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Principles of Chemistry Lab II Montgomery College, Rockville 1 Acids and Bases, pH, Buffers and Hydrolysis Introduction Acids and Bases Aqueous solutions of acids and bases are recognized as “acidic” or “basic” because they contain appreciable concentrations of either hydronium (H3O+ ) or hydroxide (OH– ) ions. Hydronium ions are produced from the reaction of covalent molecules like HCl with water. HCl (g) + H2O (l) H3O+ (aq) + Cl– (aq) Some bases are ionic compounds that dissolve in water or react with it to produce aqueous hydroxide ions, for example: NaOH (s) Na+ (aq) + OH– (aq) BaO (s) + H2O (l) Ba2+ (aq) + 2OH– (aq) Other bases such as ammonia, NH3, and related amines produce hydroxide ions when dissolved in water by accepting a proton from the solvent H2O. NH3 (g) + H2O (l) NH4 + (aq) + OH– (aq) Acids and bases are classified as either strong or weak. Strong acids or bases are those that are completely or almost completely ionized in dilute aqueous solution. We can easily calculate the concentration of either H3O+ or OH– ions in these solutions by assuming that the acid or base is completely dissociated into its constituent ions when it is dissolved in water. For example, 0.10 M hydrochloric acid is around 100% dissociated so the H3O+ concentration is 0.10 M. In aqueous solution of weak acids and bases, the undissociated species predominate and the concentration of H3O+ or OH– ions is small. For example, 0.10 M acetic acid is almost 98.7% undissociated so that the H3O+ concentration is 0.0013 M, or 1.3 × 10-3 M. CH3COOH (aq) + H2O (l) H3O+ (aq) + CH3COO– (aq) 98.7% unreacted Small amount Principles of Chemistry Lab II Montgomery College, Rockville 2 The pH Scale The pH scale was developed to express low concentrations of H3O+ without the inconvenience of using decimal numbers and negative powers. We define pH in terms of the molar concentration of H3O+ using Equation 1. pH = – log [H3O+ ] (1) The H3O+ concentration and the pH have an inverse logarithmic relationship; the higher the H3O+ concentration, the lower the pH. A change in pH of 1 unit results in a 10× change in [H3O+ ]. For example, a solution with pH of 5.00 is 100 times more acidic than a solution with a pH of 7.00. This means that a 0.1 M solution of a strong acid (completely ionized) such as HCl has a lower pH value (pH = 1.0) than a 0.01 M solution of the same acid (pH = 2.0). In aqueous solutions, there is also a relationship between the H3O+ concentration and the OH– concentration. To understand this relationship we must first understand something about the nature of water. Water “self-ionizes” to a very slight extent, as shown in the following equation. H2O (l) + H2O (l) H3O+ (aq) + OH– (aq) In this reaction one molecule of water acts as a Bronsted-Lowry acid (a proton donor) while the other acts as a Bronsted-Lowry base (a proton acceptor). In absolutely pure water the concentration of H3O+ and OH– are exactly the same. For every water molecule that dissociates, one hydronium and one hydroxide ion are formed. In pure water at 25 °C the concentration of each ion is 1.0 × 10–7 M and the pH of pure water is 7.0. Since the concentrations of H3O+ and OH– are equal, pure water is said to have a neutral pH. The product of these two concentrations is known as Kw, which is calculated in Equation 2. Kw = [H3O+ ][OH– ] = 1.0 × 10-14 (2) Equation 2 holds for most any aqueous solution. For aqueous solutions, if we know the concentration of H3O+ , we can calculate the concentration of OH– , and vice versa. We can now define acidic and basic solutions in a more quantitative manner. A solution is acidic if its H3O+ concentration is greater than 1.0 × 10-7 M or it is has a pH which is less than 7. In acidic solution, the H3O+ concentration is greater than OH– concentration. Conversely, a solution is basic if its H3O+ concentration is less than 1.0 × 10-7 M or it is has a pH which is greater than 7. In basic solutions, the OH– concentration is greater than the H3O+ concentration. Principles of Chemistry Lab II Montgomery College, Rockville 3 Indicators The pH of a solution can be measured precisely using a pH meter. Frequently, however, we need only an approximate pH value, which can be determined using a suitable acid-base indicator. Indicators are complex organic molecules which change color as they lose or gain a hydrogen ion. If we represent the acid form of the indicator as H:In, and the base form of the indicator as :In– , we can write its reaction with a base :B– or