4.3: Prelab Reading Assignment- Chemistry Review

In a chemical equation, variables that are surrounded by brackets “[“ and “]” are expressions of concentration, or the specific amount of a molecule in a given volume of solution.  For example, if you see “[\(\ce{H+}\)]” in an equation, this is read as “the concentration of hydrogen ion”.

The concentration of a solution is often expressed in units of moles per liter (mol/L).  Just as one “dozen” represents a quantity of 12 items, one “mole” represents a quantity of approximately 6.022 X 1023 items.

one dozen molecules = 12 molecules

one mole of molecules = 602,200,000,000,000,000,000,000 molecules!

Note

“n” is used in equations to indicate a quantity measured in moles.  For example if you see “nAcid” in an equation, this is read as “moles of acid”.

The term “Molarity” indicates that a solution’s concentration is in units of moles per liter. A one molar solution (1 M) contains one mole of solute within each liter of that solution. Reagents used in the laboratory will often be labeled with their concentrations expressed in terms of molarity.

The relative concentration of \(\ce{H+}\) or \(\ce{OH-}\) may change very dramatically in solutions, so a logarithmic scale (called pH) instead of a linear scale is used to express concentration.   These equations can be used to calculate the pH based on hydrogen ion concentration or vice versa:

\[\mathrm{pH}=-\log \left[\mathrm{H}^{+}\right] \quad \text { and }[\mathrm{H}+]=10^{-\mathrm{pH}} \nonumber \]

Buffers

A buffer is a mixture of a weak acid (\(\ce{HA}\)) and its salt (e.g., \(\ce{NaA}\)), and is sometimes referred to as a conjugate acid-base pair. As mentioned above, buffers have a major role in stabilizing the pH of living systems. Vertebrate organisms maintain the pH of blood using a buffer composed of a mixture of carbonic acid (\(\ce{H2CO3}\)) and sodium bicarbonate (\(\ce{Na+HCO3-}\)). The weak acid in this buffer is carbonic acid and the salt is sodium bicarbonate.  When dissolved in water, sodium bicarbonate disassociates completely into sodium ions (\(\ce{Na+}\)) and bicarbonate ions (\(\ce{HCO3-}\)). The \(\ce{H2CO3}\) is the conjugate acid of  \(\ce{HCO3-}\) and the \(\ce{HCO3-}\) is the conjugate base of \(\ce{H2CO3}\).  Together, this conjugate acid-base pair functions as the bicarbonate buffer system.   

Buffer systems are also of particular importance to experimental cell biology. 

The pH of a buffer solution may be calculated as follows:

\[\mathrm{pH}=\mathrm{pK}_{\mathrm{a}}+\log \dfrac{n_A}{n_{H A}} \nonumber \]

Where \(\mathrm{pK}_{\mathrm{a}}\) = dissociation constant of the acid, \(n_A\) = initial number of moles of salt in the buffer, and \(n_{H A}\) = initial number of moles of acid in the buffer. If you know these values, it is possible to accurately calculate the pH of a buffer system before you create it! 

The pKa of acetic acid (used in today’s experiment) is 4.75

\[n_A=\text { volume of conjugate base }(\mathrm{mL}) \times \frac{1 \mathrm{l}}{1000 \mathrm{ml}} \times \text { concentration of conjugate base }(\mathrm{mol} / \mathrm{L}) \nonumber \]

 

\[n_{H A}=\text { volume of conjugate acid }(\mathrm{mL}) \times \frac{1 l}{1000 \mathrm{ml}} \times \text { concentration of conjugate acid }(\mathrm{mol} / \mathrm{L})\nonumber \]

Use of pH Indicator Strips

The pH of a solution can be roughly approximated using strips of paper treated with color changing indicator reagents.  The strips are dipped into the solution to be tested for several seconds and then removed.  The color of the indicator strip is then compared to a reference chart, often printed on the side of the strip’s container.  The reference color on the chart that most closely matches the color of the reacted strip will have a pH value printed below it and that will be the approximate pH.  One advantage to using pH indicator strips is that they are relatively inexpensive, easy to use, and are adequate for determining pH where an error of +/- 1 pH unit is acceptable.  A more accurate method of determining pH is to use a calibrated pH meter, which can determine the exact pH to one or more decimal places depending on the quality of the device.

Use of a pH Meter

The pH meter is essentially a voltmeter designed to measure the potential difference between a reference electrode and a sensory electrode. The reference electrode will usually contain a solution of \(\ce{AgCl2}\). The external solution (or liquid to be tested) contains varying concentrations of protons. The instrument expresses the potential difference between these solutions either as a millivolt difference or translates it into a pH reading. Obtaining accurate readings with a pH meter depends on effective calibration, and the degree of static charge, and the temperature, as well as other factors. Particularly important to the proper use of a pH meter is accurate standardization, as glass electrodes must be carefully calibrated using buffers of known pH. Generally it is desirable to calibrate the meter with a standard buffer of a pH as close as possible to the pH of the solution to be measured. The problem of static charge buildup on the electrode can be eliminated by blotting (not rubbing) the surface of the electrode dry when transferring from one solution to another. 

Calibrate the pH meter for pH 4, 7, and 10 before taking measurements. If calibrated properly, your pH meter should produce measurements with an accuracy of +/- 0.06 pH units.  Always test your meter after calibration using the standard buffers and recalibrate the meter if necessary before proceeding.

 

Your instructor will demonstrate the proper calibration, care, and use of the meter.  Be sure to take good notes!