In practice, the pH is not usually measured in this way because it requires hydrogen gas at standard pressure, and the platinum electrode used in the standard hydrogen electrode is easily fouled by the presence of other substances in the solution (Ebbing). Fortunately, other practical electrode configurations can be calibrated to read the H+ ion concentration. Laboratory pH meters are often made with a glass electrode consisting of a silver wire coated with silver chloride immersed in dilute hydrochloric acid. The electrode solution is separated from the solution to be measured by a thin glass membrane. The potential which develops across that glass membrane can be shown to be proportional to the hydrogen ion concentrations on the two surfaces. In the measurement instrument, a cell is made with the other electrode commonly being a mercury-mercury chloride electrode. The cell potential is then linearly proportional to the pH and the meter can then be calibrated to read directly in pH.
In this video, Christine Muth of the North Carolina School of Science and Mathematics demonstrates how to measure pH of a water sample using two types of pH paper.
The above discussion treats changes in pH quantitatively. These changes in pH can be described qualitatively, also. A qualitative view is very useful for predicting how the pH will change in response to external conditions (such as exercise). The principle used for this qualitative view is known as Le Châtelier's Principle.
This interesting fact was also emphasized by Gärtner, with the remark that even the most practiced expert is not in a position to determine in a hybrid which of the two parental species was the seed or the pollen plant.Of the which were used in the the following are :With regard to this last character it must be stated that the longer of the two parental stems is usually exceeded by the hybrid, a fact which is possibly only attributable to the greater luxuriance which appears in all parts of plants when stems of very different lengths are crossed.
Other buffers perform a more minor role than the carbonic-acid-bicarbonate buffer in regulating the pH of the blood. The phosphate buffer consists of phosphoric acid (H3PO4) in equilibrium with dihydrogen phosphate ion (H2PO4-) and H+. The pK for the phosphate buffer is 6.8, which allows this buffer to function within its optimal buffering range at physiological pH. The phosphate buffer only plays a minor role in the blood, however, because H3PO4 and H2PO4- are found in very low concentration in the blood. Hemoglobin also acts as a pH buffer in the blood. Recall from the from Chem 151 that hemoglobin protein can reversibly bind either H+ (to the protein) or O2 (to the Fe of the heme group), but that when one of these substances is bound, the other is released (as explained by the Bohr effect). During exercise, hemoglobin helps to control the pH of the blood by binding some of the excess protons that are generated in the muscles. At the same time, molecular oxygen is released for use by the muscles.
The slope of the curve is flattest where the pH is equal to the pK value (6.1) for the buffer. Here, the buffering capacity is greatest because a shift in the relative concentrations of bicarbonate and carbon dioxide produces only a small change in the pH of the solution. However, at pH values higher than 7.1, the slope of the curve is much higher. Here, a shift in the relative concentrations of bicarbonate and carbon dioxide produces a large change in the pH of the solution. Hence, at the physiological blood pH of 7.4, other organs must help to control the amounts of HCO3- and CO2 in the blood to keep the pH relatively constant, as described above.
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However, the normal blood pH of 7.4 is outside the optimal buffering range; therefore, the addition of protons to the blood due to strenuous exercise may be too great for the buffer alone to effectively control the pH of the blood. When this happens, other organs must help control the amounts of CO2 and HCO3- in the blood. The lungs remove excess CO2 from the blood (helping to raise the pH via shifts in the equilibria in Equation 10), and the kidneys remove excess HCO3- from the body (helping to lower the pH). The lungs' removal of CO2 from the blood is somewhat impeded during exercise when the heart rate is very rapid; the blood is pumped through the capillaries very quickly, and so there is little time in the lungs for carbon dioxide to be exchanged for oxygen. The ways in which these three organs help to control the blood pH through the bicarbonate buffer system are highlighted in Figure 3, below.
As shown in Equation 11, the pH of the buffered solution (i.e., the blood) is dependent only on the ratio of the amount of CO2 present in the blood to the amount of HCO3- (bicarbonate ion) present in the blood (at a given temperature, so that pK remains constant). This ratio remains relatively constant, because the concentrations of both buffer components (HCO3- and CO2) are very large, compared to the amount of H+ added to the blood during normal activities and moderate exercise. When H+ is added to the blood as a result of metabolic processes, the amount of HCO3- (relative to the amount of CO2) decreases; however, the amount of the change is tiny compared to the amount of HCO3- present in the blood. This optimal buffering occurs when the pH is within approximately 1 pH unit from the pK value for the buffering system, i.e., when the pH is between 5.1 and 7.1.
Dye indicator solutions or paper have the advantage of being quiteinexpensive, very portable, and often suitable where only an approximatepH measurement is needed.
The take home point here is that the scientific format helps to insure that at whatever level a person reads your paper (beyond title skimming), they will likely get the key results and conclusions.
The derivation for this equation is shown in the yellow box, below. Notice that Equation 11 is in a similar form to the Henderson-Hasselbach equation presented in the introduction to the Experiment (Equation 16 in the lab manual). Equation 11 does not meet the strict definition of a Henderson-Hasselbach equation, because this equation takes into account a non-acid-base reaction (i.e., the dissociation of carbonic acid to carbon dioxide and water), and the ratio in parentheses is not the concentration ratio of the acid to the conjugate base. However, the relationship shown in Equation 11 is frequently referred to as the Henderson-Hasselbach equation for the buffer in physiological applications.