LAB 12: ACIDS AND BASES: ELECTROLYTES, PH, AND BUFFERS

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1 LAB 12: ACIDS AND BASES: ELECTROLYTES, PH, AND BUFFERS PURPOSE: To determine the ph of common acids and bases using a ph meter, ph paper, and red cabbage indicator. To test the effect of adding an acid or base to a buffer solution. SAFETY CONCERNS: Always wear safety goggles. Wash with soap and water if skin contacts acids or bases. ACIDS: An Acid is a substance that when dissolved in water will produce hydrogen ions, H +, in the solution. An acid that does not contain carbon is called an inorganic acid. An acid that contains carbon is called an organic acid. Strong acids are acids that produce lots of hydrogen ions, H 1+ s. They almost completely dissociate (break apart) and the H 1+ s combine with water to form hydronium ions, H 3 O 1+, and the companion anions. Once dissolved in water the very strong acid does not exist as a covalent compound any more since it has dissociated into its ions. Solutions of strong acids behave as strong electrolytes due to the high concentration of ions present. Dissociation of Hydrochloric acid H-Cl + H 2 O H 3 O 1+ + Cl 1- Examples of Strong Acids: Strong Acid Formula Common Source Hydrochloric Acid HCl Stomach Acid Sulfuric Acid H 2 SO 4 Battery Acid Weak acids are acids that produce only a few hydrogen ions, H 1+ s. Only some of the molecules in solution dissociate (break apart) and form hydronium ions, H 3 O 1+ at any given time. When they do break apart, the hydronium ions, H 3 O 1+ formed and the companion anions join back together again to reform the covalently bonded acid. This process of the acid breaking apart and then reforming over and over is called equilibrium. Because of the equilibrium there is a mixture of the original acid, the hydronium ion, H 3 O 1+, and the anion all in the water solution at the same time. Solutions of weak acids behave as weak electrolytes due to the low concentration of ions present. Dissociation of Acetic acid H-C 2 H 3 O 2 + H 2 O H 3 O 1+ + C 2 H 3 O 2 1- Dissociation of Carbonic acid CO 2 + H 2 O H 2 CO 3 H 1+ + HCO 3 1- CH105. Lab 12: Acids and Bases (F15) 13

2 Examples of Weak Acids: Weak Acid Formula Common Source Acetic Acid HC 2 H 3 O 2 Vinegar Carbonic Acid H 2 CO 3 In Carbonated Water Citric Acid H 3 C 6 H 5 O 7 In Lemons & Oranges Tartaric Acid H 2 C 4 H 4 O 6 In Grapes Phosphoric Acid (A weak inorganic acid) H 3 PO 4 In Cola Drinks BASES: A Base is a substance that when dissolved in water will produce hydroxide ions, OH 1-, in the solution. A base that does not contain carbon is called an inorganic base. A base that contains carbon is called an organic base. Strong bases are bases that produce lots of hydroxide ions, OH 1-. They almost completely dissociate (break apart) to form hydroxide anions, OH 1-, and the companion cations. Solutions of strong bases behave as strong electrolytes due to the high concentration of ions present. Dissociation of Sodium Hydroxide NaOH Na 1+ + OH 1- Examples of Strong Bases: Strong Base Formula Common Source Sodium Hydroxide NaOH Lye, Caustic Soda, (strong) Soda Ash, Drano Potassium Hydroxide (strong) KOH Potash Weak bases are bases that produce only a few hydroxide ions, OH 1-. Sometimes it s not obvious that hydroxide anions, OH 1-, are even produced, but the anions that are produced can react with water to produce OH 1- s. Weak bases are in equilibrium with their ions. Solutions of weak bases behave as weak electrolytes due to the low concentration of ions present. Dissociation of Sodium Bicarbonate NaHCO 3 Na 1+ + HCO 3 1- HCO H 2 O H 2 CO 3 + OH 1-14 CH105. Lab 12: Acids and Bases (F15)

