Equilibrium – Acid-Base Reactions

by Adam Le Gresley, PhD

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    00:00 or types of equations I’d like to introduce to you in this lecture.

    00:00 So, what do we mean by equilibrium? Equilibrium is very important when discussing any reaction, but it’s particularly relevant when we’re talking about acidity. So, for example, if we look at some very strong bases, it’s naturally assumed that virtually every single part of that compound is converted into H+ and its conjugate base. And that is shown here in the first equation. HA, which is our model acid, disproportionates into H+ and A-. But, that can also be, in the case of weaker... weaker acids, a backward reaction where H+ and A- come back together again to form our non-dissociated HA, which is our, again, our model acid.

    00:46 This equilibrium is shortened, as you can see at the bottom part of the board, where HA is shown to be in equilibrium with the H+ and conjugate base A-. So, in other words, what this means is we have a reaction going forward, where we have HA separating out into H+ and its conjugate base. And we have the reaction going backwards. And this is how it is often described. And, when you see these single-headed arrows going in opposite direction, you know that what you’re dealing with is a dynamic equilibrium. And that’s important for a whole number of different reasons.

    01:27 After a certain time, the speed of the forward reaction will be the same as the backward reaction. And the concentrations of the reactants and the products will remain stable or will remain steady state, that is to say, the concentration of neither the products nor the reactants will then change unless there is influence externally by varying, for example, temperature or pressure. The system is said to be, therefore, at equilibrium.

    01:58 So this is a model reaction which would be applicable to not just acid-base, but also to a whole number of different reactions. Well, we start off with two products, starting reactants A + B that are converted into C and D. So, again, irrespective of what reaction you’re talking about, reactants --> products. The equilibrium is always given as the combined concentration of products over the combined concentration of reactants. The brackets there around the ‘C’ denote the concentration of C. Likewise, the same for D, A and B.

    02:37 So, as you can see, an equilibrium constant is given by the concentration of one product multiplied by the concentration of the other product divided by the concentration of a reactant multiplied by the concentration of another reactant.

    02:54 Bear in mind, this equilibrium is a thermodynamic property. And there’s a fundamental difference between that and kinetics, which is something we may touch upon in future modules. But, crucially speaking, these are thermodynamic and therefore, this does not relate to the kinetics or the rate of formation, but rather, whether or not it is energetically more favourable or less favourable for the reaction to occur. So, here, as I’ve said, where [X] is the concentration of X, K is the equilibrium constant.

    03:31 If K is large, as you might expect, what this means is that the reaction will give mostly products with only small amounts of reactants remaining because, effectively, K is just a ratio metric measure of how much product we get in comparison to how much reactant remains. However, if K is small, only a small amount of product is formed and the majority of the compound will remain as the reactant.

    03:56 In the case of the dissociation of water, which I’ve shown you before, where you can get H+ and OH- which, for the sake of argument is represented as A- in this particular equation shown on the board, you can say that the equilibrium constant for this is the lysis of one concentration of acid HA into H+ and A-. And so, in this particular case, we can denote the equilibrium constant for this dissociation, as in this case, the acidity constant or Ka. Sometimes, you’ll come across other equilibrium constants which are specific, for example, for complexation and they’re often given as Kd, which is the equilibrium associated with dissociation. But, Ka is specifically how much of H+ is created from a given concentration of HA. The acid dissociation constant, which you’ll find in any textbooks which deal with this subject, is always given as Ka.

    05:01 If therefore a is large, it stands to reason that HA, our acid, will be almost completely dissociated giving a large concentration of H+. This is common in strong acids; strong acids being H2SO4 or sulphuric acid, HCl or hydrochloric acid or HNO3 or nitric acid.

    05:25 These are all good examples of strong acids where, because Ka is so large, it is uniformly accepted that the concentration in each of these cases directly correlates to the imparted concentration of H+ in a given solution.

    05:42 Acids, which only dissociate to a small extent, are called weak acids and these possess a small acidity constant or Ka. A good example of these are the carboxylic acids. We’ll be discussing carboxylic acids and the influences on their structure in a bit more detail in Module III, but for the moment, hopefully you can appreciate that where we have an organic acid, such as the carboxylic acid, shown in this particular case, the actual preference is largely for the formation of the acid and not dissociation to the conjugate base or carboxylate and the H+, which is obviously the measure of acidity. A good example of this is ethanoic acid. So, this is given as CH3CO2H, where the Ka is 1.8 × 10 to the -5, considerably smaller.

    06:39 The structural effects on acidity is something that relates to the concept of electronegativity and also, inductive effects, which we covered back in Module I. Since the process of dissociation generates ions, there is an advantage to stabilising the product ion to give a stronger acid, if indeed that is what you require.

    07:03 If you recall, we talked about the idea of introducing more electronegative atoms onto less electronegative atoms; an example being, of course, chlorine being attached via a sigma covalent bond to a carbon. We indicated that, in this particular case, because of the greater electronegativity of chlorine, the bond would be polarised. Indeed, the probability of finding electrons would be further moved towards the chlorine atom because it is more electronegative, pulling electrons away.

    07:34 The actual physical impact on this, as we can see in the case of this trichloroethanoic acid, is to actually render this compound more acidic. And this is because in each of the cases where we have a carbon-chlorine bond, the chlorine is pulling electron density away from that carbon which, in turn, is trying to pull electron density away from the carboxylate group, shown as the negative charge on the product side of the equation.

    08:05 By stabilising this negative charge, you’re essentially making this particular forward reaction more likely to occur resulting, of course, in a greater concentration of H+ or protons. So, therefore, the inductive effect of chlorine, in this particular case, although many halogens would have the same effect, fluorine in particular, helps to stabilise the negative charge on trichloroethanoic acid as shown above.

    08:31 And, if we were to, again, use dichloro- or monochloroethanoic acid as you can see at the bottom, we move from the strongest acid, which has the greatest amount of electronegative substituent, all the way down to the weakest acid which lacks any electronegative substituent.

    About the Lecture

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

    Included Quiz Questions

    1. 2.45 x 10-8 mol/L
    2. 9.92 x 10-7 mol/L
    3. 1.00 x 10-14 mol/L
    4. 4.08 x 10-7 mol/L
    5. 4.90 x 10-8 mol/L
    1. An acid loses one or more equivalent protons, whereas a base picks up an electron in an acid-base reaction.
    2. The acidity of a solution is the concentration of H+ or H3O+ ions in the water.
    3. An acid loses one or more equivalent protons, whereas a base picks up a proton in an acid-base reaction.
    4. During an acid-base reaction, a proton gets transferred from an acid to the base; so whole matter and charge remain conserved during the reaction.
    5. Water is a universal solvent for the vast majority of ionic reactions.
    1. An amphoteric substance or compound or system can act an acid only during a chemical reaction.
    2. Ampholytes being amphoteric molecules contain both acidic and basic groups in their molecular structure.
    3. Ampholytes can exist as zwitterions or inner salts or dipolar ions in a certain range of pH.
    4. Amino acids act as the best-known zwitterions due to the presence of ammonium and carboxylic group in the same molecule.
    5. An amphoteric substance or compound or system can act as both an acid or a base during a chemical reaction.

    Author of lecture Equilibrium – Acid-Base Reactions

     Adam Le Gresley, PhD

    Adam Le Gresley, PhD

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