Question

Draw Lewis structures for the following diatomic molecules and ions: (a) \(\mathrm{CO} ;\) (b) \(\mathrm{O}_{2} ;\) (c) \(\mathrm{ClO}^{-} ;\) (d) \(\mathrm{CN}^{-}.\)

Short answer

Question: Draw the Lewis structures for the following diatomic molecules and ions: (a) CO (b) O2 (c) ClO\(^{-}\) (d) CN\(^{-}\) Answer: The Lewis structures for the given diatomic molecules and ions are as follows: (a) CO: :C::\(\equiv\)O: (b) O2: :O::\(\equiv\)O:: (c) ClO\(^{-}\): :Cl:─:O:(\(^{-}\)) (d) CN\(^{-}\): :─C\(\equiv\)N(\(^{-}\))

Step-by-step solution

Determine the total number of valence electrons

For each molecule/ion, we first find the total number of valence electrons from the periodic table. (a) CO: C has 4 valence electrons and O has 6 valence electrons, so the total number of valence electrons is 10. (b) O2: Each O atom has 6 valence electrons, thus, the total number of valence electrons is 12. (c) ClO\(^{-}\): Cl has 7 valence electrons, O has 6 valence electrons, and the extra electron (due to the negative charge) adds 1 more electron: the total number of valence electrons is 14. (d) CN\(^{-}\): C has 4 valence electrons, N has 5 valence electrons, and the extra electron (due to the negative charge) adds 1 more electron: the total number of valence electrons is 10.

Create a skeleton structure for each molecule/ion

Now we will arrange the atoms to form a skeleton structure by attaching atoms together using single bonds. (a) CO: C and O are bonded together. (b) O2: Two O atoms are bonded together. (c) ClO\(^{-}\): Cl and O are bonded together. (d) CN\(^{-}\): C and N are bonded together.

Distribute remaining valence electrons

Now we will place the remaining valence electrons on the atoms to satisfy the octet rule (with each atom, except hydrogen, having 8 electrons in its valence shell). (a) CO: There are 6 remaining electrons after the C-O bond formation. Place 3 lone pairs on O. (b) O2: There are 8 remaining electrons after the O-O bond formation. Place 2 lone pairs on each of the O atoms. (c) ClO\(^{-}\): There are 10 remaining electrons after the Cl-O bond formation. Place 3 lone pairs on O and 2 lone pairs on Cl. (d) CN\(^{-}\): There are 7 remaining electrons after the C-N bond formation. Place 2 lone pairs on N and 1 lone pair on C.

Modify the structures to achieve octet rule if needed

After arranging the electrons, we check if the octet rule is satisfied for each atom. If not, modify the Lewis structure with multiple bonds (double or triple). (a) CO: C and O have 7 and 9 electrons respectively, instead of octet. Change a lone pair from O into a double bond with C. (b) O2: Each O atom has only 7 electrons instead of octet. Change a lone pair from both the O atoms into a double bond between the two O atoms. (c) ClO\(^{-}\): The octet rule is satisfied for all the atoms. (d) CN\(^{-}\): C and N have only 7 and 9 electrons respectively, instead of octet. Change the lone pair from N into a triple bond with C.

Draw the final Lewis structures

Now we can draw the final Lewis structures for each molecule/ion with the completed skeleton structure and the proper electron distribution. (a) CO: :C::\(\equiv\)O: (b) O2: :O::\(\equiv\)O:: (c) ClO\(^{-}\): :Cl:─:O:(\(^{-}\)) (d) CN\(^{-}\): :─C\(\equiv\)N(\(^{-}\))

Key concepts

Valence Electrons

Valence electrons play a crucial role in determining how atoms interact and bond with each other. They are the electrons present in the outermost shell of an atom. These electrons are responsible for the chemical properties of the elements.

When atoms form bonds, their valence electrons can either be shared with or transferred to other atoms, thereby achieving more stable electron configurations. To find the number of valence electrons, you look at the periodic table where you can see the group numbers:

• Elements in group 1 have 1 valence electron.
• Elements in group 2 have 2 valence electrons.
• Groups 13-18 have 3-8 valence electrons, respectively.

In a molecule, knowing the total number of valence electrons helps us distribute these electrons properly and draft the initial structure, as seen with

• CO having 10 valence electrons.
• O sub2 having 12 valence electrons.
• ClO sup{-} having 14 valence electrons.
• CN sup{-} having 10 valence electrons.

This is why understanding and identifying valence electrons is our first and pivotal step towards drawing Lewis structures.

Octet Rule

The octet rule is a fundamental concept in chemistry that describes how atoms tend to combine in such a way that they each have eight electrons in their valence shell, achieving a noble gas configuration. This rule is key to understanding how molecules are structured and how chemical bonds are formed.

The octet rule helps in predicting how atoms will share or transfer electrons to achieve the stable electron configuration. Atoms, other than hydrogen, which only requires 2 electrons, strive to have 8 valence electrons.

• In CO, the octet rule is initially not satisfied after the first step. Carbon and oxygen need their electron count adjusted so they have a full octet.
• For O sub2, each oxygen follows the octet rule by forming a double bond.
• In ClO sup{-}, both chlorine and oxygen satisfy their octets naturally with lone pairs and the existing single bond.
• CN sup{-} requires adjustments using a triple bond to satisfy the octet rule for carbon and nitrogen.

Thus, the octet rule explains why elements form specific bonds, adjusting accordingly to satisfy this criterion during Lewis structure construction.

Multiple Bonds

In some molecules, single bonds can't satisfy the octet rule for all involved atoms. In these cases, multiple bonds, which include double and triple bonds, are used to share more than one pair of electrons between atoms.

Multiple bonds are crucial for understanding molecules in which atoms need more electrons to achieve their octets. Here's what they typically involve:

• Double bonds: These occur when two pairs of electrons (four electrons in total) are shared between two atoms. An example is the oxygen molecule (O sub2), where a double bond allows each oxygen atom to achieve an octet.
• Triple bonds: Here, three pairs of electrons (six electrons in total) are shared. A classic case is the cyanide ion (CN sup{-}), where the triple bond is required for both carbon and nitrogen to attain octet.

Multiple bonds not only fulfill the octet rule but also influence the physical characteristics of molecules such as bond length and bond strength. Knowing when and how to form these bonds is an essential skill in drawing accurate Lewis structures.