3 Examples of Weak Bases: Weak Base Formula Common Source Magnesium Hydroxide Mg(OH) 2 In Milk of Magnesia Ammonium Hydroxide NH 4 OH In Glass Cleaner Sodium Bicarbonate NaHCO 3 Baking Soda Calcium Carbonate CaCO 3 Antacids, Sea Shells, Egg Shells, Limestone & Marble PH: The ph of a solution is an indicator of the number of hydrogen ions, H 1+ s, present in a solution. We measure the concentration of hydrogen ions, H 1+ s, in moles per liter, M, which we symbolize with square brackets, [H 1+ ]. Pure water contains a small amount of both hydrogen ions, and hydroxide ions. In pure water the concentrations of hydrogen ion and hydroxide ions are the same. [H 1+ ] = [OH 1- ] = 0.000,000,1 M. It is awkward to work with numbers as small as 0.000,000,1 M. It helps to report the hydrogen ion concentration, [H 1+ ], in scientific notation. To simplify reporting the concentration of H 1+ s even more, we use the positive value of the exponent only and call it the ph. (in scientific notation) [H 1+ ] [H 1+ ] = 1x10 -ph (log means opposite of exponent on base 10) ph = -log [H 1+ ] 0.000,1 M 1 x 10-4 M ,000,1 M 1 x 10-7 M ,000,000,1 M 1 x M 10 If the [H 1+ ] = [OH 1- ] then the solution is considered to be neutral and the ph = 7 If the [H 1+ ] > [OH 1- ] then the solution is considered to be acidic and the ph < 7 If the [H 1+ ] < [OH 1- ] then the solution is considered to be basic and the ph > 7 [H 1+ ]= ph = Increasingly Acidic [H 1+ ] > [OH 1- ] Neutral [H 1+ ] = [OH 1- ] Increasingly Basic [H 1+ ] < [OH 1- ] CH105. Lab 12: Acids and Bases (F15) 15

4 The [H 1+ ] will not always be a simple 1 x 10 whole number value. In such cases we can estimate the ph range or we can calculate the ph exactly using the ph equation: ph = -log [H + ]. Examples: [H 1+ ] = 1x10 -ph ph = -log [H 1+ ] Given [H 1+ ] If [H 1+ ] = 0.000,032M = 3.2 x 10-5 M Calculated [H 1+ ] Then [H 1+ ] = 1 x = 2.5 x 10-9 M We can estimate ph to be between 4 & 5 We can estimate [H 1+ ] to be between 10-8 & 10-9 Calculated ph Then ph = -log(3.2x10-5 ) = 4.5 Given ph If ph = 8.6 INDICATORS: The ph of a solution is often measured by observing how the acid or base causes the color of certain organic molecules to change. Litmus, a chemical found in lichen, is one of many acid-base indicators. Litmus turns from blue to red in acidic solutions and from red to blue in basic solutions. Phenolphthalein, a laxative, is colorless when acidic but turns brilliant pink above ph 7. Anthocyanin indicators are common in flower petals, berries, and purple cabbage. The colors change over the entire spectrum of ph making it a universal indicator. The chart below shows the relative color changes of anthocyanin pigments over a wide ph range. Colors of Anthocyanin (Red Cabbage) Indicator ph = Color Red Fushia Rose Purple (neutral) Blue Aqua Green Greenbrown Yellow NEUTRALIZATIONS: A Neutralization reaction is a reaction between a strong acid and a strong base in which the product formed is neither acidic nor basic. For example, if hydrochloric acid (an acid) is mixed with sodium hydroxide (a base), the products will be sodium chloride (table salt) and water. Neutralization: Strong Acid + Strong Base Salt + Water HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O 16 CH105. Lab 12: Acids and Bases (F15)

