Drawing the Lewis structure of hydrogen bromide comes down to counting valence electrons and placing a single bond between the hydrogen and bromine atoms. Because hydrogen brings one valence electron and bromine brings seven, the combined molecule contains eight valence electrons, or four electron pairs.
Hydrogen bromide ($HBr$) forms a simple diatomic molecule featuring a single covalent bond. Since bromine has a much higher electronegativity than hydrogen, the shared electron pair spends more time near the bromine atom. This unequal pull creates a permanent dipole moment, turning the molecule into a classic example of a polar covalent bond.
Calculating Valence Electrons in Hydrogen Bromide
Before sketching any chemical bond, you need to tally the total available valence electrons from each constituent atom.
Hydrogen sits in Group 1 of the periodic table, contributing a single valence electron from its $1s$ orbital. Bromine belongs to Group 17, the halogens, and brings seven valence electrons from its outer shell ($4s^2 4p^5$).
Adding these values together gives a total of eight valence electrons for the $HBr$ molecule. Since a stable chemical structure requires satisfying the octet rule for main-group elements (and a duet rule for hydrogen), these eight electrons organize into one shared bonding pair and three unshared lone pairs residing entirely on the bromine atom.
Step-by-Step Drawing Process
Constructing the Lewis structure requires a methodical approach to ensure electron counts and octets check out correctly.
- Place the atoms: Write the symbol for hydrogen ($H$) next to the symbol for bromine ($Br$). In diatomic molecules, the atoms simply sit adjacent to each other.
- Form a single bond: Draw a single line between $H$ and $Br$ to represent a covalent bond. This line accounts for 2 of your 8 available valence electrons.
- Distribute remaining electrons: You have 6 electrons left. Place all three remaining electron pairs as lone pairs around the bromine atom. Bromine now holds a complete octet of eight electrons in its valence shell.
- Verify hydrogen's duet: Hydrogen shares the 2 electrons in the single bond, filling its single $1s$ orbital to reach a stable duet configuration.
To check your work, review official chemistry references like LibreTexts Chemistry or the Royal Society of Chemistry to see standard molecular representations.
Polarity and Molecular Geometry
The geometry of hydrogen bromide is strictly linear because any diatomic molecule only connects two points in space. However, its electronic distribution is far from symmetrical.
Bromine's electronegativity measures roughly 2.96 on the Pauling scale, while hydrogen sits lower at 2.20. That difference of 0.76 triggers a significant shift in electron density. The bonding pair pulls closer to bromine, generating a partial negative charge ($\delta^-$) on the bromine side and a partial positive charge ($\delta^+$) on the hydrogen side.
This permanent dipole moment makes liquid and gaseous $HBr$ exceptionally polar, influencing how molecules interact with one another through dipole-dipole forces.
Physical and Chemical Properties
The electronic structure of $HBr$ dictates its macroscopic behavior in laboratory and industrial settings.
At standard temperature and pressure, hydrogen bromide exists as a colorless gas with a pungent, irritating odor. It condenses into a liquid at $-66.9^\circ C$ and freezes at $-86.9^\circ C$. When dissolved in water, it forms hydrobromic acid, which acts as a strong, highly corrosive acid capable of ionizing completely in aqueous solution.
Chemically, $HBr$ participates heavily in organic synthesis. Chemists use it to add across double bonds in alkenes via Markovnikov's rule, yielding alkyl bromides. It also serves as a potent reducing agent and a catalyst in industrial manufacturing processes.
Common Mistakes to Avoid
Students often stumble on a few frequent pitfalls when drawing or analyzing this specific molecule.
- Giving hydrogen an octet: Hydrogen can only hold two electrons in its valence shell. Never add lone pairs to hydrogen or try to surround it with eight electrons.
- Misplacing lone pairs: All three lone pairs belong exclusively on the bromine atom, not on hydrogen or floating in the middle of the bond.
- Assuming nonpolar behavior: Because both atoms are nonmetals, beginners sometimes assume the bond is nonpolar. Always check electronegativity differences rather than relying solely on atom types.
Frequently Asked Questions
Why doesn't bromine share more electrons with hydrogen?
Hydrogen only has a $1s$ orbital, which accommodates a maximum of two electrons. Once hydrogen forms a single covalent bond, its valence shell is completely full, meaning it cannot accept additional electron pairs from bromine.
Is hydrogen bromide an ionic or covalent compound?
It is a polar covalent compound. Although the electronegativity difference is large enough to create a strong dipole, it isn't high enough to strip an electron completely away from hydrogen to form distinct ions in the gas phase.
How does HBr differ from HCl in its Lewis structure?
The Lewis structures look identical on paper—both feature a single bond and three lone pairs on the halogen. The primary difference lies in the size of the halogen atom and the bond length, as bromine is larger than chlorine, resulting in a longer and weaker bond in $HBr$.
What is the formal charge of the atoms in HBr?
Both hydrogen and bromine carry a formal charge of zero in a neutral $HBr$ molecule. You can calculate this by taking the valence electrons of the free atom, subtracting the assigned lone pair electrons, and subtracting half of the shared bonding electrons.
