Very basic question here, but I'm confused why Delta-H seems to be the reverse of what I would expect for bond enthalpy. For example, here's a problem in my textbook where the goal is to find delta-H:
$$\ce{H2(g) + Br2(l) -> 2HBr(g)}$$
Using the table I have, I see that an $\ce{H-H}$ bond has an average enthalpy of 436, $\ce{Br-Br}$ has 193, and $\ce{H-Br}$ has 366. 436+193 = 629, and 2*366 = 732, so the reactants have a total of 629, compared to the products, which have 732.
My intuition (and Google, as far as I can tell) would tell me that, because the H value increased when the reaction occured, the answer would be positive 103. But my textbook says the answer is -103. This makes no sense to me - what am I missing here? Why do all the Delta-H (AKA Change in H) values seem to be negative what you'd expect?