5 BUFFERS: A Buffer is a solution that will not drastically change in ph even when strong acids or bases are added to it. A buffer is made by combining a weak acid and a weak base in a water solution. By having both an acid and a base in the solution together there will always be an acid to react with any base added, or there will be a base to react with any acid added. For example, a simple buffer could be made by combining equal amounts of the weak acid, Acetic Acid, and the weak base, Sodium Acetate. Buffer = Weak Acid + Weak Base HC 2 H 3 O 2(aq) + NaC 2 H 3 O 2(aq) (weak acid) (weak base) + NaOH If a strong base gets added to the buffer, it reacts with the weak acid present (rather than giving the solution more OH- s) and gets turned into NaC 2 H 3 O 2(aq) a weak base that does not drastically change the ph of the solution. + HCl If a strong acid gets added to the buffer, it reacts with the weak base present (rather than giving the solution more H+ s) and gets turned into HC 2 H 3 O 2(aq) a weak acid that does not drastically change the ph of the solution. As long as the number of OH 1- s and H 1+ s does not drastically change then the ph does not drastically change. Buffer solutions are important in keeping the ph of biological solutions in a narrow range. ph of a buffer: Consider a generic weak acid, HA. The equation for the dissociation of HA in water would be: HA + H 2 O H 3 O + + A - The K a expression for this equation would be: H O A Rearranging the K a expression gives us HA H 3O K a A A buffer could be made by adding equal amounts of HA and its conjugate base, A -. If the concentrations of HA and A - are HA 1 equal then A and so our equation goes to K a 3 HA [H 3 O + ] = K a. Since ph = -log [H 3 O + ] then the ph of our buffer is ph = -log K a CH105. Lab 12: Acids and Bases (F15) 17

6 PROCEDURES: ACTIONS: I. DETERMINATION OF PH OF COMMON ITEMS: 1. Arrange your large test tubes (except use small tubes for samples of toilet cleaner and lye) 3 and into each one put a different sample of a liquid 1 (about 1 inch deep) or a solid (about a dime sized scoop) of household product to be tested. 2. Add to each sample enough deionized water so that the probe of your ph meter can just be immersed into the liquid. Stir or stopper and shake each tube to dissolve. 3. Use these sample tubes for the sections A, B, & C where the ph of each will be determined by various methods. A. By Meter: 1. Clean the electrodes of your ph meter with deionized water between each sample test. 2 Be gentle with the electrodes as they are delicate and can easily break or scratch. 2. Carefully submerge the cleaned electrode of a ph meter into the samples indicated (but not into toilet cleaner or lye). 3 Read and record the ph on your report sheet to the accuracy of your meter. 3. Calculate the [H 1+ ] from your ph meter reading and record it on your report sheet. B. By Paper: 1. Tear the universal indicator ph paper into 1 cm square pieces and arrange them on a watch glass or paper towel Dip a glass stirring rod into the sample to be tested and touch the wet stirring rod to the ph paper. 3. Compare the color produced on the wet 5 paper to the color on the ph paper container label. Record the results. C. By Cabbage (Anthocyanins): 1. Tear a piece of purple cabbage 6 into small pieces and place in a glass beaker Add deionized water to barely cover the cabbage and then boil it using a hot plate, burner, or microwave oven until the water is dark purple. Remove from heat. NOTES: 1 If the liquid sample is highly viscous, (thick like shampoo) treat it as a solid. 2 We must have clean electrodes for each test. If you do not rinse off the previous sample then the new sample gets contaminated and the ph can be altered. 3 Toilet cleaner and drano/lye are very strong and so tax the flexibility of our meters. To avoid recalibrating our meters we will put these into small tubes in which the meter electrodes don t fit and we will test these with other methods. 4 Bring the sample to be tested to the ph paper rather than inserting the ph paper into the solution. Putting the ph paper into the solution washes pigment from the paper into the sample. 5 Compare the color when the paper is wet. The color will change as the paper dries and not compare accurately. 6 Roses, hydrangeas, and blueberries also contain anthocyanin pigments and will work as indicators. We use red cabbage because it is so inexpensive and readily available. 7 If working at home use a coffee mug or glass container. Metal pans cause the purple pigments to turn blue. 8 Because tea and coffee are so dark in color it is difficult to see any change in the cabbage pigment. So we will test these by other methods. 9 Compare the colors of your substance-cabbage mix with the colors on the chart in the discussion section. This is not exact but just allows you to predict a general ph range. 3. Into each test tube or beaker of sample to be tested, 8 add a few mls of purple cabbage juice and observe the color change. Record the ph range CH105. Lab 12: Acids and Bases (F15)

