Strength and Concentration – Acid-Base Reactions

by Adam Le Gresley, PhD

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    00:00 So, note the difference between strength and concentration. Strength relates directly to Ka, the acidity constant for a given dissociation. Concentration, on the other hand, relates purely to the amount per... per litre of a given ion. Strength refers to Ka and concentration refers to the amount of acid actually in solution. So, therefore, it's possible to have a concentrated solution of a weak acid or a dilute solution of a very strong acid. Concentration is measured in moles per litre or moles per dm3. And what you often find is it's referred as mol dm-3 and, sometimes, just as M (molar).

    00:54 Acid strength can be measured using Ka, as we've already indicated. The larger the Ka, the greater the degree of dissociation of the acid into the conjugate base and, of course, the important H+. But, these are not generally convenient ways of measuring it because you have to use, again, standard form in this case. 1.8 ×10 to the -5 is a little bit more difficult to work with, especially if you're looking at it from a biological perspective.

    01:18 And a lot of the calculations may appear to be confusing and may even result in some errors as a consequence.

    01:25 So, as a consequence of this, pKa is used with the small letter p denoting negative log (-log). That's all p means in this context. So, we take whatever the value of Ka is and then we carry out the negative log to the base 10 of that value.

    01:45 By carrying out the negative log to the base 10, we get a value which is a bit more easy to deal with. Usually, a value running from 1 through to 14 or, at the very least, a decimal value which is easy to deal with.

    01:58 If we look at this equation, what we mean is, by taking pKa, we are taking the negative log to the base 10 of the concentrations of H+ multiplied by a conjugate base A- divided by the non-dissociated concentration of our acid HA.

    02:15 So, if, for example, we were to look at ethanoic acid, we see that the pKa of this calculation results in a value of 4.77. The stronger acid, trichloroethanoic acid, results in a pKa of 0.70. And, as we actually go down in terms of pKa value, we're increasing the degree of dissociation and increasing the concentration of H+ that is actually formed.

    02:48 When an acid or a base is added to water, the concentration of H+ will change. And this is the origins of the term pH. pH, like pKa in this case, just means that we're taking the negative log to the base 10 of the concentration of H+ in solution. Hence, the term pH derived from the German expression potenz Hydrogen.

    03:14 Pure water has a pH of 7. And there's a relatively easy rule of thumb to bear in mind with this. When the concentration, for example, is approximately 1 x 10 to the minus 7 in terms of moles per litre, then you can actually just remove that -7 and remove the negative value. So, in other words, if I have 1 x 10 to the minus 7 moles of H+ in a litre of water, I will have a pH of 7. So, sometimes, there's a general rule of thumb which doesn't require you to use a calculator. Therefore, if we had a concentration of 1 x 10 to the minus 2, we can just take that power, remove the negative value and we end up with a pH of 2, which you should appreciate is an acidic pH.

    04:06 Remember, pure water has a pH of 7 and it also dissociates to a small extent. Acidic solutions will always give a greater concentration of H+ that exists in water and have a pH of less than 7. Basic solutions will have a smaller concentration of H+ free in solution and have a pH greater than 7, always.

    04:31 Typically, physiological conditions result in a pH of around 7, although obviously, that varies depending on whether or not you're talking, for example, about urine samples, saliva samples or indeed plasma blood samples.

    About the Lecture

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

    Included Quiz Questions

    1. 8.95 * 10^ -4
    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 have 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 acid dissociation in an aqueous solution.
    3. The concentration of an acid refers to the amount of the acid in the solution denoted by mol/dm3.
    4. The potential of hydrogen (pH) is a measure of hydrogen ion (H+) concentration and is given by the equation pH = -log [H+].
    5. Pure water dissociates to a smaller extent at pH 7.0.
    1. … the given acid is strong and dissociates into ions readily.
    2. … the given acid is weak and dissociates into ions readily.
    3. … the given acid is strong and does not dissociate into ions readily.
    4. … the given acid is weak and partially dissociates into ions.
    5. … the given acid is weak and does not dissociate into ions even at elevated temperatures.
    1. K = {[H+] × [C3H7COO-]} / [C3H7COOH]
    2. K = {[H+] × [C3H7COOH]} / [C3H7COO-]
    3. K = [H+] / {[C3H7COOH] × [C3H7COO-]}
    4. K = {[C3H7COOH] × [C3H7COO-]} / [H+]
    5. K = [C3H7COOH] / {[H+] × [C3H7COO-]}
    1. Carboxylic acids are stronger acids than mineral acids, such as HCl or HNO3.
    2. Weak acids dissociate to a small extent in an aqueous solution.
    3. Due to the generation of small amounts of H+ during dissociation, weak acids possess less acidic character than H2SO4.
    4. Ethanoic acid possesses a low acidity constant (Ka = 1.8 × 10-5).
    5. Oxalic acid gives a less stable conjugate base, hence it undergoes partial dissociation.

    Author of lecture Strength and Concentration – Acid-Base Reactions

     Adam Le Gresley, PhD

    Adam Le Gresley, PhD

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