond with other atoms. These electrons are found in the atom's highest energy level and are involved in forming bonds because they can be gained, lost, or shared.
Group numbers in the periodic table can help you identify the number of valence electrons for main-group elements. For example:

• Group 1 elements have 1 valence electron.
• Group 14 elements have 4 valence electrons.
• Group 15 elements, like nitrogen, have 5 valence electrons.
• Group 17 elements, such as chlorine, have 7 valence electrons.

Knowing the number of valence electrons is the first step in drawing a Lewis structure, as it helps determine how atoms will bond to achieve a stable electron configuration.

Octet Rule

The octet rule is a chemical rule of thumb reflecting the tendency of atoms to prefer having eight electrons in their valence shell. This arrangement is associated with the stability similar to noble gases, which naturally have complete outer electron shells.
When drawing Lewis structures, adhering to the octet rule is imperative, although there are some exceptions such as hydrogen and helium which aim to reach a duet (2 electrons) instead. Here’s how it generally works:

• Atoms will share, gain, or lose electrons in pursuit of a full valence shell.
• Covalent bonds often contribute shared electrons to help fill these shells.
• In some molecules, atoms might have less or more than 8 electrons due to resonance or extended octets, but typically, achieving an octet helps gauge electron allocation around central atoms.

Incorporating the octet rule helps to predict how molecules form and maintain stability through electron sharing.

Covalent Bonds

Covalent bonds are the connections formed when two atoms share pairs of electrons. This sharing allows the atoms to fill their respective valence shells, helping them achieve the stable electronic configuration prescribed by the octet rule.
Covalent bonds are distinct in that:

• They typically occur between non-metal atoms, which have similar electronegativities.
• Each covalent bond is represented by a line in a Lewis structure, symbolizing one pair of shared electrons.
• In molecules like $\text{Double subscripts: use braces to clarify}$, all involved atoms engage in single covalent bonds, sharing two electrons for each connection.

Covalent bonds can be single, double, or triple depending on the number of shared electron pairs. Thus, understanding covalent bonding is essential for grasping how molecules stabilize and interact.

Molecular Geometry

Molecular geometry is the three-dimensional arrangement of atoms within a molecule. This geometry is dictated by the number and type of bonds and lone pairs surrounding the central atom in a compound.
For example, to determine the molecular geometry:

• First identify the central atom, usually the least electronegative.
• Count both bonding and lone pairs of electrons around the central atom.
• Apply VSEPR theory (Valence Shell Electron Pair Repulsion), which states electron pairs will repel each other, arranging as far apart as possible.

As a result, different molecular shapes arise such as linear, bent, trigonal planar, tetrahedral, etc. For instance, $SiC{l}_4$ adopts a tetrahedral geometry due to four single bonds to chlorine atoms, all repelling each other equally around the silicon.