Bases – Acid-Base Reactions

by Adam Le Gresley, PhD

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    00:00 saliva samples or indeed plasma blood samples.

    00:00 Let’s get onto bases. So, we’ve talked about acids, we’ve talked about the dissociation of acids into conjugate base and H+. Now, let’s talk about bases.

    00:12 So, there are a number of different bases. These can be charged bases, such as hydroxide, or these can be uncharged bases, for example, like ammonia and amine derivatives.

    00:23 What happens in the case of a base, as we said before, is that it can attract H+ or a proton resulting in the formation, as you can see from this particular equation on the board, of BH+ and OH-. Okay? BH+ being the conjugate acid in this particular case as opposed to the conjugate base as we saw in the case of acidity.

    00:47 As with the acids, the strength of the base can be shown by the equilibrium constant, in this case Kb. And this results in pKb also being able to be determined. However, and this is something to bear in mind, more commonly, you’ll find that, in the interest of clarity, a base strength is actually measured by pKa. Now, this is not, contrary to understanding, the loss of a hydrogen or proton from the base. This is the loss of the H+ from the formed conjugate acid. So, this equation here, as you can see on the board, this shows that BH+, which is our conjugate acid dissociating to give us our free base and our H+, correlates actually to a dissociation or acidity constant pKa. The equation for this is shown on the board.

    01:44 Weak bases, therefore, have very low pKa values and strong bases have very high pKa values.

    01:53 And when we consider any acid or base, we need to know or should be aware that pKa + pKb = 14.0.

    02:06 Strong bases include things like sodium hydroxide. Indeed, mostly every hydroxide is regarded as a strong base. And as we said before in case... in the case of hydrochloric acid, sulphuric acid and nitric acids, they’re considered to be almost completely dissociated in water. And so, therefore, in the case of our NaOH, we see that sodium ion and hydroxide ion are formed in solution.

    02:32 Now, of course, what you’ll be talking about and may wonder is, well, we were talking about pH being potenz Hydrogen, we’re talking about H+. So, where does OH- come into this? We’re not measuring OH- when we’re looking at pH, we’re measuring H+. And this is because OH- will bind with any residual or large amount of residual H+ in water and convert it back into water because of the equilibrium of H+... OH- in water and the preference for it to exist as non-ionised water. This decreases the amount of H+ in solution and therefore, increases the pH, making it more basic.

    03:13 Amines are also examples here of weak bases. Remember, we had weak acids. We can also get weak bases. And here we see the pKa. Remember, that is the equilibrium constant associated with the dissociation of the conjugate acid being 9.25. Methylamine is 10.66 pKa. And the equilibrium in this particular case for ammonia, and this is how you would form ammonium hydroxide solution, is where we have our base NH3, which is a neutral base, abstracting a proton from H2O to give us NH4+, a complex cation otherwise known as the ammonium ion and OH-, hydroxide ion.

    About the Lecture

    The lecture Bases – Acid-Base Reactions by Adam Le Gresley, PhD is from the course Ionic Chemistry.

    Included Quiz Questions

    1. 1.0 x 10-3
    2. 9.5 x 10-2
    3. 3.1 x 10-3
    4. 7.6 x 10-4
    5. 2 x 10-3
    1. Acidic solutions give smaller concentrations of H+ ions, while basic solutions contain greater amounts of free H+ ions.
    2. The strength of an acid pertains to the extent of dissociation in an aqueous solution.
    3. The concentration of an acid refers to the amount of an acid in the solution denoted by mol/dm3.
    4. The potential of hydrogen (pH) is a measure of hydrogen ion (H+) concentration to measure the acidity or basicity of a solution and is given by the equation pH = -log [H+].
    5. Pure water dissociates to a smaller extent at pH 7.0.

    Author of lecture Bases – Acid-Base Reactions

     Adam Le Gresley, PhD

    Adam Le Gresley, PhD

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