Draw a reasonable Lewis structure for each of the following formulas. a. H2S b. CHBr3 c. NF3 d. Br2O e. H2CO3 f. H2S2 g. HOCl h. BBr3 i. CH3CH2CHCH2
Here are the reasonable Lewis structures for each given molecular formula, followed by an explanation.
a. H₂S (Hydrogen sulfide)
- Sulfur is the central atom with two single bonds to hydrogen atoms.
- Sulfur has two lone pairs.
Structure:
H–S–H (with two lone pairs on S)
b. CHBr₃ (Bromoform)
- Carbon is central, bonded to one hydrogen and three bromine atoms.
- Carbon has four single bonds; bromines have three lone pairs each.
Structure:
H–C–Br₃ (tetrahedral geometry)
c. NF₃ (Nitrogen trifluoride)
- Nitrogen is central, bonded to three fluorine atoms with one lone pair.
- Each fluorine has three lone pairs.
Structure:
F–N–F with a lone pair on N and three lone pairs on each F
d. Br₂O (Dibromine monoxide)
- Oxygen is central, bonded to two bromine atoms with two lone pairs.
- Each bromine has three lone pairs.
Structure:
Br–O–Br
e. H₂CO₃ (Carbonic acid)
- Carbon is central with one double bond to oxygen and two single bonds to OH groups.
- Each O has two lone pairs.
Structure:
HO–C(=O)–OH
f. H₂S₂ (Disulfane)
- Two sulfur atoms bonded together.
- Each sulfur also bonds to a hydrogen and has two lone pairs.
Structure:
H–S–S–H
g. HOCl (Hypochlorous acid)
- Oxygen is central, bonded to hydrogen and chlorine.
- Oxygen has two lone pairs; chlorine has three.
Structure:
H–O–Cl
h. BBr₃ (Boron tribromide)
- Boron is central, bonded to three bromine atoms.
- Boron has no lone pairs; each Br has three.
Structure:
Br–B–Br (trigonal planar)
i. CH₃CH₂CHCH₂ (1-butene)
- A four-carbon chain with a double bond between the third and fourth carbons.
- Hydrogen atoms fill in to satisfy valence.
Structure:
CH₃–CH₂–CH=CH₂
Explanation:
Lewis structures show how valence electrons are arranged among atoms in a molecule. They help predict geometry, reactivity, and polarity. In each of the molecules above, the central atom is usually the least electronegative (except hydrogen, which can never be central).
For example, in H₂S, sulfur forms two single bonds with hydrogen and keeps two lone pairs, following the octet rule. In CHBr₃, carbon has four bonds forming a tetrahedral structure, with bromines completing octets with lone pairs.
NF₃ resembles ammonia (NH₃) but with fluorines. Nitrogen forms three bonds and retains one lone pair. Fluorine, being highly electronegative, doesn’t form more than one bond and completes its octet with lone pairs.
In Br₂O, oxygen is more electronegative and central, forming two single bonds with bromine. H₂CO₃, an acid, has a central carbon double-bonded to one oxygen and single-bonded to two hydroxyl groups.
H₂S₂ is similar in bonding to hydrogen peroxide, with a sulfur–sulfur single bond and each sulfur also bonded to a hydrogen. Lone pairs on sulfur affect molecular shape and reactivity.
HOCl is a small molecule where oxygen bonds to hydrogen and chlorine. BBr₃ is an electron-deficient compound where boron forms three bonds without achieving a full octet, making it a strong Lewis acid.
Lastly, 1-butene (CH₃CH₂CHCH₂) contains a carbon–carbon double bond, creating a region of high electron density and affecting the geometry with sp² hybridization around the double bond.
These Lewis structures follow valence rules and help in understanding chemical behavior.
