I'm a bad person.

Having taken Gen Chem in 2008 (which in my world is like a BILLION years ago), I'm kind of having internal panic attacks about Biochemistry. I forgot all about the Henderson Hasselbalch equation and how to do titrations and all that shit so I've spent the last week and a half relearning that. In case anyone else is in my boat, here's how to bring back old memories.

Water dissociates into H+ and OH- ions.

At 25' C, the concentration [H+] ions (what we worry about) is 10^-7 M

Apparently, pH is the negative log of the [H+] and using log rules, you can calculate the pH.

Here's some more crap.

Talking about Equilibrium Constants this time (ratio of concentrations of crap in solution when the equilibrium gets to tell you that the rate of the reaction going in one direction = rate of reaction going in reverse direction)

Kw (equilibrium constant) calculated by [Products]/[Reactants], with H20 being 1, is 10^-14

Then we use some logs again to get the pkw, log calculation to the side for anyone who isn't log-savvy, and we end up with a pkw = 14

Some more very basic crap. Getting the pH of a 1M solution of a strong acid.

Overly complicated way to solve a very quick issue...getting the pH of a 1.5M solution of HF, given the Ka value.

Do the Ka = Products/Reactants, the icebox to the lower left afterwards.

You start out with 1.5M HF, and 0 of H+, F-

You use up "x" of the HF, meaning you gain "x" of H+ and F-

Final concentrations are 1.5-x of HF and x of H+ and x of F-

Plug that all into the Ka = Products/Reactants formula, neglect x for the denominator, and do some really simple algebraic crap. pH you get is about 1.49.

Whole purpose of that is to just show relationship between pka and pkb

It's unfortunate that this is so blurry since I definitely remember in gen chem that my life felt like it depended on knowing how to go from the Ka to the pH using the Products/Reactants crap.

This is REALLY important. 

Once I get back on my feet, there will be a LOT more where this came from.