Health Science Chemistry I CHEM-1180 Experiment No. 12 Acids, Bases, ph, Hydrolysis and Buffers (Revised 05/27/2015) Introduction Hydrogen Ion Concentration The acidity of aqueous solutions and its measurement is intimately associated with the autoionization of water. Pure water ionizes to a very slight extent according to the equation: The equilibrium expression for the reaction is given by: H2O(l) H + (aq) + OH - (aq) (Equation 1) [H ][OH ] [H O] 2 K (Equation 2) ion Since the concentration of water in most solutions is essentially constant at 55.5 mol/l, the expression can be rewritten as [H + ] [OH - ] = Kion [H2O] = Kw (Equation 3) It is important to note that hydrogen ion, written H +, does not exist as such in aqueous solutions. The small size and intense positive charge of a naked proton cause it to bond to any species having a lone electron pair. In aqueous solutions, that means that H + actually exists as the hydronium ion, H3O +. The formulas are used interchangeably in this laboratory discussion. Either representation is correct; [H + ] always has the same meaning, units and numerical value as [H3O + ]. At 24 C in pure water, when the hydrogen ion concentration, [H + ] and that of hydroxide ion, [OH - ] are identical, they each have a value of 1.00 x 10-7 mol/l. Thus, the value of the constant Kw is 1.00 x 10-14. For clarity and simplicity, molar units are customarily omitted from equilibrium constant values. However, molar units are never omitted from concentration values. Always use molar units, M or mol/l when reporting [H + ] and [OH - ]. Equation 3 holds for any H + and OH - concentrations. In any aqueous solution at 24 C, their product always equals 1.00 x 10-14. [H + ] for any solution can be found if [OH - ] is known and vise-versa. In the special case when [H + ] = [OH - ] = 10-7 M, the solution is neutral. When [H + ] > [OH - ] the solution is acidic. When [OH - ] > [H + ], the solution is basic. Equation 3 in the form [H + ] [OH - ] = Kw can be solved for either [H + ] or [OH - ] when the other quantity is known. The number of significant figures in any [H + ] value will be the same as the number in its corresponding [OH - ] value. Example 1: [H + ] = 10-3 M. What is [OH - ]? K 1.00 x 10 [H ] 10 M 14 W 11 [OH ] 10 M 3 Since [H + ] is greater than [OH - ], the solution is acidic. 1
Example 2: A solution has [OH - ] = 10-5 M. What is [H + ]? 14 KW 1.00 x 10 9 [H ] 10 M 5 [OH ] 10 M Since [H + ] < [OH - ], the solution is basic. Example 3: [H + ] = 3.0 x 10-8 M. Calculate [OH - ] for the solution. 14 KW 1.00 x 10 7 [OH ] 3.3x10 M 8 [H ] 3.0 x10 M This solution is slightly basic. Example 4: The [OH - ] of a solution equals 5.80 x 10-12 M. What is the solution [H + ]? 14 KW 1.00 x 10 3 [H ] 1.72 x10 M 12 [OH ] 5.80 x10 M The solution is acidic. Calculation of ph from [H + ] [H + ] can range from 10 0 M (1 M) to 10-14 M. Such a wide variation in [H + ] values is cumbersome to use and express mathematically. The ph concept was developed as a more concise way of expressing [H + ]. The ph scale does away with the necessity of stating the base ten, the negative sign of the exponent and molar units. ph is defined as the negative logarithm of the hydrogen ion concentration. ph = -log [H + ] (Equation 4) Since a logarithm is an exponent to the base 10, whole number phs are easy to calculate. Use Equation 4 to convert [H + ] to ph. For instance the log of 10-5 is -5 so the ph of a 10-5 M H + solution is 5. ph values for [H + ] between 10 0 M and 10-14 M are expressed as ph 0 to ph 14. Example 5: [H + ] = 10-4 M. What is the solution ph? ph = -log [H + ] = -log (10-4 ) = -(-4) = 4 When [H + ] values between whole number exponents or when fractional phs are encountered, a log table was once necessary to carry out the calculations. Now you may use a scientific calculator. The logarithms of numbers between 1 and 10 lie between 0 and 1. Refer to a suitable mathematics text for a detailed discussion of logs if you want a full understanding of ph calculations. Significant figures in ph calculations are handled as follows: The exponential term in any scientific notation expression does not contain any significant figures; only the significant part of the expression does. Since the significant part is properly expressed as