acid H:B as the following reversible equation. H:In (aq) + :B– (aq) :In– (aq) + H:B (aq) Acid form of indicator Base Base form of indicator acid Each indicator has its characteristic tendency to ionize; therefore, its color change will occur over a specific range of pH values. Table I shows some common indicators with their color changes and the pH ranges at which they occur. Table I. Common Acid-Base Indicators Indicator pH range of color change Color in acid Color in base Methyl violet 0 – 2.0 Yellow Violet Congo red 3.0 – 5.0 Blue Red Litmus 5.0 – 8.2 Red Blue Phenolphthalein 8.3 – 10.0 Colorless Red/pink Alizarin Yellow R 10.1 – 12.0 Yellow Red Hydrolysis Weak acids are poorly ionized in aqueous solution; that is, they are poor proton donors or BronstedLowry acids. Conversely, the bases formed when weak acids dissociate (their conjugate bases) have relatively strong affinity for protons and are comparatively strong Bronsted-Lowry bases. Consider the dissociation of H2CO3 below to produce the conjugate base HCO3 – . Correspondingly, the conjugate acids formed when weak bases react with H2O are relatively strong. For example, consider dissolving HCl in water. Principles of Chemistry Lab II Montgomery College, Rockville 4 Since HCl is a strong acid, Cl– , its conjugate base, is weak. Further, since we know HCl is strong and does dissociate completely, it is a stronger acid than H3O+ . Likewise, since NH4 + , the conjugate acid of NH3, does in fact react with OH– to produce NH3, NH4 + is a stronger acid than H2O. Similarly, H3O+ is a stronger acid than H2CO3. Table II lists some common acids and bases in order of their relative acid and base strengths. This table will aid you in your interpretation of the observations you will make in the hydrolysis experiments. Table II. Relative strengths of acids and their conjugate bases Acid Conjugate Base Strongest H2SO4 HSO4 – Weakest HCl Cl– H3O+ H2O HSO4 – SO4 2– H3PO4 H2PO4 – HC2H3O2 C2H3O2 – H2CO3 HCO3 – H2PO4 – HPO4 2– NH4 + NH3 H2O OH– HCO3 – CO3 2– Weakest HPO4 2– PO4 3– Strongest Principles of Chemistry Lab II Montgomery College, Rockville 5 Buffers The proper function of biochemical systems requires that the pH be maintained within a few tenths of a pH unit, yet many cellular reactions produce hydronium ions which could radically alter the pH. Blood normally has a pH of about 7.4; any significant variation can result in death. Many chemical reactions occurring in the laboratory or in industry also require careful control of pH. The pH of a system can be controlled by the use of a buffer. A buffer is a substance or combination of substances capable of consuming limited amounts of added H+ or OH– , thereby preventing significant changes in the pH of the system. Consider one of the major buffers in blood, the carbonic acid/bicarbonate ion system. Carbonic acid (aqueous CO2) is a weak acid that ionizes as follows. H2CO3 (aq) + H2O (l) H3O+ (aq) + HCO3 – (aq) If we prepare a solution containing both carbonic acid and bicarbonate ion (from NaHCO3) at equal concentrations, as shown below, we see that added OH– simply reacts with H2CO3. H2CO3 (aq) + OH– (l) H2O (aq) + HCO3 – (aq) Similarly, added H+ reacts with HCO3 – . HCO3 – (aq) + H+ (l) H2O (aq) + H2CO3 (aq) Acid or base added to the system is removed by reacting with one component of the buffer system, converting it into the other buffer component. Thus, as long as we do not exceed the buffering capacity (which depends on the concentration of the buffer components), addition of H+ or OH– to the buffer solution results in only very slight pH changes. Summary In this experiment we will • Observe the conductivity of a series of acids and bases as a measure of the degree of dissociation and strength of the acid or base • Use indicators and a pH meter to estimate the pH of acidic and basic solutions • Study the control of pH by means of buffers • Study pH changes that result from the hydrolysis of salts. Principles of Chemistry Lab II Montgomery College, Rockville 6 Pre-laboratory Questions (to be answered in order in blue or black in in the lab notebook; due at the beginning of lab) 1. You will work with 0.10 M acetic acid and 17 M acetic acid in this experiment. What is the relationship between concentration and ionization? Explain the reason for this relationship. 2. Explain hydrolysis, i.e, what types of molecules undergo hydrolysis (be specific) and show equations for reactions of acid, base, and salt hydrolysis not used as examples in the introduction to this experiment. 