7 II. PH OF CARBON DIOXIDE, CO 2, SOLUTION: 1. Place about 10 mls of deionized water into a large test tube or small beaker and determine the ph using a ph meter Keeping the ph electrode in the tube, place a straw in the water and blow bubbles into it. 11 Watch the ph meter to observe changes. 3. Continue to blow bubbles into the water for a couple of minutes. Record any ph changes on the report sheet. III. ELECTROLYTIC PROPERTIES: 1. Obtain 2 small ( ml) beakers and label them A and B. Into the designated beaker pour in the following samples to a depth of about 1 cm: Beaker A: 0.1M hydrochloric acid (HCl) Beaker B: Glacial (100%) acetic acid (HC 2 H 3 O 2 ). 2. Plug a conductivity tester into an electrical outlet, being careful not to touch the electrodes. 12 Immerse 13 the tester into the solution A, 0.1M HCl, and observe and record the effect on the light bulb Unplug the conductivity tester and wash the electrodes with deionized water. 4. Repeat the conductivity test with beaker B, the glacial acetic acid. Record your results. 5. With the conductivity tester still immersed in the glacial acetic acid, beaker B, gradually add 50 mls of deionized water to the beaker a little at a time. Note and record the results. 15 IV. BUFFERS: 1. Obtain 4 large test tubes and label them Tubes A-D. Into both tubes A and B place about 10 mls of deionized or distilled water. Into both tubes C and D make buffer solutions by combining 8 mls of 0.1 M Sodium Acetate, (NaC 2 H 3 O 2 ) with 8 mls of 0.1 M Acetic acid, (HC 2 H 3 O 2 ) and mix well. A. Water with Acid: 2. Measure and record the ph of the water in tube A using a meter Add 1 drop of 1.0 M HCl, Hydrochloric Acid, and swirl the sample to mix. Record the ph. 10 The deionized water may not be at a ph of 7 depending on the amount of CO 2 gases already dissolved in it. Record the ph as you find it and use this as your starting point. 11 As you blow into the tube you are adding Carbon Dioxide, CO 2 to the water. CO 2 + H 2 O forms H 2 CO 3. Review the disassociation of Carbonic acid, H 2 CO 3, in the laboratory discussion of weak acids. 12 To avoid sever electrical shock, do not touch the electrodes. 13 When testing the solution, be sure the electrodes are immersed well beneath the surface of the solution. 14 A strong electrolyte will light a 60 watt bulb, a weak electrolyte will light a 15 watt bulb, and a nonelectrolyte will light neither bulb. Any degree of illumination will be assumed to be a positive test for an electrolyte. Whether the electrolyte is strong or weak may also be predicted from the brightness of the light. 15 Adding water to concentrated acid is generally an unsafe procedure however this particular addition may be done with caution. 4. Continue to mix in 1 drop of HCl at a time, recording the ph after each drop until a total of 5 drops have been added. Discard the solution in the sink. CH105. Lab 12: Acids and Bases (F15) 19

