I have no quibbles with Buck’s answer, which is correct (and I upvoted it), but consider another plausible scenario. Suppose the OP’s stated multiple choice question is simply an exam question from an exam given during the first semester of a typical two semester general chemistry course. Perhaps the course sequence is for science majors but not chemistry majors, per se. This will be assumed below.
Then the putative exam question would be quite reasonable because students would have learned about exothermic and endothermic reactions and would likely have seen some of these common reactions in lecture demonstrations. Four specific examples from my own experience, teaching such a course multiple times, include 1) burning magnesium in air (lots of heat and blinding white and UV light released), 2) throwing calcium metal into not much water in a big beaker (lots of heat produced in the roiling and boiling reaction mixture, the beaker gets too hot to touch, hot water vapor spews up and out of the beaker), 3) the super-saturated sodium acetate ‘hot ice’ demo (which is used in some consumer ‘cold packs’) and 4) dissolution of ammonium nitrate in water in a beaker (the solution gets quite cold and students are told that this is another way that ‘cold packs’ have been made).
Students are told that the first three reactions are exothermic, as are lots of other reactions familiar from everyday life, e.g., burning wood in a fireplace or combustion of gasoline in an automobile engine, are also exothermic and the reactions are performed specifically because heat is desired. The dissolution of ammonium nitrate is described as endothermic and the reason for this is deferred to the next semester because students have not been introduced to entropy.
And there is the issue: students are introduced to thermodynamic concepts in much the same fashion as hospital patients are introduced to intravenous drips: slowly, for the common good of them and the teacher. During the first semester, students encounter the concepts of system, surroundings, energy and energy changes, enthalpy and enthalpy changes, Hess’s law and standard enthalpies of reaction. But it is not until the second semester that students are introduced to entropy, free energy, changes in them and equilibrium.
Arguments abound about best practice in teaching the beginning material in general chemistry and whether or not it should even be taught at all: some favor just starting with organic chemistry. People go apoplectic about some of the simplifications and admitted over-simplifications, e.g., VSEPR with d orbitals, that are perpetrated in the first year course. But it is not an upper level undergraduate course or a graduate course and students would be poorly served if everything had to be meticulously and profoundly correct. This does not mean that teaching incorrect material is justified. It simply means that there is a place for zero order approximations in the pedagogy. Of course, rigorous quantum mechanics is correct in so far as we know. But, in light of the great Richard Feynman’s experience teaching physics to insanely bright freshman CalTech students, it would take amazing courage to attempt such an endeavor in first year chemistry, especially for non-majors.
Now, with regard to the OP’s question, the answer is simply “D”: heat is released in exothermic reactions. For this question and its implicit course context, it does not matter where the heat goes: if the beaker gets hot (as in the calcium and water demo), or you can get burned by touching it, then it was exothermic. So answer “C” is wrong. The first two answers are obviously incorrect: burning magnesium (or wood or paper or ...) does not involve bubbles and light is not released when calcium metal is dropped into water.