3. In Part C: Hydrolysis of Salts, you will calibrate the pH probe prior to testing the pH of the various salts. In a few sentences, summarize and explain the necessity of the calibration process. Why is it necessary to work with three buffers? 4. You prepare a buffer by adding 10.0 mL of 0.10 M acetic acid and 10.0 mL of 0.10 M sodium acetate. Calculate the pH that you expect for this buffer. Write the chemical equation for the equilibrium. Show your work using an ICE table. 5. To the solution from #4, you add 0.50 mL of 0.10 M HCl. Calculate the new pH of the buffer solution. Write the chemical equations for the reaction and the equilibrium. Show your work using an ICE table. 6. To the solution from #4, you add 0.50 mL of 0.10 M NaOH. Calculate the new pH of the buffer solution. Write the chemical equation for the equilibrium. Show your work using an ICE table. 7. Beginning on a new page, create data tables in your lab notebook using a straightedge to. Do not submit the tables with your pre-lab questions, as you will record your data in the tables during lab. Procedure You will be working in groups for this experiment. Someone in each group will be required to submit a photo I.D. in order to check out a Lab Quest unit and box containing the cables, a conductivity probe, and a pH probe. The lab Quest and other equipment is available from the prep room. Your I.D. will be returned when the Lab Quest unit and other equipment is returned to the prep room at the end of the lab session. Part A: Conductivity of Strong and Weak Acids and Bases Test the conductivity of each of the following solutions and record the results in your notebook. 0.10 M hydrochloric acid (HCl) 0.10 M sodium hydroxide (NaOH) 0.050 M sulfuric acid (H2SO4) 0.10 M ammonia (NH4OH) 0.10 M acetic acid (HC2H3O2) 0.10 M methylamine (CH3NH2) 17 M acetic acid (HC2H3O2) Deionized water 0.050 M oxalic acid (H2C2O4) Distilled water Tap water Principles of Chemistry Lab II Montgomery College, Rockville 7 Lab Quest Instructions 1. Connecting the equipment and Lab Quest together a. Connect the power cord to the Lab Quest unit (right side) and plug it into a power source. b. Connect the conductivity probe into one of the four channels (CH1, CH2, etc.) located on the left side of the Lab Quest unit. Make sure the conductivity probe is set to the 0-20000 µs position. c. Turn on the Lab Quest unit by pushing the power button on the top of the unit. d. In the meter mode, the screen should display a meter reading with the conductivity reading in µs/cm. A stylus can be used to enter information and move between different screens and commands. The stylus is connected via a cord to the Lab Quest unit. Do not use pens, etc. on the screen. At the top of the screen are a series of tabs consisting of, from left to right, Meter (File sensors), Graph (File Graph Analyze), Table (File Table), and Note Pad (File Notes). At the bottom of the screen are a series of icons consisting of, at far left, Data Collection (to start and stop the data collection process), and the Home, Battery status, and finally the time. Tapping on the appropriate tab will move you to that screen. 2. Collecting the data a. Pour about 20 mL of the solution to be tested into a 50 mL beaker. b. In the meter screen, tap on mode (upper right of the screen) and change the mode setting from “Time based” to “Events with Entry” (tap on the down arrow). For “Name”, type in solution, leave the “Units” area blank, and then select OK. c. To start the data collection, tap on the data collection icon (lower left of the screen). The screen should change to a graph. Tapping on the meter tap will change the screen back to displaying the conductivity meter reading. d. Collect the conductivity readings for the various solutions. Carefully raise the beaker with the solution around the conductivity probe until the hole near the end of the probe is completely submerged in the solution being tested. This is necessary because the two electrodes are positioned on either side of the hole. Briefly swirl the beaker contents. e. When the conductivity reading is fairly stable (note: the value will fluctuate a little bit) tap keep and then enter the name of the tested solution in the new window that appears. Then select OK. Note: set the switch on the conductivity probe to the 0-200 position when measuring the following solutions: deionized water, distilled water, and 17 M acetic acid f. Clean the electrodes by rinsing with deionized water. Carefully dry the outside of the probe with a Kimwipe tissue. It is not necessary