Carbon tetrachloride (CCl4) contains polar C–Cl bonds, yet the molecule has no overall dipole moment. What is the main reason for its net dipole moment being zero?

Introduction / Context:
Molecular polarity depends not only on differences in electronegativity between atoms but also on the three dimensional arrangement of bonds. Carbon tetrachloride, CCl4, is a classic example used to illustrate how a molecule with polar bonds can still have zero overall dipole moment if the geometry is highly symmetric. This question asks for the correct explanation of why CCl4 has no net dipole moment even though the C–Cl bonds are individually polar.

Given Data / Assumptions:

• Carbon tetrachloride has the formula CCl4.
• The electronegativity of chlorine is greater than that of carbon, so each C–Cl bond is polar.
• The molecular geometry of CCl4 is predicted by valence shell electron pair repulsion theory.
• Dipole moment is a vector quantity that depends on both magnitude and direction of bond dipoles.

Concept / Approach:
The central carbon atom in CCl4 has four bonding pairs and no lone pairs. VSEPR theory predicts a tetrahedral geometry with bond angles of about 109.5 degrees. Each C–Cl bond dipole points from carbon towards chlorine, but because the four chlorine atoms are arranged symmetrically around the carbon, the vector sum of these four bond dipoles is zero. This complete cancellation leads to a nonpolar molecule with zero net dipole moment, despite each bond being polar. The explanation does not depend on equal atomic sizes or identical electron affinities, but on the symmetric tetrahedral geometry.

Step-by-Step Solution:
Step 1: Recognise that in CCl4, carbon is surrounded by four chlorine atoms and forms four single covalent bonds. Step 2: Use VSEPR theory to predict that with four electron pair regions and no lone pairs, the geometry is tetrahedral. Step 3: Recall that each C–Cl bond is polar because chlorine is more electronegative than carbon. Step 4: Visualise the four bond dipoles pointing symmetrically outwards from the central carbon towards the corners of a tetrahedron. Step 5: Note that due to this symmetry, the vector sum of the four equal dipoles is zero, so the molecule has no net dipole moment.

Verification / Alternative check:
Compare CCl4 with CHCl3 (chloroform). In CHCl3, the geometry is still approximately tetrahedral, but three positions are occupied by chlorine and one by hydrogen. The bond dipoles do not cancel perfectly because hydrogen has different electronegativity and polarisation compared to chlorine. As a result, CHCl3 has a net dipole moment and is polar, while CCl4, with four identical C–Cl bonds arranged symmetrically, has zero net dipole moment. This comparison supports the explanation based on tetrahedral symmetry.

Why Other Options Are Wrong:
The sizes of carbon and chlorine atoms are not equal, and size alone does not determine dipole moment. The molecule is not planar; it is three dimensional tetrahedral. Carbon and chlorine do not have identical electron affinities or electronegativities, which is why the bonds are polar in the first place. It is incorrect to say that there are no polar bonds in CCl4; each C–Cl bond is indeed polar. Thus, those statements do not explain the observed zero dipole moment.

Common Pitfalls:
Students sometimes think that any molecule with polar bonds must be polar and forget to consider geometry and vector addition of dipoles. Another mistake is to assume that tetrahedral molecules are always polar, ignoring cases like CCl4 and CF4. To avoid these errors, always draw or visualise the three dimensional shape, mark the bond dipoles, and then consider whether they cancel due to symmetry. Recognising highly symmetric geometries (linear, trigonal planar, tetrahedral, octahedral) often helps explain zero dipole moments.

Final Answer:
Carbon tetrachloride has no net dipole moment because its regular tetrahedral structure makes the four identical C–Cl bond dipoles cancel out.