Among the following molecules, which one contains the fewest covalent bonds?

Rationale

Chlorine molecule (Cl2) contains the fewest covalent bonds.

The chlorine molecule (Cl2) consists of two chlorine atoms connected by a single covalent bond, making it the molecule with the fewest covalent bonds among the given choices. In contrast, the other molecules listed contain multiple covalent bonds.

A (Water molecule (H2O)) YOUR ANSWER

The water molecule (H2O) consists of two hydrogen atoms covalently bonded to one oxygen atom, resulting in a total of two covalent bonds. This configuration allows water to exhibit its characteristic properties, such as polarity and hydrogen bonding.

B (Methane molecule (CH4)) YOUR ANSWER

The methane molecule (CH4) is formed by a carbon atom covalently bonded to four hydrogen atoms, resulting in a total of four covalent bonds. This tetrahedral arrangement maximizes the stability of the molecule and is fundamental to its chemical behavior.

C (Chlorine molecule (Cl2)) CORRECT

The chlorine molecule (Cl2) is composed of two chlorine atoms sharing a single covalent bond. This bond is the only connection between the two atoms, making Cl2 the molecule with the fewest covalent bonds among the options presented.

D (Ammonia molecule (NH3)) YOUR ANSWER

The ammonia molecule (NH3) consists of one nitrogen atom covalently bonded to three hydrogen atoms, resulting in a total of three covalent bonds. This configuration leads to the characteristic trigonal pyramidal shape of ammonia.

Conclusion

In summary, among the molecules listed, the chlorine molecule (Cl2) contains the fewest covalent bonds with only one bond between its two chlorine atoms. In contrast, water, methane, and ammonia contain two, four, and three covalent bonds, respectively. Understanding covalent bonding is essential in predicting molecular behavior and interactions in chemical reactions.

Among the following molecules, which one contains the fewest covalent bonds?

Chlorine molecule (Cl2) contains the fewest covalent bonds.

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