to dry the inside of the hole. g. Repeat the procedure for each of the other solutions. When you have finished collecting data for all 11 solutions, tap the data collection icon to stop the process. h. Tap on the Table icon (upper right of the screen) to see the data table of conductivity readings and solutions tested. Copy down the conductivity data into your lab notebook. Principles of Chemistry Lab II Montgomery College, Rockville 8 Part B: Indicators 1. Indicator papers: Place approximately 1 mL of each of the following acids and bases (0.1 M solutions) into separate, labeled test tubes: hydrochloric acid, acetic acid, methylamine, and sodium hydroxide. Test the color change produced by each solution with each of the following indicator papers: Congo Red, red and blue litmus, and wide-range pH paper. This is conveniently done by placing strips of the test papers (a separate strip for each solution tested) on a sheet of white paper, and transferring one drop of solution to the test paper by means of a glass rod. Note the color change that occurs as soon as the solution wets the indicator paper. (Colors may change later due to air-borne contaminants.) Be sure to rinse the glass rod between each test. Record your observations in the data table in your notebook. After you have determined the color change of each indicator with the four solutions, add 1 drop of phenolphthalein solution to each test tube. Note any color change which occurs and record it in your notebook. 2. Universal Indicator: By combining several indicators of suitable pH ranges, a mixed or universal indicator can be prepared that changes colors at intervals across the entire pH range. The widerange indicator paper you used has been impregnated with a universal indicator. Prepare another set of labeled test tubes containing 1 mL of each of the same acids and bases you tested in Part B procedure 1. To each solution, add 1 drop of universal indicator. Note the colors which develop upon mixing and their estimated pH values and record them in your notebook. Part C: Hydrolysis of Salts The salts listed below will be tested. To prepare the solutions, pour about 10 mL of one of the salt solutions into a 50 mL beaker. If a prepared solution is not available, weight out approximately 1 g of the salt in a 50 mL beaker, then add 10 mL of deionized water. Stir the mixture until the solid has completely dissolved. NaC2H3O2 NaHCO3 NaH2PO4 NaCl NH4Cl Na2CO3 Na2HPO4 Lab Quest Instructions 1. Connecting the equipment and Lab Quest together a. Connect the pH probe into one of the four channels (CH1, CH2, etc.) located on the left side of the Lab Quest Unit. b. In the meter mode, the screen should display a meter reading with a pH value. A stylus can be used to enter information and move between different screens and commands. 2. Preparing the pH electrode for use a. Remove the storage bottle from the electrode by loosening the lid. Slide the lid upward on the probe body, exposing the O-ring. If present, carefully slide the O-ring up also. Do not remove the O-ring and the cap completely off the probe body. Rinse the lower section of the probe, especially the region of the bulb, with deionized water. If you notice air bubbles Principles of Chemistry Lab II Montgomery College, Rockville 9 in the reference reservoir, shake the electrode in a downward motion until all bubbles disappear, or else see your instructor. b. Clamp the pH probe around the bottle cap using a utility clamp. When the probe is not in use it can be stored for short periods of time (up to 24 hours) in a pH 4 or pH 7 buffer solution. It should never be stored in deionized water. For this experiment, use a pH 7 buffer solution as the electrode solution. 3. Calibration of the pH electrode a. In the meter mode, tap on Sensors, then tap on Calibrate and then on pH. In the new window, tap on Calibrate Now to do a 2-point calibration. b. Rinse the tip of the electrode with deionized water and catch the rinse water in a 100-mL beaker. Carefully blot dry the electrode with a Kimwipe. Do not touch the glass electrode; it is extremely fragile. c. Pour 10-20 mL of a pH 4 buffer solution into a 50 mL beaker. Raise the solution to the pH electrode, swirling it gently around the electrode. A voltage reading is displayed in the upper right of the screen. When the voltage reading is steady, type in a pH value of 4.0 for “Reading 1 known value”, and tap on “Keep.” Rinse the electrode with deionized water, and blot dry with a Kimwipe. d. Repeat step c) with a pH 10 buffer solution. When the voltage reading has stabilized (this will probably take a couple of minutes to reach a final voltage reading of ~1.0 V), type in a pH value of 10.0 for “Reading 2 known value”, tap on “Keep”, and then tap on OK. The electrode should now be calibrated. Rinse the electrode with deionized water, and blot dry with a Kimwipe. *The pH amplifier produces a voltage of 1.75 V in a pH 7 The results of this assessment are attached. buffer solution with an increase of ~0.25V/pH unit decrease and a decrease of ~0.25V/pH unit increase. e. Check the calibration with 10-20 mL of the pH 7 buffer solution in a 50 mL beaker. In meter mode, the pH reading should be within 0.5 pH units of 7.0. 