8 B. Water with Base: 5. Measure and record the ph of the water in tube B using a meter Add 1 drop of 1.0 M NaOH, Sodium Hydroxide, and swirl the sample to mix. Record the ph. 16 Not all buffers are neutral. The ph depends on the weak acid and the weak base used. Each buffer will have its own unique ph. Record the ph of yours and use this as your starting value. 7. Continue to mix in 1 drop of NaOH at a time, recording the ph after each drop until a total of 5 drops have been added. Discard the solution in the sink. 8. Analyze the results and report your conclusions on the report sheet. C. Buffer with Acid: 9. Measure and record the ph of the buffer solution in tube C using a meter Add 1 drop of 1.0 M HCl, Hydrochloric Acid, and swirl the sample to mix. Record the ph. 11. Continue to mix in 1 drop of HCl at a time, mixing and recording the ph after each drop until a total of 5 drops have been added. Discard the solution in the sink. D. Buffer with Base: 12. Measure and record the ph of the buffer solution in tube D using a meter Add 1 drop of 1.0 M NaOH, Sodium Hydroxide, and swirl the sample to mix. Record the ph. 14. Continue to mix in 1 drop of NaOH at a time, recording the ph after each drop until a total of 5 drops have been added. Discard the solution in the sink. 15. Analyze the results and report your conclusions on the report sheet. 20 CH105. Lab 12: Acids and Bases (F15)

9 LAB 12: ACIDS AND BASES: PRE LAB EXERCISES: NAME DATE Which of the following will provide the most accurate means of measuring ph? A. anthocyanin pigments B. litmus paper C. ph paper D. a ph meter E. All of these are the same. A solution has a ph of 8.7. The solution is A. Acidic B. Basic C. Neutral Which solution has the smaller hydrogen ion concentration? A. Solution A which has a hydrogen ion, [H 1+ ], concentration of 3.8 x B. Solution B which has a hydrogen ion, [H 1+ ], concentration of 2.3 x C. Not enough information to predict. Which solution has the lower ph? A. Solution A which has a hydrogen ion, [H 1+ ], concentration of 3.8 x B. Solution B which has a hydrogen ion, [H 1+ ], concentration of 2.3 x C. Not enough information to predict. Which solution has the higher ph? A. 0.1 M HCl. B. 0.1 M HC 2 H 3 O 2 C. Neither as both solutions have the same ph. Sodium hydroxide turns purple cabbage to a yellow color. Ammonium hydroxide turns purple cabbage green. Which solution has the lower ph? A. 0.1 M NaOH. B. 0.1 M NH 4 OH C. Neither as both solutions have the same ph. Which statement about a ph 5 solution is correct? A. The solution is neutral. B. The solution is a buffer. C. The hydroxide ion concentration is greater than the hydronium ion concentration. D. The hydronium ion concentration is greater than the hydroxide ion concentration. A buffer is A. a solution of acid and base that should always have a ph of 7. B. formed when a strong acid neutralizes a strong base. C. a solution of a weak acid and a weak base that resists change in ph. D. all of these. 9. Write the equation of Hydrochloric acid (HCl) with sodium acetate (NaC 2 H 3 O 2 ) in a buffer solution. 10. Write the equation of Sodium Hydroxide (NaOH) with acetic acid (HC 2 H 3 O 2 ) in a buffer solution. CH105. Lab 12: Acids and Bases (F15) 21

10 22 CH105. Lab 12: Acids and Bases (F15)

11 LAB 12: ACIDS AND BASES REPORT: I: Determination of ph by Meter, Paper, and Cabbage: Solution A. ph by Meter Water (for comparison) Deionized water used [H 1+ ] = 10 -ph B. ph by Paper NAME PARTNER DATE C. New Cabbage color produced ph by Cabbage Foods & Edibles Soda Pop brand Tea or Coffee (circle which) Orange Juice Vinegar Baking Soda NaHCO 3 Milk of Magnesia Antacid brand Cleaners: Shampoo Detergent Ammonia Bleach Toilet Cleaner Lye Salts: Table Salt: NaCl Kosher Salt: NaCl Washing Soda: Na 2 CO 3 Do not insert the meter into this solution Do not insert the meter into this solution OMIT as color is too dark to see a difference with cabbage. Epsom Salt: MgSO 4 Results Summary: 1. In general, foods and beverages are mostly and_ cleaning supplies are mostly A. Acidic B. Basic C. Neutral 2. Complete the following equations for formation of the tested salts and label the parent compounds and the salts produces as strong acid (SA), weak acid (WA) strong base (SB), weak base (WB) or neutral (N). + NaCl + H 2 O + Na 2 CO 3 + H 2 O Explanation/Analysis: Why were the results as they were? Explain any anomalies. CH105. Lab 12: Acids and Bases (F15) 23

