Lewis structure and molecular geometry for BeF2, NOF , SbCl3 and SiHBr3
The Correct Answer and Explanation is:
1. BeF₂ (Beryllium Difluoride)
Lewis Structure:
Be is the central atom, bonded to two F atoms.
Each F has 3 lone pairs (6 electrons), and Be forms two single bonds.
Be has no lone pairs.
Molecular Geometry: Linear
Bond Angle: 180°
2. NOF (Nitrosyl Fluoride)
Lewis Structure:
N is central, double bonded to O and single bonded to F.
O has 2 lone pairs, F has 3 lone pairs, and N has 1 lone pair.
Molecular Geometry: Bent (angular)
Bond Angle: Slightly less than 120°
3. SbCl₃ (Antimony Trichloride)
Lewis Structure:
Sb is the central atom with 3 single bonds to Cl atoms and 1 lone pair.
Each Cl has 3 lone pairs.
Molecular Geometry: Trigonal Pyramidal
Bond Angle: Approximately 101–102°
4. SiHBr₃ (Silyl Tribromide)
Lewis Structure:
Si is central, bonded to 1 H and 3 Br atoms with single bonds.
Each Br has 3 lone pairs. Si has no lone pairs.
Molecular Geometry: Tetrahedral
Bond Angle: Approximately 109.5°
Explanation
Lewis structures show how valence electrons are arranged around atoms in a molecule. The central atom is usually the least electronegative, and single or multiple bonds form to satisfy the octet rule, except for elements like Be, which often have fewer than 8 electrons.
In BeF₂, beryllium is an exception to the octet rule. It forms only two bonds with fluorine atoms, using four electrons total. Each fluorine completes its octet with three lone pairs. With no lone pairs on beryllium and two regions of electron density, the molecule adopts a linear geometry with a bond angle of 180°.
NOF is more complex. Nitrogen is bonded to both oxygen and fluorine. A double bond forms with oxygen and a single bond with fluorine. Nitrogen also holds a lone pair. The three electron regions around nitrogen form a trigonal planar electron arrangement, but the lone pair causes a bent shape due to repulsion, lowering the bond angle below 120°.
SbCl₃ features antimony with three bonds and one lone pair. This gives it four regions of electron density, resulting in a trigonal pyramidal shape. The lone pair pushes down on the bonding pairs, slightly compressing the bond angles below the ideal tetrahedral angle.
SiHBr₃ has four bonded atoms and no lone pairs on silicon. This gives a tetrahedral shape with ideal bond angles of 109.5°. The difference in size and electronegativity between H and Br does not significantly alter the geometry.
In summary, Lewis structures help visualize bonding and lone pairs, while molecular geometry is determined by the number of regions of electron density around the central atom.
