Unbelievable video always struggled with buffers due to the questionable teaching methods by my teacher. Thanks to this man I shall walk into my A2 exam with confidence.
Phenomenal video - genuinely so much easier spending 10 minutes watching these videos as opposed to trying to process the content of the text book (which are generally very poor)
Wouldn’t the equilibrium shift to the right for when acid is added to the acidic buffer, as the addition of H* ions will make more CH3COOH, so to counteract this change the equilibrium will shift to the right to produce more CH3COO- and H*. Could you help me out with this I am a bit confused.
i think it’s because it forms more CH3COOH on the RHS so to equal out the concentrations it shifts to the left, i’m not sure, this is how i remember it though.
Learn definitions of buffers (weak acid or base PLUS its corresponding salt) and weak acids/bases (ie only partially dissociate into H+ or OH- ions in water hence WEAK, so establishing an equilibrium/reversible reaction with them). Then apply LeChateliers principle - which is just like negative feedback mechanism in biology. Homeostasis is the term used to describe 'keeping things the same' - for example body temperature (an internal thermostat) or blood pH. Carbonic acid/hydrogen carbonate ions in blood is great example of an acidic buffer, which can resist SMALL changes in pH maintaining blood pH around 7.4 (vital to staying alive!). I hope that helps to give context for learning/understanding buffers! With regards to basic buffer though, you just apply LeChateliers principle in the same way as you would for acidic buffers - since XS OH- ions are added, the equilibrium must shift in the direction to OPPOSE this action(ie remove the XS OH-). So SOME (but not all) of the ammonium ions NH4 + present in solution (formed from both weak base ammonia and its salt) react with the XS OH- ions to reform ammonia NH3 (and water H20), so shifting equilibrium to the lefthand side. In this way, a buffer maintains a constant pH - but only for SMALL additions of acid or alkali! I hope this helps to clarify as buffers are tricky I know!
Thank you for the video, however I don’t get why the solution has to be a buffer solution in the situation where e.g OH- is added to an acidic buffer, cause the conjugate base from the salt doesn’t play a role in the reaction in resisting the change in pH
Am I correct in saying when you add an X-moles of acid to an acidic buffer, the moles of acid increase by X and the moles of salt decrease by X? ( And vice versa for adding an alkali to an acid) And then those new mole values are what you use for Ka calculations? Thank you :)
A buffer solution means resisting change in Ph right? So when an acid(H+) is added to that solution, the H+ reacts with the COO- and produces a higher concentration of COOH? Doesn't this mean the value of Ph decreases?
+Aashish_Menon Yes that's right. But remember pH is a measure of H+ ions. COOH dissociates weakly and hence does not produce many H+ ions. For this reason the pH value doesn't go lower. Hope this helps?
+Allery Chemistry Hi, I'm really confused about a certain part in the acidic buffer section. How come CH3COO-NA+ dissociates completely? I thought that for an acidic buffer, the salt has to be weak as well?
12:00 i dont get whats being opposed. surely if OH- reacts with NH4+ then you are reducing the concentration of NH4+ so i would assume the equilibrium shifts right to make more of the NH4+ thats been used up by reacting with OH-?