4. Collecting the data a. In the meter mode, the screen will be displaying live pH readings. b. Rinse the pH electrode with deionized water. Blot it dry. Again, be careful since the pH probe is fragile. c. Raise the solution to be tested to the electrode and swirl gently. When the pH reading has stabilized, write down the pH value in the appropriate data table in your notebook. d. With the same procedure, determine the pH of the other solutions. Remember to rinse and blot dry the electrode between solutions. Also remember to copy the data displayed on the screen into your notebook. Part D: Buffers Use the Lab Quest instructions from Part C when taking pH measurements in Part D. 1. pH changes in aqueous solution Principles of Chemistry Lab II Montgomery College, Rockville 10 a. Using the pH probe, measure the pH of distilled water. b. Pour 10 mL of distilled water into a 50 mL beaker. Add 10 drops of 0.1 M HCl, mix well, and measure the pH. c. Pour 10 mL of distilled water into a second 50 mL beaker. Add 10 drops of 0.1 M NaOH, mix well, and measure the pH. d. Record the pH of each solution in the data tables for Part D. 2. pH changes in acetic acid/sodium acetate buffer a. Pour 10 mL of 0.1 M acetic acid in a 50 mL beaker and measure the pH. b. Pour 10 mL of 0.1 M sodium acetate in a second 50 mL beaker and measure the pH. c. Combine the acetic acid solution with the sodium acetate solution. Mix well and measure the pH of the resulting buffer solution. d. Divide the buffer solution into 2 equal portions of about 10 mL each. To one portion add 10 drops of 0.1 M HCl, mix well, and measure the pH. To the second portion add 10 drops of 0.1 M NaOH, mix well, and measure the pH. e. Record the data in the data tables for part D. 3. pH changes in dihydrogen phosphate/monophosphate buffer a. Pour 10 mL of 0.1 M sodium dihydrogen phosphate into a 50 mL beaker and measure the pH. b. Pour 10 mL of 0.1 M sodium monohydrogen phosphate into a second beaker and measure the pH. c. Combine the two solutions. Mix well and measure the pH of the resulting buffer solution. d. Divide the solution into 2 equal portions of about 10 mL each. To one portion add 10 drops of 0.1 M HCl, mix well, and measure the pH. To the second portion add 10 drops of 0.1 M NaOH, mix well, and measure the pH. e. Record the data in the data tables for Part D. Clean-up When you have finished the entire experiment, rinse the pH probe with deionized water, return it to the storage solution bottle, and screw the cap securely onto the bottle. Push and hold down the power button until the Lab Quest unit turns off. Disconnect the components and return them to the box. Return the Lab Quest unit and other equipment to the prep room and collect your I.D. Data (your data tables must be constructed in your lab notebook prior to lab) Record the following data tables in your laboratory notebook in blue or black ink. Principles of Chemistry Lab II Montgomery College, Rockville 11 Part A: Conductivity of Strong and Weak Acids and Bases Solution Conductivity Acid or Base? Weak or Strong? 0.10 M HCl 0.050 M H2SO4 0.10 M HC2H3O2 17 M HC2H3O2 0.050 M H2C2O4 0.10 M NaOH 0.10 M NH3 CH3NH2 Tap water Deionized water Distilled water Part B: Indicators Indicator HCl HC2H3O2 CH3NH2 NaOH Congo Red Red Litmus Blue Litmus Wide-range pH paper Phenolphthalein Universal Indicator Principles of Chemistry Lab II Montgomery College, Rockville 12 Part C: Hydrolysis of Salts Salt Solution pH value (probe) Universal indicator color and approximate pH NaCl NaC2H3O2 NH4Cl NaHCO3 Na2CO3 NaH2PO4 Na2HPO4 Part D: Buffers Solution pH Distilled water H2O + HCl H2O + NaOH HC2H3O2 NaC2H3O2 HC2H3O2 + NaC2H3O2 buffer Buffer + HCl Buffer + NaOH NaH2PO4 Na2HPO4 NaH2PO4 + Na2HPO4 buffer Buffer + HCl Buffer + NaOH Principles of Chemistry Lab II Montgomery College, Rockville 13 Acids and Bases, pH, Buffers, and