a number between 1 and 10, its logarithm will be a decimal fraction with a value less than 1. The logarithm of any number between one and ten is properly expressed with as many significant figures as the original number. The logarithm of the exponential term is a negative integer and the logarithm of the significant part is a positive decimal fraction, which are then combined into a single term. For example, the calculation of the log of the number 3.4 x 10-9 can be broken down to the steps: log (3.4 x 10-9 ) = log 3.4 + log 10-9 = 0.53 + 9 = 8.47. Note that the two significant figures of the original number (the hydrogen ion concentration) become two decimal places in the logarithm. 2
To carry out [H + ] to ph conversions with a calculator, follow these steps: a) Enter the [H + ] value, preferably in scientific notation. b) Press the log key. c) The number on the display will be negative ph. Press the +/- key to convert it to the ph value. d) When calculating ph, to keep as many decimal places in the ph value as there are significant figures in the significant part of the original [H + ] value. Example 6 (Broken into steps to track significant figures): [H + ] = 2 x 10-6 M. What is ph? ph = -log [H + ] = -log (2 x 10-6 ) = -(log 2 + log 10-6 ) = - (0.3 + -6) ph = -(-5.7) = 5.7 Example 7 (As done on the calculator): [H + ] = 7.8 x 10-13 M. What is ph? ph = -log [H + ] = -log (7.8 x 10-13 ) ph = -(-12.11) = 12.11 Since the Example 7 hydrogen ion concentration 7.8 x 10-13 M lies between the values 1 x 10-12 and 1 x 10-13 M, its ph must lie between 12 and 13. Calculation of [H + ] from ph: To calculate a hydrogen ion concentration from a ph value, it is necessary to take the antilogarithm (antilog) of a number. An antilogarithm is simply the number for which a given logarithm stands. For instance, if log x equals y, then x is the antilogarithm of y. Hydrogen ion is calculated from the expression Example 8: The ph of a solution = 12. What is its [H + ]? [H + ] = antilog(-ph) = antilog(-12) = 10-12 M [H + ] = antilog(-ph) (Equation 5) The same significant figure rules apply here that were given earlier to convert [H + ] to ph, but in reverse order. The number of significant figures in the calculated hydrogen ion concentration should be the same as the number of decimal places in the ph. To carry out ph to [H + ] conversions with a calculator, do the following: a) Enter the ph value. b) Change the positive ph to a negative value with the +/- key. c) Find the antilog of the -ph value by pressing the second function (2nd) or inverse (inv) key and then pressing the log key. d) The [H + ] value will appear in the display. Express it in proper scientific notation, keeping only as many significant figures as there are decimal places in the original ph value. Do not forget to include molar units with every [H + ] value you report. Example 9: The ph of a solution = 7.8. What is its [H + ]? [H + ] = antilog(-ph) = antilog(-7.8) = 2 x 10-8 M Example 10: ph = 1.20. What is the solution [H + ]? [H + ] = antilog(-ph) = antilog(-1.20) = 6.4 x 10-2 M 3
Measurement of ph A simple way to tell if a solution is acidic or basic is to test a drop of the solution with litmus paper. You may have already done this in previous chemistry experiments. Litmus is a natural vegetable dye that changes color in response to hydrogen ion concentration. Blue litmus turns red in the presence of acid; red litmus turns blue in the presence of base. Other than those changes, litmus does not indicate even approximate ph values. (Litmus paper moistened with distilled water may also be used to test for the presence of acidic gases such as SO2 and the basic gas NH3.) Indicator or ph papers supplied by various manufacturers are impregnated with several dyes that turn a range of colors in response to ph changes in the range 0 to 12 or 0 to 14. The color of the test strip is compared to a standard card to get a ph reading to within one unit. Short-range papers exist that measure ph across a one or two unit range to within 0.1 unit. These papers may not work well in deeply colored or turbid solutions that mask the dye colors. Soluble indicator dyes are also used to determine solution phs. Indicators were once widely used to signal endpoints in acid-base titrations. For