No. You'll get the same temperature rise from an ideal frictionless gas.
(This was one of the exercises in fluid mechanics class in college.)
The bit about the air cannot get out of the way fast enough comes from its inertia, not friction.
You'll find the same temperature rise in high speed aircraft like the SST. The max temperature rise is at the stagnation point where the air velocity is zero.
(Compress any gas and it warms up, this is how refrigerators work.)
What makes the air too slow to get out of the way? I would have naively thought friction, and that if everything was frictionless the air would seamlessly move to replace the vacuum behind the aircraft, and thus no compression would occur.
To put it bluntly, air molecules aren't psychic. The only way they know what happens to neighboring air molecules is by running into them. At every-day velocities, this happens fast enough that we can treat it as though it is instantaneous. That the "vacuum" is a thing that sucks air in.
At high velocities, this intuition breaks down. The communication time between molecules is limited to the speed of sound, the speed at which air molecules travel and bump into each other. If you are traveling faster than the speed of sound, (or several times faster, as is the case for re-entry), the air molecules don't have any information that an object is coming to them, because there hasn't been enough time for molecules hitting the object to then hit them.
In an ideal frictionless gas (zero viscosity), the flow after the object has passed through will be identical to the flow before the object, so in that sense the air moves "seamlessly" around the object. But during this process, the pressure of the gas still changes. You can think of this as just newton's second law: in order to change the velocity of an air packet, you must apply a force to it, and pressure is just force per area. So any time a gas flows in anything other than a straight line, you know that there was a pressure gradient involved and the pressure was not constant.
It's a common example in fluid mechanics textbooks to derive the flow of a zero-viscosity gas around a cylinder, you can get an exact solution e.g. here: http://brennen.caltech.edu/fluidbook/basicfluiddynamics/pote... . Note that in this example there is no drag, but there are two zones of high pressure in front of and behind the cylinder. Those zones would have higher-than-ambient temperature.
Heating via compression is not the same as heating by friction.
When you squeeze a gas into a smaller volume it heats up because the heat energy of the gas which was spread out over the larger volume is now packed into a smaller one.
Boyle's Law describes a behavior without positing the mechanism.
What does it mean for a gas to heat up? Is the quantity of heat energy increasing, or the temperature?
A thermometer generally measures the temperature of molecules adjacent to the thermometer, which is a function of the amplitude of those molecules' vibration. More molecules within a given space increases the heat capacity, and the total heat energy represented by a temperature of the total quantity of molecules, but the temperature is a function of the average heat energy per individual molecule.
The temperature increase described by Boyle's law is a consequence of the work required to compress a gas—somewhere there has to be a compressor imparting kinetic energy to individual atoms.
Think of it this way, you have two containers of heat energy represented by 1L of an ideal gas at 1 atmosphere and 20°C. If you magically transported the contents of both containers into a single 1L container, and they continued to express a temperature of 20°C, you would have twice as much stuff, and contain twice as much heat energy in 1L. If on the other hand you had twice as much stuff and a higher temperature, then you would have increased the total amount of energy present in the universe.
The temperature comes from the extra energy put into the system to compress the gas rather than expressly from the density of the heat energy.