Hydrolysis Lab Grade Sheet Pre-lab questions Questions are numbered, answered in order, and answered correctly in pen in the lab notebook Calculations include correct units and significant figures Work is shown for all calculations 25% Data Includes appropriate data tables written in the lab notebook in advance of lab Data recorded in blue or black ink 20% Post-lab Worksheet Any calculations include correct units and significant figures Work is shown for calculations References are cited according to ACS Guidelines 45% Student Comportment On time (not 1-5 minutes late), remembers drawer number and lock combination Prepared, understands procedure, works efficiently, and independently When necessary, works collaboratively and contributes equally to lab pair Works with high regard to personal safety and the safety of others Respects, maintains, and cares for lab glassware and equipment Thoroughly cleans personal and fair share of common glassware, equipment, and areas of lab Completes experimentation and clean-up by the end of lab 10% Total 100% Principles of Chemistry Lab II Montgomery College, Rockville 14 Principles of Chemistry Lab II Montgomery College, Rockville 15 Name: _________________________ Name of lab partner: _________________________ Date the experiment was conducted: _____________ ACIDS AND BASES, PH, BUFFERS AND HYDROLYSIS Post-Lab Worksheet Your work must be submitted on this worksheet (not other paper) and be legible. Acids and Bases 1. In your own words, explain the following acid/base theories and give an example of each. Acid Base Arrhenius: Explanation: Example: BrønstedLowry: Explanation: Example: Lewis: Explanation: Example: 2. In your own words, explain what makes an acid or base strong or weak. Conductivity 3. From your conductivity measurements, determine the strongest electrolyte. ___________________ Explain why it is the strongest electrolyte by using what you know about the properties of atoms, ions, and molecules. 4. From your conductivity measurements, determine the weakest electrolyte. ____________________ Explain why it is the weakest electrolyte by using what you know about the properties of atoms, ions, and molecules. Principles of Chemistry Lab II Montgomery College, Rockville 16 5. Acetic acid conductivity a. Which was more conductive (circle one)— 0.10 M acetic acid and 17 M acetic acid? b. Explain the chemistry behind the measurements you observed. Indicators 6. Describe a situation in which you might want to use Congo Red as an indicator instead of phenolphthalein. 7. Look up one of the indicators mentioned in this lab. a. Indicator name a. pH range b. chemical formula b. molecular structure b. Use the chemical formula and molecular structure to explain why the molecule acts as an acid/base indicator. Principles of Chemistry Lab II Montgomery College, Rockville 17 The pH Scale 7. Calculate the pH value for a sample of distilled, deionized water at 20 °C and 30 °C. Use your lecture notes or other reference materials to obtain any values needed. Cite references according to ACS Guidelines. pH at 20 °C pH at 30 °C 8. Is pure water acidic, basic, or neutral at 20 °C? _____________ At 30 °C? _____________ Explain. Salts 9. Explain why spectator ions such as sodium and chloride do not affect the pH of an aqueous salt solution. Principles of Chemistry Lab II Montgomery College, Rockville 18 Buffers 10. Results a. Describe your results for the acetate buffer solution and for the phosphate buffer solution. b. Use your results to explain how buffers work and why a buffer must contain a weak acid (or weak base) and its conjugate (or conjugate acid). Use equations to enhance your explanation. Continued Principles of Chemistry Lab II Montgomery College, Rockville 19 11. Determine whether each of the following systems has pH < 7, pH >7, or pH = 7. For each, complete the reaction and identify whether the reactants (R) or products (P) will predominate and explain your choice. All species are aqueous unless otherwise noted. pH 7 <, >, = Reactants Favored: R or P Explain your choice of R or P a. H2SO4 + H2O(ℓ) Complete the equation H2SO4 + H2O(ℓ) ⇌ b. NH3 + H2O(ℓ) Complete the equation NH3 + H2O(ℓ) ⇌ c. NH4 + + H2O(ℓ) Complete the equation NH4 + + H2O(ℓ) ⇌ d. NH4 + + CO3 2- Complete the equation NH4 + + CO3 2- ⇌ e. NH3 + HF Complete the equation NH3 + HF ⇌ f. benzoic acid and acetic acid Complete the equation C6H5COOH + CH3COOH ⇌ g. hypochlorous acid and methylamine Complete the equation HClO + CH3NH2 ⇌ Principles of Chemistry Lab II Montgomery College, Rockville 20