Henderson-Hasselbalch equation = where it comes from & when and how to use it скачать в хорошем качестве

Henderson-Hasselbalch equation = where it comes from & when and how to use it

pH is a measure of, acidity, which refers to how many free protons (H⁺) are hanging out in a solution. I say “free” but in reality most protons in a water-based (aqueous) solution quickly latch onto a water molecule to become a hydronium ion (H₃O⁺). An acid (in one definition) is something that donates proton(s) and a base (in one definition) is something that accepts protons. So the pH is a combination of the effects of all the acids & bases present in a solution. pH is an inverse log of the proton concentration: pH = log10(1/[H⁺]). Dividing by something bigger gives you something smaller, and all the log part does is make the numbers smaller and easier to deal with, so the more protons there are (the more acidic the solution) the lower the pH. We call pH 7 “neutral” - above it (fewer protons) & we call a solution BASIC or ALKALINE - lower than 7 (so more protons) and we call it ACIDIC I like to think of a sort of “proton cookie jar” where the jar fullness corresponds to pH (fuller jar → lower pH → more acidic) & how full the jar is at any time depends on how many proton-givers (acids) & proton-takers (bases) there are. But that’s not all - it also depends on how “stingy” they are - how willing are to give or take? An acid is a molecule that can donate a proton (H⁺), but that doesn’t mean it always wants to…Instead of thinking in terms of acids wanting to get rid of H⁺, it sometimes helps to think of how reluctant they are to give them up. An acid’s willingness to give up H⁺ depends on its surrounding environment - how much “free” H⁺ is already out there. If there are plenty of H⁺ around, the acid’s desire to donate is low. BUT if there aren’t many H⁺ around, the acid “feels guilty” & donates 1 to help out. pKa is a measure of “how extreme” (i.e. basic/alkaline) conditions must be in order for an acid to give up its H⁺ quick note on notation. An acid can be anything that can donate a proton, so we often use “HA” as a generic shorthand for an acid and A⁻ as a generic shorthand for its deprotonated form. So HA → H⁺ + A⁻. But this is reversible. A⁻ can now snatch up a proton to become HA again, H⁺ + A⁻ → HA, and since proton-snatchers are called bases, we can call A⁻ the “conjugate base” of HA. This HA/A⁻ notation is typically used for what are referred to as “weak acid/bases.” Theoretically, any acid/base reaction can go both ways, but some acid/bases are sooooo much happier in one form (i.e. a strong acid is way happier deprotonated and a strong base is way happier protonated) that protonation/deprotonation is virtually “irreversible.” Examples of such strong acids are: HCl (hydrochloric acid), HNO₃ (nitric acid), and H₂SO₄ (sulfuric acid) and a couple strong bases are: NaOH (sodium hydroxide) and KOH (potassium hydroxide). You might come across these in the lab in biochemistry, but when it comes to our bodies, we’re usually dealing with weak acids & bases, which can be in either their conjugate acid or conjugate base forms. And figuring out how much of which form is present means dealing with their pKa. ⚠️ Don’t confuse pKa w/pH. pH is a measure of the total concentration of protons ([H⁺]) from any source. This includes: water itself: H₂O ⇌ H⁺ + OH⁻ check out this post for more on such “auto-ionization of water:” http://bit.ly/phacidbase AND/OR acid(s)(HA) dissolved in it: HA ⇌ H⁺ + A⁻ ⚠️ pKa & pH are different: pKa is a (constant) property of a molecule while pH is a (changeable) property of a solution pH & pKa are different BUT directly related through the Henderson-Hasselbalch equation: pH = pKa + log[A⁻]/[HA] pKa is the pH @ which 1/2 of the acid molecules have given up a H⁺ @ any pH higher than an acid’s pKa, a molecule of that acid is more likely to be deprotonated than protonated the chances increase the further above pKa you are. For example, 1% deprotonated @ 2pH units below pKa & 99% deprotonated @ 2pH units above pKa. Different acids have different “willingness thresholds” STRONGER acids are more generous & have LOWER pKas - even a slight deficit of H⁺ in surroundings & they’ll give 1 up WEAKER acids are “greedier” - they have HIGHER pKas meaning they won’t give up H⁺ until there’s a big deficit For example, say you have 3 acids: X w/pKa of 3, Y (pKa of 7), & Z (pKa of 9) at neutral pH (7), most of X will be deprotonated, 1/2 of Y, BUT almost none of Z But if we raise pH (so there’s fewer free H⁺ available) even stingy Z will give up its H⁺ If you want to know what proportion of an acid will be deprotonated ([A⁻]/[HA]) you can calculate that using a rearrangement of the Henderson-Hasselbalch equation (if you know the pH and the pKa) pH = pKa + log[A⁻]/[HA] [A⁻]/[HA] = antilog(pH-pKa) when pH = pKa, you get a ratio of 1 (equal amounts of protonated form (conjugate acid) & deprotonated form (conjugate base) when pH is greater than pKa, [A⁻] is greater than [HA] when pH is less than pKa, [A⁻] is less than [HA] You can plug in numbers to get values.