instance, phenolphthalein indicator is deep pink above ph 10 and colorless below ph 8.2. Phenolphthalein solution added to a solution of unknown ph can tell you whether the solution ph is above or below about 8.2. Several indicator dyes used with separate samples of the same solution can narrow the range of a solution ph value. Neutral red is pink below ph 6.8 and yellow above ph 8.0. If a solution gives no color with phenolphthalein and a yellow color with neutral red, then its ph must lie between 8.0 and 8.2. Thymol blue can be used to verify or narrow down these results since thymol blue is yellow below 8.0 and green in the ph range 8.0 to 9.6. The most accurate and precise instrument for measuring ph is the ph meter. A ph meter assembly consists of two probes, called the ph electrode and the reference electrode (sometimes combined into one housing), and an electronic meter with reads in the range 0 to 14 to a precision of 0.01 or even 0.001 ph units. Most ph meters read reliably to 0.1 unit; achieving a precision of 0.01 or more requires a more expensive meter, careful calibration and well cared-for electrodes. Hydrolysis The Arrhenius model for acids and bases defines acids are substances that donate hydrogen ion and bases as substances that donate hydroxide ion. This model proved to be deficient soon after its acceptance because it failed to explain why solutions of salts such as sodium acetate and potassium nitrite, while they contain no hydroxide ion in their formulas, give basic solutions. Equally puzzling was the acidity of ferric chloride solutions when the ferric chloride formula contains no hydrogen ion. To explain these phenomena, let us first review strong and weak acids and bases. A strong acid such as hydrogen chloride ionizes completely as shown: or HCl(aq) H + (aq) + Cl - (aq) (Equation 6) HCl(aq) + H2O(l) H3O + (aq)+ Cl - (aq) (Equation 7) Chloride ion has little affinity for or ability to bond covalently to H + ion. Since all the hydrogen atoms in aqueous HC1 become hydrogen ions, we refer to HC1 as a strong acid. On the other hand, most of the ionizable hydrogen atoms in the acetic acid molecule remain unionized in water. CH3COOH (aq) CH3COO - (aq) + H + (aq) (Equation 8) 4
Acetic acid is a weak acid because only a small percentage of the acetic molecules ionize. Acetic acid does not ionize readily because of the great affinity of acetate ion for hydrogen ion, H +. In the case of bases, NaOH is classified as strong because Na + does not bond covalently to OH -. Aqueous NaOH dissociates completely. Fe(OH)3 is an extremely weak base because Fe 3+ ion bonds strongly to OH - so most of the Fe(OH)3 remains insoluble and unionized. NH3 is a weak base because it produces few OH - ions when it interacts with water according to the equation: NH3 (aq) + H2O(l) NH 4 (aq) + OH - (aq) (Equation 9) In a solution of a salt such as NaC1, neither the Na + ion nor the Cl - ion interacts with water so the ph remains unchanged. In a solution of sodium acetate, the acetate ion removes some H + from water molecules, leaving behind OH - ions. The solution becomes basic due to hydrolysis or the reaction of acetate ion with water. The reaction is summarized as: CH3COO - (aq) + H2O(l) CH3COOH(aq) + OH - (aq) (Equation 10) To help you remember that such a solution is basic, think of sodium acetate as the product of a strong base (NaOH) and a weak acid (CH3COOH). Imagine that the properties of the stronger half of the pair (in this case, the base) predominate. For analogous reasons, FeC13 solutions are acidic. The Fe +3 ion reacts with water to produce an iron(iii) hydroxo complex and hydrogen ion: Fe +3 (aq) + H2O(l) Fe(OH) +2 (aq) + H + (aq) (Equation 11) Even though the ferric chloride formula contains no hydrogen atoms, the ferric ion generates protons by its reaction with water. NH4Cl solutions are acidic because the NH 4 ion dissociates slightly in water to give H + and the weak base NH3 according to the equation: NH 4 (aq) + H2O(l) NH3(aq) + H3O + (aq) (Equation 12) Chloride ion does not hydrolyze because it is unable to remove protons from water. To remember that FeC13 and NH4Cl are acidic salts, remember that each is the product of a weak base and a strong acid. Buffers Often it is desirable to have a solution whose ph is