Experiment 23   Acids, Bases, and Salts

 

 

Discussion

Aqueous solutions, which are so common, almost always are acidic or basic to some degree.  In fact, large amounts of dissolved salts influence or determine the pH of these solutions.  (The pH scale is used as a measure of the acidity of a solution.)  Photosynthesis and respiration, the two most important biological processes on earth, depend on acid-base reactions.  Carbon dioxide, CO2, is the most important acid-producing compound in our atmosphere; rainwater is therefore naturally slightly acidic because of the dissolved carbon dioxide.  Further acidification may result form pollutants, such as NO2 and SO2.  On the other hand, our oceans are slightly basic because of dissolved salts such as carbonates and phosphates. 

 

The main components of the digestive fluid in the human stomach are pepsinogen (an inactive form of the enzyme pepsin) and hydrochloric acid, HCl.  The HCl promotes the conversion of pepsinogen into its active form and provides the necessary environment in which the enzyme pepsin can break down protein molecules.  When excessive digestive fluid (gastric juice) is secreted it may contribute to the formation of an ulcer in the stomach lining.  One of the most common remedies for excessive digestive fluid (stomach acidity) is antacid tablets.  Stomach secretions generally have a pH ranging from 1.0 to 2.0; acid indigestion occurs at a lower pH.  Antacids neutralize (or buffer) the excess hydronium ion, H3O+, in the stomach to relieve this discomfort.

 

The terms acidic and basic are used to describe solutions with varying concentrations of hydrogen ions.  A hydrogen ion is commonly written as H+ even though hydrogen ion doesn’t actually exist in solution.  Rather, it exists covalently attached to water to produce H3O+, the hydronium ion.  Even though it is more precise to describe acidic and basic solutions in terms of their hydronium ion concentration, we will use the hydrogen ion for simplicity. 

 

 

The pH concept

According to the Brønsted-Lowry concept, an acid is a substance that donates an H+ and a base is a substance that accepts an H+ in an aqueous solution.  The concentration of the hydrogen ion [H+] is then a measure of the acidity of a solution and the pH of a solution is defined by the following equation:

 

                                    pH = - log [H+]                                                                                      (1)

 

The use of pH has the advantage that one does not deal with small fractions or very small numbers in the molarity expression.  For example, a solution with a molar concentration of H+ = 1 x 10-7 has a pH of 7.  Natural rainfall has a pH of 5.6.  This is due to carbon dioxide which, in an aqueous solution, exists in equilibrium with carbonic acid.  Rainwater with a pH of 2.6 has been observed in parts of Europe.  A pH of 2.6 is 1,000 times more acidic than a pH of 5.6!

Basic solutions are also expressed in terms of pH, or in terms of the pOH.  The pOH is defined in much the same way as the pH and can be written as

 

                                    pOH = - log [OH-]                                                                                             (2)

 

In any aqueous solution, the following equilibrium relation will always be obeyed:

 

                        [H+] x [OH-]= Kw = 1 x 10-14 at 25 oC                                                  (3)

or,

                        pH + pOH = 14 at  25 °C                                                                                  (4)

 

Kw is the water ionization constant and has a value of 1 x 10-14 at 25 °C.  In deionized water, [H+] always equals [OH-].  So, by equation (2), [H+] must be 1 x 10-7 M.  Therefore, ideally, the pH of deionized water is 7.  Aqueous solutions with a pH of 7 are said to be neutral. 

                        neutral solution                      [H+] = 10-7 M    pH = 7

                        acidic solution                        [H+] > 10-7 M    pH < 7

                        basic solution             [H+] < 10-7 M    pH > 7

 

The common pH scale runs from 1 to 14.  Like a Richter scale for measuring earthquakes, each unit of increase or decrease on the pH scale corresponds to a concentration change equal to a factor of 10. 

