The behavior of real gases is often analyzed to see how closely they follow the ideal gas laws, particularly as these laws are approached under varying conditions of temperature and pressure. The correct answer in context to the temperature at which a real gas behaves most like an ideal gas over a wide range of pressures is not the critical temperature, inversion temperature, or reduced temperature. Instead, it is the Boyle temperature.

To elaborate, let's look at each option:

Critical temperature is the temperature above which a gas cannot be liquefied by pressure alone. At this temperature, the properties of the gas change significantly, which does not necessarily imply ideal gas behavior.
Boyle temperature is the specific temperature at which a real gas adheres to Boyle's Law (PV = constant) over a wide range of pressures. At this temperature, the attractive and repulsive forces between molecules balance out in such a way that the gas's P-V behavior (pressure-volume) mimics that of an ideal gas.
Inversion temperature is related to the Joule-Thomson effect, which describes how the temperature of a real gas changes when it expands or compresses without the exchange of heat with its surroundings. This is different from behaving like an ideal gas over a range of pressures.
Reduced temperature is a term used in corresponding states theory to correlate the properties of substance under various conditions. It is dimensionless and given by $$ T_r = \frac{T}{T_c} $$ where $ T $ is the absolute temperature and $ T_c $ is the critical temperature.

Since the question specifically asks about the temperature at which the behavior of a real gas is similar to that of an ideal gas across a broad range of pressures, the answer is Option B: Boyle temperature. This is the temperature where deviations from ideal gas law due to intermolecular forces are minimized in the behavior of real gases, thus aligning more closely to the predictions by the ideal gas law (PV=nRT).