insensitive to the addition of small amounts of acid or base. This is especially true in biochemical systems where ph must remain in a narrow range for an organism to survive. Buffer solutions are widely distributed in nature and are often used for chemistry laboratory experiments. They act by their ability to absorb H + or OH - ions added to the system. Buffer solutions contain a weak acid and a salt of that acid with a strong base. This is equivalent to saying that buffers consist of a weak acid and its conjugate base. A typical buffer is a solution 0.1 M in acetic acid and 0.1 M in sodium acetate (acetate ion). The acetic acid component will react with small amounts of added base (hydroxide ion) to produce acetate ion (a base weaker than OH - ). The acetate ion component reacts with small amounts of added acid (H3O + ) to produce acetic acid, a weak 5
acid. Since the products of both reactions are already present in high concentrations, they exert only a small effect on the solution ph. Buffers can be tailor made to maintain any desired ph value. A buffer ph depends principally on the strength of the weak acid used to prepare it and to a lesser extent on the concentration ratio between that acid and it conjugate base. Buffers also happen to be insensitive to concentration changes caused by dilution because it is the ratio of the component concentrations that determines their ph. The ph of human blood is closely maintained between 7.35 and 7.45 by the action of three buffer 1 systems. They are the H2CO 3 / HCO 1 2 3 pair (carbonate buffer), the H2PO 4 / HPO4 pair (phosphate buffer) and the carboxylate and amino side groups on protein chains. Experimental Note: To reduce congestion at the ph meters and at the laboratory supply area, do not do the procedures in the order listed below. Procedures A and B must be done sequentially, but C, D and E should be done in random order so you can avoid waiting for supplies and equipment. A. ph Determination Using Indicator Paper Test each of the solutions in Table 1 with ph indicator paper. To avoid contaminating the solutions, do not put the paper indicator strips into them. Instead, dip a clean dry glass stirring rod into each solution and transfer a drop to a small piece of test paper. Immediately compare the color of the strip to the standard colors on the dispenser package and record the ph to the nearest unit. Table 1 Solutions to be Tested with Indicator Paper and the ph Meter 0.01 M HC1 0.1 M CH3COONa 0.01 M NaOH 0.1 M NH4C1 1.0 M CH3COOH 0.1 M NaHCO3 1.0 M NH3 0.1 M Na2CO3 Club Soda 0.1 M Na3PO4 Distilled Water 0.1 M NaCl Pour about 100 ml of each solution from its supply bottle into a clean, labeled plastic beakers. Be sure the stirring rods and ph electrodes placed in the solutions are clean. Do not return any solution to its original supply bottle. Report any possible contamination of a solution to your instructor. B. Measurement of ph by ph Meter The ph meter and its electrodes constitute a precision instrument that must be handled with care. Your instructor will demonstrate their proper handling and use. Newer ph electrodes have a glass ph electrode and a so-called reference electrode combined in one housing, protected by a perforated plastic sleeve. (ph electrodes have a delicate glass membrane that is easily damaged.) Test the solutions in plastic, not glass beakers to minimize the chance of electrode damage. Rinse the combined electrode before every use with distilled water from a wash bottle or by dipping it into a beaker of distilled water. Then gently blot the bottom of the electrode dry with a soft tissue. Store electrodes in water or buffer solution between readings to keep them from drying out and becoming unresponsive. 6
