ACIDS & BASES

Other than salts (such as table salt, sodium chloride), the next most well known example of ionic based bonding is found in acids and bases. But first, what is an acid? What is a base? Acids are defined as substances that are capable of donating a hydrogen ion, and bases as substances capable of accepting a hydrogen ion. Acids and bases can be either strong or weak. A strong acid (of which hydrochloric acid is a well known example) will be completely ionized when added to water. There will be no hydrochloric acid molecules left in the water, only chlorine ions (chloride) and hydrogen ions. The latter part of the last statement is a slight over simplification. Free hydrogen ion (H+) attaches itself to a water molecule. As seen in Figure 2, the actual form of hydrogen ions is that of the hydronium (H3O+) ion.

pH

When we measure pH we are actually measuring the concentration of the hydronium ion in the water(power of Hydronium, pH). Because the concentration of hydrogen ion is typically quite small, it is easier to refer to pH in terms of the negative logarithm of the concentration rather than the concentration itself (pH 7 is a little easier to recite than pH 0.0000001). This also explains why high acidity is expressed by low pH values (an otherwise counter intuitive convention). A weak acid or base is one that does not completely ionize when dissolved in water. A well known example of a weak base is ammonia. This example can also be used to highlight the central duality between acids and bases ( Figure 3).

When ammonia (NH₃), a base, accepts a hydrogen ion, it produces NH₄+, the conjugate acid of NH₃. When NH4+, an acid, releases a hydrogen ion, it forms the conjugate base, NH₃. The transformation of one into the other and back again results in an equilibrium. The relative balance between the two forms is described by an equilibrium constant, K. This balance is determined by the ratio of the product of the concentrations of the chemical products to the product of the concentrations of the chemical reactants ( Figure 4).

This constant is a function of the relative stability of the products and reactants in the selected solvent (typically water). The pH of the solvent can influence the relative position of the equilibrium but not the equilibrium constant itself. In other words, the equilibrium constant, K, does not change. The reason pH can influence the position of the equilibrium is explained by LeChatelier's principle, which says: "If a system at equilibrium is disturbed by a change in temperature, pressure, or the concentration of one of the components, the system will tend to shift its equilibrium position so as to counteract the effect of the disturbance."

This principle applies to any process involving some type of equilibrium. In ( Figure 5 we have simplified the ammonia/ammonium equilibrium in terms of only those components that see an appreciable change in their concentration. The concentration of water is basically constant, thus it can be incorporated into the equilibrium constant, K, to yield a new value, Ka (the acid dissociation constant). The negative log of Ka is the pKa, which for ammonium is 9.25. However, the main purpose of Figure 5 is to illustrate LeChatelier's principle by showing that lowering the pH (increasing the hydrogen ion concentration) results in the ammonia level dropping and the ammonium level rising. These changes occur in order to reestablish the equilibrium and return the ratio of product and reactant to its equilibrium value.

BUFFERS

A buffer consists of two parts (i) a weak acid or base and (ii) a salt of that weak acid or base. The equilibrium constant for the weak acid or base dictates what pH the buffer will most strongly maintain. The salt of the weak acid or base allows us to set the desired ratio between the conjugate acid and conjugate base so that the buffer will start out at the desired pH. For example, a bicarbonate based buffer employs a weak acid (carbonic acid) and a salt of that acid (sodium bicarbonate). To determine the resultant pH we use the Henderson-Hasselbalch equation ( Figure 6) which helps clarify how the ratio of acid and base affect the pH. In this equation we see that when the acid and base components are in equal concentrations the pH is equal to the pKa of the acid (6.35 for carbonic acid). Since the relationship between the acid and base ratio is expressed logarithmically, otherwise large shifts in the amount of acid or base produce very small changes in the final pH. However, once one exceeds more than a 5-fold excess of one component over the other the pH change becomes large. If the absolute amount of the buffer components is large, then the capacity for the buffer to resist a change in pH is large also.

GAS LAWS

Henry's law ( Figure 8) illustrates the relationship between the pressure of a gas above a solvent and the concentration of the gas in that solvent. As the pressure of the gas (Pg) increases, so does the concentration of the gas (Cg) in the solvent (k is a proportionality constant).

As temperature increases, solubility decreases for gases because dissolving gas into water is an exothermic process (gives off heat). Putting heat into the system will result in an increase of free gas as LeChatelier requires. At elevated temperatures the levels of dissolved CO₂ and O₂ (oxygen) decrease (see ( Figure 8: CO₂ concentration is given assuming a 10% CO₂ atmosphere which is the upper limit found on most CO₂/HCO₃- charts, O₂ is given assuming 100% O₂ atmosphere).

CO₂ and KH

The use of a carbonate buffer can be quite tricky because of one unique aspect that is not found in other buffer systems. Normally one gains or loses the acid component in a buffer with a change in pH. But with a carbonate buffer one can also gain or lose the acid component (carbonic acid) with a change in CO₂ concentration. CO₂ concentration is affected by temperature, aeration, or CO₂ pressure in the gas over the water. This brings us back to LeChatelier's principle. A decrease of CO₂ in the water will result in a decrease of carbonic acid in the water (see ( Figure 9).

The carbonic acid concentration decreases so as to bring the system back into equilibrium with the new lower level of CO₂. When the carbonic acid level drops, the bicarbonate level drops as well which in turn yields a higher pH (see Figure 9, bicarbonate: dropping levels of bicarbonate necessarily produce a drop in hydronium ion). Likewise, an increase in CO₂ will result in an increase in carbonic acid, followed by an increase in bicarbonate and its concomitant hydronium ion, which results in a pH decrease.

ALKALINITY and KH
There is much confusion over these terms. We will not delve into the unfortunate reasons for this right now, but rather will try to simply clear the air by defining and clarifying these terms: