This is my first post since February, partially because I haven't recently read any riveting books to review, but mostly because I've been quite busy with chemistry and calculus and OH DEAR GOD IN HEAVEN, LET IT END!! IF I SEE ANOTHER TITRATION PROBLEM I SWEAR I'LL-
Sorry. That slips out sometimes. I do love what I study, I swear, but BOY, is it tedious.
Something that I really enjoy about my chemistry class is that I finally understand how acids and bases work. When you're in elementary and middle school, you're told that some substances are acids, and some substances are bases, and they react together. Acid-ness or base-ness is measured on this mysterious spectrum called the "pH scale". They turn litmus paper different colors, too.
But when I asked, "What exactly makes a molecule an acid or a base?", my teacher could not find an answer past "Well, acids turn litmus paper blue litmus paper red, and bases turn red litmus paper blue." It frustrated me for years. So if you have suffered the same frustration, I am here to enlighten you! Let me be your guide in the magical microscopic world of hydrogen-swapping!
There are three kinds of acids and bases. The most basic (not in the chemical sense) type is called Arrhenius acids and bases. An Arrhenius acid is a molecule with a single hydrogen atom dangling off the side, and when it is mixed in water, the water molecules, with their funky bendy shape, yank off that hydrogen. The water molecules that do this then form what's called a hydronium ion, or H3O+. (The "+" means it has a positive charge, like a magnet.) An Arrhenius base is similar; it has an OH- (a hydroxide, an oxygen and a hydrogen) dangling off of the side. This, too, is yanked off by water's wily ways. (The "-" means it has a negative charge, also like a magnet.) BUT! Things get interesting when you put them TOGETHER!
Let's take a typical strong acid, HCl, and a typical strong base, NaOH. But let's make this exciting; let's put some adventure into there. Let's say you are a mad scientist who put more energy into studying physics and engineering than chemistry. You ride in on your stealth helicopter to the local college and use your laser pistol to carve a hole in the window of the chemistry lab. You use your super-duper-lockpicking-invention to break into the chemical storage room, and you pick out two amber bottles that say "HYDROCHLORIC ACID" and "SODIUM HYDROXIDE". You've heard of those. They sound DANGEROUS! But could they be more dangerous TOGETHER? You mix some together in equal amounts in a beaker. Little do you know that a chemical jitterbug is going on on that beaker's dance floor:
HCl + NaOH ---> H2O + NaCl
You laugh manically when an idea hits you- drinking it could potentially kill you, but it could also give you superpowers, like Spiderman. Being a mad scientist, emphasis on mad, you have little impulse control and decide to take a swig. As soon as it fills your mouth, you recoil and spray it out, partially through your nose- it tastes like salt water! It IS salt water! You rush to rinse out your mouth and grab the nearest bottle labeled "water", but that tastes nasty, too; it's deionized water. You're not having the best night. Eventually it dawns on you that mixing these kinds of acids and bases generally makes water and whatever else the other atomic legos in the mix make when they are snapped together.
Next in line we have Bronstead-Lowry acids and bases! A Bronstead-Lowry base is what we call a proton-accepter. Remember that dangling hydrogen? Hydrogen is the first element on the periodic table- it has one proton in the middle, and one single electron whizzing around it. (Its lack of a second electron makes it highly unstable, which is why it likes to explode, but that's for another post.) When it's hanging out on the edge of a molecule, that single electron is doing double duty zipping around its nucleus AND zipping between the other electrons in the molecule to keep it attached. That electron is working SO hard to keep it in place that it can rarely ever stop by to visit, so that lonely hydrogen is dangling off the side, barely attached, essentially a naked proton. When you throw the molecule into water, the water molecules, or whatever base is in there, looks at that naked proton and goes "Oh! Poor thing, it must be freezing! Let's give it some of our electrons!", and before you know it, that proton has been spirited away into another molecule, and it has effectively been "donated" by its home molecule.
Things get reeeaaally different with Bronstead-Lowry bases. Remember the Arrhenius base with the hydroxide hanging off the side? While all Arrhenius acids are also Bronstead-Lowry acids, that rule is flat-out thrown away when it comes to bases! A Bronstead-Lowry base is a proton-acceptor. It finds those naked protons dangling off the edge acids, and its electrons slurp them up into its molecule. A Bronstead-Lowry acid is what took the proton in the previous analogy! Now, the hydroxide ion does do this, but plenty of molecules don't have any hydroxide ions and are perfectly capable.
Now comes the finale, the LEWIS acids and bases! If you've ever taken high school chemistry, you've probably made Lewis models where you wrote the chemical formula for an element and drew little dots around it, representing the electrons in outermost orbit. Usually in molecules electrons are pretty wrapped up in keeping the atoms attached to each other, but occasionally you'll see what's called a "lone pair", where there are two electrons hanging out, not really assigned to any atomic bond. A molecule with this is a Lewis base, an electron-pair donor. Those electrons getting bored whizzing around with no atoms to hold onto, so whenever they see a stray atom with few electrons to speak of, they tackle it and attach it to their molecule! (Sound familiar?) I find "donating an electron pair" to be a misnomer, because electrons are essentially the glue between the atoms. But I like to see it as donating a parking space for a lonely atom to come and hang out in. And what kinds of atoms come and hang out in those electron parking spots? Lewis acids! You don't really need to know a whole lot about these kinds of acids and bases, but I thought I'd throw it in there because it exists, and it was in my book.
Then there's that pH scale thingy. What the heck is it? There's actually a formula:
pH = -log[H+]
This means that it is the negative logarithm of the hydronium concentration. Acidity is a measurement of how many little protons manage to escape their molecules. A low pH means that it's SUPER acidic and ALL of them got loose, and as the pH gets higher, that concentration decreases by factors of 10, like the Richter scale. That's how that logarithm makes it work out. A high pH means that it's basic, containing lots of OH- ions. (Some of those Bronstead-Lowry bases can react with the water to make OH-! It retuuuurrrns!) There's another scale based on the OH- concentration called the pOH scale, but that's rarely ever used.
It's important to note that none of these definitions make the other definitions "wrong". They were developed at different times, and they all have proven to be useful, so we keep all three definitions around. You simply use whichever definition applies to your situation. You can use these to do some crazy math, like titration. Goodness, I never want to see a titration problem again.
Tune in next time for another episode of Science Student Explains Science! Leave questions in the comments!
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