A ph meter must be calibrated prior to use. If your meter has not been calibrated, choose the standard buffer solution that is closest to the ph you will be measuring. For instance, if you want to measure a low ph, use a meter calibrated with ph 3 or ph 4 buffer. Put the clean electrode in the appropriate buffer solution, swirl the beaker, and after the (analog) needle stops moving or the digital display stabilizes, adjust the calibration setting so the needle points to the same ph value as that of the buffer. If the meter has a temperature-compensating knob, set it to room temperature, about 22 C. If you later want to measure a high ph, use a different meter calibrated with a high ph buffer or your reading may not be accurate. Ideally, ph meter response should be linear over the whole ph scale, but often it is not. Therefore, it is important to calibrate with a buffer near the ph you expect to measure. Measure the ph of each solution in Table 1. Use the approximate ph values you determined with ph paper to choose the most appropriate ph meter, that is, the one calibrated with a buffer closest to the solution ph value. Between readings, rinse the ph electrode with distilled water and gently blot it dry. Place the rinsed electrode in a beaker of each solution and swirl the beaker. Wait until the needle stops moving before you record a measurement. (Most analog and digital ph meters read to 0.01 ph unit.) C. ph Determination Using Indicator Solutions Use the indicator dyes in Table 2 to determine a ph range for the four solutions in Table 3. Place about 1 ml of one of the solutions in a clean test tube. Add 1 drop of crystal violet indicator, shake the test tube gently from side to side and note the color. If, for example, the mixture turns yellow, the solution is at ph 0 (or below) and you need test no further. However, if crystal violet gives a blueviolet color, the solution ph can be anywhere between 1.8 and 14 and you need to narrow the range by using the higher ph range indicators. Thymol blue is listed twice in the table because it undergoes two color changes, one at low and one at high ph. Next, test one ml of fresh solution with a drop of thymol blue indicator. A yellow color would mean the ph is 2.8 or above. If methyl orange gives a pink color, meaning ph 3.2 or below (between 0 and 3.2), you can stop testing the solution because you have narrowed the ph range to 2.8-3.2. On the other hand, if methyl orange turns yellow, it indicates only that the solution ph is 4.4 or above. Continue testing with higher range indicators until you find an upper ph value. Occasionally a solution turns an indicator a color between its two extremes. Bromothymol blue, for instance, is green between ph 6.0 and 7.6. With crystal violet, a green color means the solution ph is between 0.0 and 1.8. Table 2 ph Indicator Solutions Indicator Name Color ph Range Color ph Range Crystal Violet yellow 0 or below blue-violet 1.8 14 Thymol Blue pink 0 1.2 yellow 2.8 8.0 Methyl Orange pink 0 3.2 yellow 4.4 14 Bromothymol Blue yellow 0 6.0 blue 7.6 14 Neutral Red pink 0 6.8 yellow 8.0 14 Thymol Blue yellow 2.8 8.0 blue 9.6 14 Phenolphthalein colorless 0 8.2 deep pink 10.0 14 7
Table 3 Solutions to be Tested with Indicator Dye Solutions 0.01 M HC1 0.01 M NaOH 0.1 M NH4C1 0.1 M NaC2H3O2 D. Effect of a Buffer Solution on ph Stability Obtain about 100 ml of phosphate buffer solution, consisting of equal concentrations of sodium 1 dihydrogen phosphate and sodium monohydrogen phosphate ( H2PO 2 4 and HPO 4 ions). As mentioned earlier, these ions comprise one of the buffer systems in the human bloodstream. You will test the ability of this buffer solution to minimize ph changes in the following way: Use any one of the laboratory ph meters to measure the phosphate buffer solution ph. (It is not necessary to use a meter calibrated with a particular buffer.) Add one drop of 6M HCl to the buffer, swirl the solution and measure its ph again. Then add 1 ml of 6M HCl, swirl and measure the ph a third time. Using the same meter, repeat the experiment with 100 ml of distilled water instead of buffer solution. (Do not expect distilled water to be at ph 7.00. The ph of supposedly pure distilled water is usually lower than 7 because of dissolved CO2). Measure the ph of distilled water alone, distilled water with an added drop of 6 M HC1 and then distilled water mixed with 1 ml of 6 M HCl. Swirl each solution before you take the readings. E. ph of Unknown Solutions Obtain two unknown solutions. Record their numbers and measure their approximate ph values with indicator paper. Use this information to select the best ph meter for each solution, that is, the meter calibrated with a buffer nearest in ph to each unknown solution. Measure the ph of each unknown solution with the appropriate meter. (If the ph paper shows that your two unknown solutions have the same ph, ask the instructor to replace one of your unknowns.) Safety You must wear chemical splash goggles and a waterproof apron from the very beginning to the very end of the laboratory period. Disposal Unless your instructor directs otherwise, small amounts of the solutions used in this experiment may be safely flushed down the drain with water. Cleanup At the end of the laboratory period, wipe down all your work surfaces with a wet sponge. Return the unknown vials to the laboratory supply area. 8