12 II. ph of Carbon Dioxide, CO 2, in water: ph of Deionized Water = ph of Breath Carbonated Water = Complete and balance the equation for the reaction of carbon dioxide with water to form carbonic acid. CO 2 + H 2 O Results Summary: (Why did the CO 2 from your breath have the effect that it did?) 1. The ph of the original deionized water in part II may not be exactly 7.00 because A. there is dissolved carbon dioxide gas in the water that forms carbonic acid. B. there are dissolved hard water ion impurities in the water. C. water is naturally acidic due to the 1x10-7 M of H 1+ ions present. D. all of these. III. Electrolytic Properties: Observation 0.1 M Hydrochloric Acid: HCl Glacial (100%) Acetic Acid: HC 2 H 3 O 2 Acetic acid diluted with H 2 O HC 2 H 3 O 2 + H 2 O Complete the equation for the acid-base reaction between acetic acid and water: (Remember charges) HC 2 H 3 O 2 + H 2 O Results Summary: 1. My results show that Hydrochloric Acid is A. a strong electrolyte B. a weak electrolyte C. not an electrolyte. 2. My results show that Glacial (100%) Acetic Acid is A. a strong electrolyte B. a weak electrolyte C. not an electrolyte. 3. My results show that a solution of Acetic Acid in water is A. a strong electrolyte B. a weak electrolyte C. not an electrolyte. 4. When water is added to glacial acetic acid the equilibrium A. shifts to the right making more ions in solution. B. shifts to the left making more ions in solution. C. shifts to the right causing acetic acid to completely dissociate into ions. C. does not shift as it was already liquid in the beginning. Explanation/Analysis: Why were the results as they were? Explain any anomalies.. 24 CH105. Lab 12: Acids and Bases (F15)

13 IV. Buffers: Solution: ph Solution: ph A. Water with Acid C. Buffer with Acid Deionized Water NaC 2 H 3 O 2 :HC 2 H 3 O 2 Buffer After 1 drop 1.0M HCl After 1 drop 1.0M HCl After 2 drops 1.0M HCl After 2 drops 1.0M HCl After 3 drops 1.0M HCl After 3 drops 1.0M HCl After 4 drops 1.0M HCl After 4 drops 1.0M HCl After 5 drops 1.0M HCl After 5 drops 1.0M HCl B. Water with Base D. Buffer with Base Deionized Water NaC 2 H 3 O 2 :HC 2 H 3 O 2 Buffer After 1 drop 1.0M NaOH After 1 drop 1.0M NaOH After 2 drops 1.0M NaOH After 2 drops 1.0M NaOH After 3 drops 1.0M NaOH After 3 drops 1.0M NaOH After 4 drops 1.0M NaOH After 4 drops 1.0M NaOH After 5 drops 1.0M NaOH After 5 drops 1.0M NaOH Explanation/Analysis: (Why were the results as they were? When adding HCl or NaOH to water and a sodium acetate/acetic acid buffer solution, what caused the ph to change or not change as it did? Explain the chemistry that was taking place. Explain any anomalies) Results Summary: 1. In general, adding a strong acid or a strong base to water A. neutralizes the solution making a solution that is neither acidic nor basic. B. drastically changes the ph to be either strongly acidic or strongly basic. C. causes only a slight change in the original ph making the solution a only a little more acidic if strong acid is added, or only a little more basic if strong base is added. D. has variable results as there is no general trend 2. In general, adding a strong acid or a strong base to a buffer solution A. neutralizes the solution making a solution that is neither acidic nor basic. B. drastically changes the ph to be either strongly acidic or strongly basic if strong base is added. C. causes only a slight change in the original ph making the solution a only a little more acidic if strong acid is added, or only a little more basic if strong base is added. D. has variable results as there is no general trend 3. A. The K a for acetic acid (HC 2 H 3 O 2 ) is 1.8x10-5. Calculate the expected ph of a buffer solution made by mixing 5 mls of 0.1 M Acetic acid with 5 mls 0.1 M Sodium Acetate, (NaC 2 H 3 O 2 ). Show your calculations. B. How does the ph of the buffer you made in lab compare to that theoretical calculation? A. it was right on (within the range of appropriate significant figures.) B. it was close (within +/- 1 ph increments) to theoretical. C. it was within +/- 2 ph increments of theoretical. D. it was > +/- 2 ph increments from theoretical. Analysis: Why were the results as they were? Explain any anomalies between the theoretical ph and your experimental ph. CH105. Lab 12: Acids and Bases (F15) 25

