Difficulty: Easy
Correct Answer: SiCl4 is a nonpolar molecule because the polar Si–Cl bonds cancel in a tetrahedral shape
Explanation:
Introduction / Context:
Understanding whether a molecule is polar or nonpolar is a key part of chemical bonding and molecular structure. Polarity affects boiling point, solubility, and many physical properties. Silicon tetrachloride, SiCl4, is a common example used in textbooks to show how a molecule can contain polar bonds but still be nonpolar overall. This question asks you to identify whether SiCl4 is polar or nonpolar and to connect that to its geometry and bond dipoles.
Given Data / Assumptions:
Concept / Approach:
The approach is to combine two ideas. First, each Si–Cl bond is polar because chlorine is more electronegative than silicon, so electrons are drawn toward the chlorine atom. Second, the overall molecular geometry determines whether these individual bond dipoles cancel or reinforce each other. According to valence shell electron pair repulsion ideas, a molecule with four bonding pairs and no lone pairs on the central atom adopts a tetrahedral geometry. In a perfect tetrahedron with identical outer atoms, the bond dipoles are symmetrically arranged and cancel out, giving a nonpolar molecule even though each bond is polar.
Step-by-Step Solution:
Step 1: Determine the electron pair arrangement around silicon. Silicon in SiCl4 forms four single covalent bonds and has no lone pairs on the central atom.
Step 2: Use valence shell electron pair repulsion rules: four bonding pairs and no lone pairs lead to a tetrahedral molecular geometry.
Step 3: Recognise that each Si–Cl bond is polar because chlorine attracts bonding electrons more strongly than silicon, giving a partial negative charge near each chlorine.
Step 4: In a regular tetrahedral arrangement with four identical polar bonds, the bond dipole vectors are oriented symmetrically and cancel each other, resulting in no net dipole moment for the whole molecule.
Step 5: Therefore, SiCl4 is classified as a nonpolar molecule, even though the individual bonds are polar.
Verification / Alternative check:
Another way to verify this conclusion is to recall similar molecules such as carbon tetrachloride, CCl4, which is also nonpolar for the same geometric reason. Experimental measurements of dipole moment for SiCl4 show a value close to zero, supporting the nonpolar classification. In addition, the solubility behaviour is typical of nonpolar molecules; SiCl4 is more soluble in nonpolar organic solvents than in water, which is consistent with an overall nonpolar character.
Why Other Options Are Wrong:
Option A is wrong because although Si–Cl bonds are polar, the shape is not asymmetric; it is symmetric tetrahedral, so the dipoles cancel. Option C is incorrect since SiCl4 is not an ionic salt; it is a covalent molecular compound. Option D claims the molecule is linear, which contradicts the known tetrahedral shape. Option E is wrong because it states that the bonds are nonpolar; the bonds themselves are polar, but the net molecule is nonpolar due to symmetry. Only option B correctly explains that SiCl4 is nonpolar because of cancellation of bond dipoles in a tetrahedral geometry.
Common Pitfalls:
A frequent mistake is to assume that if all bonds in a molecule are polar, then the molecule must automatically be polar. Another pitfall is to ignore the molecular shape and focus only on electronegativity differences. Some learners also misremember the geometry and think SiCl4 is planar or linear rather than tetrahedral. Always determine the shape first, then combine bond polarity with symmetry to decide overall polarity. For symmetrical molecules with identical outer atoms, bond dipoles often cancel.
Final Answer:
The correct description is that SiCl4 is a nonpolar molecule because the polar Si–Cl bonds cancel in a tetrahedral shape and the net dipole moment is essentially zero.
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