Health Science Chemistry I CHEM-1180 Experiment No. 12 Acids, Bases, ph, Hydrolysis and Buffers Data A. ph of Solutions Measured with Indicator Paper Solution ph Solution ph 0.01 M HC1 0.1 M NaC2H3O2 0.01 M NaOH 0.1 M NH4C1 1.0 M CH3COOH 0.1 M NaHCO3 1.0 M NH3 0.1 M Na2CO3 Club Soda Distilled Water 0.1 M Na3PO4 0.1 M NaC1 B. ph of Solutions Measured with a ph Meter Solution ph Solution ph 0.01 M HC1 0.1 M NaC2H3O2 0.01 M NaOH 0.1 M NH4C1 1.0 M CH3COOH 0.1 M NaHCO3 1.0 M NH3 0.1 M Na2CO3 Club Soda Distilled Water 0.1 M Na3PO4 0.1 M NaCl 9
Health Science Chemistry I CHEM-1180 Experiment No. 12 Acids, Bases, ph, Hydrolysis and Buffers Data C. ph Ranges of Solutions Measured with Soluble Indicator Dyes Solution Tested 0.01 M HC1 ph Indicator Solutions Used Colors Observed with Each Indicator ph Range indicated by Each Individual Indicator Solution Final Solution ph Range 0.01 M NaOH 0.1 M NH4C1 0.1 M NaC2H3O2 10
Health Science Chemistry I CHEM-1180 Experiment No. 12 Acids, Bases, ph, Hydrolysis and Buffers Data D. Effect of a Buffer on ph Stability Solution ph Measured with the ph Meter Phosphate Buffer alone Phosphate Buffer plus one drop of 6 M HCl Phosphate Buffer plus one ml of 6 M HCl Distilled Water Alone Distilled Water plus one drop of 6 M HCl Distilled Water plus one ml of 6 M HCl E. ph of Unknown Solutions Unknown Solution Number ph Measured with ph Indicator Paper ph Measured with the ph Meter 11
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Health Science Chemistry I CHEM-1180 Experiment No. 12 Acids, Bases, ph, Hydrolysis and Buffers Report 1) Write the ph value measured by the ph meter for each of the 12 solutions in experimental part B. (4 points) Solution Name Measured ph Solution Name Measured ph 0.01 M HC1 0.1 M NaC2H3O2 0.01 M NaOH 0.1 M NH4C1 1.0 M CH3COOH 0.1 M NaHCO3 1.0 M NH3 0.1 M Na2CO3 Club Soda Distilled Water 0.1 M Na3PO4 0.1 M NaCl 2) Indicate which salt solutions in part B undergo hydrolysis and which ones do not. To answer this question properly, consider two factors for each salt solution: a) The measured ph value of each solution, and b) Your knowledge of which ions hydrolyze and which do not. (Note: Answer this question only for the salt solutions, not for the acid and base solutions. (6 points) 13
3) State the unknown number and the ph (measured with a ph meter) of your two unknown solutions. In the space below, calculate [H + ], [OH - ] and poh for each solution and state whether the solution is acidic or basic. For full credit, show clear setups for your calculations. (4 points) Unknown No. Unknown No. (4 points) 4) Summarize and explain the results you observed in part D when you added HCl to the phosphate buffer solution and to unbuffered distilled water. Explain the difference in the behavior of the two systems in terms of the chemical species involved. (2 points) 14
Health Science Chemistry I CHEM-1180 Experiment No. 12 Acids, Bases, ph, Hydrolysis and Buffers Prestudy 1) Calculate [H + ] for each solution and state whether the solution is acidic, basic, or neutral. (2 points) [OH - ] = 10-6 M [OH - ] = 8.81 x 10-8 M 2) Calculate [OH - ] for each solution and state whether the solution is acidic, basic, or neutral. (2 points) [H + ] = 4 x 10-2 M [H + ] = 1.9 x 10-13 M 3) Calculate [H + ] for each solution and state whether the solution is acidic, basic, or neutral ph = 10 (2 points) ph = 9.72 4) Calculate the ph for each solution and state whether the solution is acidic, basic, or neutral. (2 points) [H + ] = 10-13 M [H + ] = 7.30 x 10-5 M 5) What precautions are necessary when using ph electrodes? (1 point) 6) What is the ph or ph range of the solution that turns thymol blue indicator yellow and neutral red indicator yellow? (1 point) 15