Inorganic Chemistry - Water
Acids, Bases, pH and Buffers
Recall that the bonds that bind the oxygen and hydrogen together in water are polar covalent bonds and that covalent compounds typically do not dissociate. However, the polarity of water allows it to form hydrogen bonds with other water molecules in which the negative (oxygen) end of one water molecule is attracted to the positive (hydrogen) end of another water molecule. Although this is a weak attraction, occasionally the oxygen of one water molecule is able to steal the hydrogen from another water molecule, splitting the water molecules into ions. When this happens it results in the formation of a hydrogen ion (H+) and a hydroxide ion (OH-). Realize that in pure water, only a very, very few water molecules split—about 1 out of every 554,000,000 (who counted?). We can write the equation for this process like this:
H2O <-- --> H+ + OH-
Note that as with all chemical reactions, the reactants and products are in equilibrium and if that equilibrium is disturbed, the reaction will proceed until a new equilibrium is reached; hence the two headed arrow in the equation (<-- -->).
Acids and Bases
In pure water at 25o C, the concentration of H+ is always equal to the concentration of OH-. Both have a concentration of 1.0 x 10-7 Molar. (Placing the symbol for a chemical in brackets [H+] is chemical shorthand for “concentration of”, therefore [H+] is read “the concentration of hydrogen ion”). If we add a substance that results in an increase in [H+] we say that substance is an acid. If we add a substance that results in a decrease in [H+], we say that substance is a base. An acid is any substance that when added to, an aqueous solution increases the [H+] of the solution and a base is any substance that when added to, an aqueous solution decreases the [H+] of the solution. A common acid, for example is hydrochloric acid, HCl. When HCl reacts with water it dissociates into a H+ and a chloride ion (Cl-) thus increasing the [H+]. HCl is considered a strong acid because when placed in water it completely dissociates into its two ions.
HCl --> H+ + Cl-
A weak acid, such as acetic acid (CH3COOH) dissociates into H+ and CH3COO-, however, most remain as acetic acid and there is a chemical equilibrium between the CH3COOH and the H+ + CH3COO-.
CH3COOH <-- --> H+ + CH3COO-
An example of a base is ammonia (NH3), which will combine with H+ to form an ammonium ion (NH4+), thus removing H+ from the solution.
NH3 + H+ --> NH4+
Another common base is sodium hydroxide (NaOH). How is this a base? When it dissolves it dissociates into a sodium ion (Na+) and OH-, no change in [H+], right? However, the OH- will combine with H+ to form water, thus removing H+ from the solution.
NaOH --> Na+ + OH-
H+ + OH- --> H2O
pH
Why do we care about the [H+] anyway? What is special about this particular ion? Well, it turns out that either too much or too little H+ can cause us serious problems. If the [H+] is too low it causes an excitation of the nervous system, resulting in constant contraction of our muscles, including the respiratory muscles, and that is a problem. On the other hand, if the [H+] is too high it can result in depression of the nervous system, leading to coma. We use the terms acidic and basic to describe these conditions. If the [H+] of the solution is greater than 1.0 x 10-7, we say the solution is acidic, and if the [H+] is less than 1.0 X 10-7, we say that the solution is basic.
Because the [H+] is so important and because it is rather cumbersome to say things like, “the [H+] of the fluid is 1.0 X 10-7 Molar,” chemists have developed a shorthand to express the [H+]. This shorthand expresses the [H+] as the pH of the solution. The pH of a solution is the negative logarithm of the [H+] (concentration expressed as moles per liter, M). So, if the [H+] is 1.0 x 10-7 M the pH of that solution would be 7 (-log 10-7 is –(-7) or 7). Since this is the pH in which the [H+] and [OH-] are equal we say that this is a neutral solution. One thing that is a little confusing when we use pH is that as the [H+] of a solution goes up, the pH goes down. Suppose that a solution has a [H+] of 1.0 X 10-6 M, the pH of the solution would be 6 and the [H+] would be 10 times greater than a solution of pH 7. Thus, an acidic solution is any solution with a pH<7. Likewise, any solution that has a pH>7 is a basic solution. Here is an image that shows the pH of some common solutions.
Downloaded from Wikimedia Commons Fall 2014; Author: OpenStax College; License: Creative Commons Attribution 3.0 Unported license.
So, two important lessons from this, the lower the pH the higher the [H+] and a change in pH of one unit, 7 to 6 for example, is a 10 fold change in [H+]. Just for reference, the normal pH of our blood is slightly basic, 7.4 (range = 7.35 – 7.45). If the pH of the blood rises above 7.45 the person is in a state of alkalosis (not enough H+) and if it drops below 7.35, the person is in a state of acidosis (too much H+). In mammals, the pH range of the blood that is considered to be compatible with life is from 6.8 to 7.8. A pH above or below these values usually results in death.
Buffers
Because it is absolutely essential that blood pH be maintained within the narrow range, the body has several mechanisms to regulate the H+ concentration of the blood. One important defense employed by the body is the various buffer systems. Buffers are chemicals that tend to resist changes in pH. Note that buffers do not prevent changes, they resist changes. Let’s see if we can figure out how this works.
A typical buffer system is composed of a weak acid and the conjugate base of that acid. Remember, weak acids are those that do not dissociate completely but reach an equilibrium between the reactants and the products of the reaction. An important buffer system in our blood is the bicarbonate buffer system. The components of this system are shown below.
H 2CO3 <-- --> H + + HCO3-
Carbonic Acid Hydrogen Ion Bicarbonate Ion
In this case, the carbonic acid is the weak acid and the bicarbonate ion is its conjugate base. The entire reaction is in equilibrium. If the equilibrium is disrupted by the addition of more hydrogen ions, the reaction will proceed to the left until equilibrium is restored. When it proceeds to the left, some of the excess hydrogen ions will combine with bicarbonate forming carbonic acid, hence removing some of the excess hydrogen ions from the solution. Essentially, the buffer has “soaked up” some of the extra hydrogen ions, thus preventing a large change in pH. Another way of thinking of this system is to assume it behaves like a teeter-totter. If we have equal weights on each side, the teeter-totter is balanced (in equilibrium). If we add excess weight to one side (excess hydrogen ions), it will be out of balance. The only way to restore balance (equilibrium) is to move some of the excess weight to the opposite side until the teeter-totter is balanced again (equilibrium restored). Obviously, in this simple example, we realize that we cannot move all of the added weight to the opposite side because it would again be out of balance, but if some of the excess weight is moved to the other side, balance can be restored. Like the teeter-totter, when extra hydrogen ions are added, not all can be combined with bicarbonate, so there will still be a few more hydrogen ions than at the beginning (This is why buffers resist pH changes instead of prevent changes in pH). The pH will decrease, but not nearly as much as it would have if all added hydrogen ions were allowed to remain without being buffered. We could use the same analogy to see what happens when hydrogen ions are removed from the solution by the addition of a base. Since the equation is again out of equilibrium, the reaction will proceed, this time to the right until some of the hydrogen ions have been replaced. Again, there will be a slight increase in pH, but not nearly as great as would happen in the absence of the buffer.
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