Key idea: order of reaction vs molecularity

• Order of reaction is determined experimentally from the rate law.
• Molecularity is the number of species colliding in an elementary (single-step) reaction.
• For a single-step reaction, order and molecularity are related.
• For a multistep reaction, the overall order depends on the mechanism and need not directly reflect stoichiometry.

Analyze each option

Option A: "A first order reaction is always a single step reaction"

This is false.

A first order reaction has rate law: $r = k[A]$ But such a rate law can arise from either:

• an elementary unimolecular step, or
• a multistep mechanism whose rate-determining step leads to first-order dependence.

So, first-order reaction is not always single step.

Option B: "A zero order reaction is a multistep reaction"

This is generally true in the context of elementary reactions.

For an elementary single-step reaction, rate depends on concentration of reactants involved, so order cannot normally be zero. Zero-order kinetics arise in special situations such as:

• surface-catalyzed reactions,
• photochemical reactions,
• enzyme reactions at saturation, where the rate becomes independent of concentration due to the mechanism.

Hence zero-order reactions are associated with complex/multistep mechanisms, not a simple elementary step.

Therefore, B is correct.

Option C: "A zero order reaction is a single step reaction"

This is false.

As discussed, a true elementary single-step reaction cannot have zero order with respect to all reactants in the usual kinetic sense.

Option D: "A second order reaction is always a multistep reaction"

This is false.

A second-order reaction can be a single elementary bimolecular step, for example: $A + B \to \text{products}$ with rate law $r = k[A][B]$ or $2A \to \text{products}, \quad r = k[A]^2$

So second-order reactions are not always multistep.

Conclusion

The correct option is: $\boxed{\text{B}}$