14 LAB 12: ACIDS AND BASES NAME Related Exercises: Use the following equilibrium equation for the formation and dissociation of carbonic acid to answer questions 1-5: CO 2 + H 2 O H 2 CO 3 H HCO A sample of deionized water had a ph of 6.0 until it was boiled. Once this water was boiled and cooled it had a ph of 7. Which statement might best explain why this could happen? A. Boiling will decrease the quantity of a gas that will dissolve in water so boiling would cause the CO 2 to escape thus driving the equilibrium left and lowering the concentration of H 1+. B. Boiling will increase the quantity of a gas that will dissolve in water so boiling would cause the CO 2 to become more soluble thus driving the equilibrium right and raising the concentration of H 1+. C. Boiling has no effect on the quantity of CO 2 gas that will dissolve in water so this change in ph must have been a fluke. D. Boiling would kill bacteria in the water that is responsible for acidic ph. Acidosis is a condition in which the blood ph drops below normal. Acidosis can be caused by metabolic or respiratory conditions. Which of the following best explains the role of carbon dioxide in producing respiratory acidosis? A. Decreased concentration of CO 2 can shift the equilibrium to the left causing a decrease in carbonic acid which in turn causes a decrease in the [H + ] causing a decrease in blood ph. B. Decreased concentration of CO 2 can shift the equilibrium to the left causing a decrease in carbonic acid which in turn causes a decrease in the [H + ] causing an increase in blood ph. C. Increased concentration of CO 2 can shift the equilibrium to the right causing an increase in carbonic acid which in turn causes an increase in the [H + ] causing a decrease in blood ph. D. Increased concentration of CO 2 can shift the equilibrium to the right causing an increase in carbonic acid which in turn causes an increase in the [H + ] causing an increase in blood ph. A. A person who is hyperventilating could suffer from A. alkalosis B. acidosis C. Neither of these B. Explain your answer in terms of the given carbonic acid equilibrium equation given above and ph. C. Explain in terms of the equilibrium why breathing into a paper bag is a treatment for hyperventilation. 4. Bacteria in the mouth convert sugar into acids that are capable of dissolving tooth enamel. Repeated drinking of carbonated beverages, even those that do not contain sugar, can also contribute to loss of enamel. Explain why carbonated beverages can act in this way and cause dental caries (tooth decay). 5. Show the reaction with stomach acid that explains how a Tums antacid (CaCO 3 ) settles an upset stomach. Account for any gas formed. 6. A student measured the ph of Caffeinated coffee as ph 5.6 and Decaffeinated coffee as ph 5.0. On the basis of this data, would you classify caffeine as an acid or base? A. acid B. base Reference Search: 7. Use a reference source of your choice to find the following information about Washing soda, Na 2 CO 3 Solubility in water at room temp: Common uses: g/100 ml Reference Used: 26 CH105. Lab 12: Acids and Bases (F15)