5 posts • Page 1 of 1
I understand that in a equation, acid+base<->conjugate acid+conjugate base, the acid by the Bronsted definition, donates a proton to the base, but what does the conjugate bases or conjugate acids do? Do they follow the Lewis or Bronsted definitions of an acid or base? I can't see the relationship.
The conjugate base is the compound formed once the acid donates a proton. A conjugate acid would be formed when a base accepts a proton.
This also confused me a little, but reading the textbook helped clarify. In the example of HCO3- + O2-, HCO3- loses its H and donates it to O2- so it becomes OH- and HCO3- becomes CO3-. HCO3- is the Bronsted acid because it donated a proton, and O2- is the Bronsted base because it accepted a proton. CO3- is left after this Bronsted acid donates the H, which is the definition of the conjugate base. OH- is what remains after O2- accepts the proton, making it the conjugate acid. They're what's formed by the Bronsted acids and bases, so they're products of the acid and base reaction. The terminology is a little confusing at first but if you understand what's actually happening in the reaction then it makes a lot more sense. Hope this helps!
No I don't think so. However, usually when there's a strong conjugate acid/base, the "second step" of the equation involves the conjugate acid/base to take a proton from water and reform into it's original base/acid, and making the solution either basic or acidic.
Who is online
Users browsing this forum: No registered users and 1 guest