Apparent contradiction in the type of deviation of an ethanol-water mixture
As suggested in the comments, there is an entropic contribution to the excess free energy which should not be overlooked, as it is the excess Gibbs free energy that determines whether the deviation from Raoult's law is positive or negative. All of the statements made by the OP are correct, except that negative $\Delta H$ (exothermicity) and $\Delta V$ (volume contraction) do not necessarily imply a reduced volatility (negative deviation from Raoult's law), that is they are not sufficient conditions for a negative deviation, a negative excess Gibbs free energy instead being the sufficient condition. Again, as alluded to in the comments, a simple analysis based entirely on $\Delta H$ is not appropriate here, largely because of the hydrogen bonding (associative) nature of the liquids. In fact, the negative mixing volume and enthalpy might partly explain the negative excess entropy (and vice-verse: they are correlated). The excess entropy of mixing of water and ethanol is highly negative over the entire range of composition at room temperature, leading to a net positive excess Gibbs free energy. 


  Does the mixture actually exhibit positive deviation? Which interactions are stronger - unlike or like?


Therefore the mixture does exhibit a positive deviation from Raoult's law, even as interaction enthalpies between unlike molecules are stronger.