 

 

 

The following table shows pH values and corresponding hydrogen ion and hydroxide ion concentrations of everyday acids and bases. 

pH

[H+]

[OH]

Acid-Base Rating

Common Example

0

10-0 M

10-14 M

Acidic

Swimming pool acid

1

10-1 M

10-13 M

Acidic

Battery acid

2

10-2 M

10-12 M

Acidic

Stomach acid, lemon juice

3

10-3 M

10-11 M

Acidic

Vinegar, soft drinks

4

10-4 M

10-10 M

Acidic

Apple juice, acid rain

5

10-5 M

10-9 M

Acidic

Urine

6

10-6 M

10-8 M

Acidic

Unpolluted rainwater

7

10-7 M

10-7 M

Neutral

Pure water

8

10-8 M

10-6 M

Basic

Fresh egg white, bile

9

10-9 M

10-5 M

Basic

Baking soda solution

10

10-10 M

10-4 M

Basic

Milk of magnesia

11

10-11 M

10-3 M

Basic

Ammonia solution

12

10-12 M

10-2 M

Basic

Washing (laundry) soda

13

10-13 M

10-1 M

Basic

Dilute lye solution

14

10-14 M

10-0 M

Basic

Strong lye solution

 

 

The pH can be measured in several ways.  Most easily and precisely, it is done with a pH meter.  A pH meter is a device that uses a glass electrode sensitive to changes in [H+] .  The potential between the electrodes is related to the pH and is converted into a pH reading. 

Other methods include the use of pH paper and various indicators.  pH paper and indicator solutions contain chemicals that change colors in response to different [H+] concentrations.  By comparing the color to a chart, the pH can be determined.

 

 

 

 

 

 

 


 

Acid and Base Behavior in Aqueous Solution

Some acids will undergo almost complete ionization in water.  These acids are called strong acids.  In the same way, some bases will undergo substantial dissociation in water.  These bases are called strong bases.  We can represent the dissolving and ionization of a strong acid, such as HCl in water with an equation:

 

                        HCl(aq)  +  H2O(l)  ®  H3O+(aq)  +  Cl-(aq)                                           (5)

 

The dissociation of a strong base such as NaOH is represented by the equation

 

                        NaOH(aq)  ®  Na+(aq)  +  OH-(aq)                                                                   (6)

 

Other acids and bases, because of incomplete ionization, are called weak.  Acetic acid, CH3CO2H, is a common weak acid, and ammonia, NH3, is a common weak base. 

The equations for the ionization reactions of these two substances are given by

 

            CH3COOH(aq)  +  H2O(l)  Û  H3O+(aq)  +  CH3COO-(aq)                                (7a)

 

            NH3(aq)  +  H2O(l)  Û  NH4+(aq)  +  OH-(aq)                                                       (7b)

 

Note that a double arrow, Û, is used to indicate that this is an equilibrium reaction. 

At equilibrium, the reaction of the acid can be expressed in terms of an equilibrium constant and is written as

 

                                    Ka[H3O+] x [CH3COO-]                                                                   (8)

                                                    [CH3COOH]

 

where Ka is called the acid ionization constant, and has a constant value at a given temperature.   It follows that the larger the ionization constant, the stronger the acid. 

 

A similar equation can be derived for bases.  The corresponding equilibrium constant expression is

 

                                    Kb = [NH4+] x [OH-]                                                                              (9)

                                                    [NH3]

 

The equilibrium constant, Kb, is called the base ionization constant for the base.  The larger the ionization constant, the more product-favored the ionization reaction, and the stronger the base. 

 

 

In addition to Ka, another useful measure of the strength of a weak acid is the percent ionization, defined as the concentration of the acid that ionizes divided by the initial concentration of the acid times 100%.  For a weak acid, HA, it is calculated

 

            Percent Ionization = [HA] ionized      x 100%

                                                  [HA] initial

 

In general, the percent ionization depends on the acid and increases with increasing value of Ka.  For a given weak acid, the percent ionization also increases with increasing dilution.  For example, a 0.0100 M CH3COOH is 4.2% ionized, whereas a 1.00 M CH3COOH is only to 0.42% ionized. 

 

 

 

Salt Behavior in Aqueous Solution

A salt is one of the products of a chemical reaction between an acid and a base (called a neutralization reaction). 

 

                        KOH(aq)  +  HCl(aq)     ®         KCl(aq)  +  H2O(l)                                             (10)

 

A salt that is formed from a strong acid and a strong base, such as in reaction (9), will dissolve in water to produce a solution of neutral pH. 

 

A salt derived from a weak acid and a strong base will produce a basic solution and will give a pH different from the pH of the original water.  This is because the salt reacts with water in a reaction known as hydrolysis.  For example, sodium acetate is an example of a basic salt.  It is prepared in a reaction similar to (10) by the reaction of sodium hydroxide with acetic acid (11).  Its hydrolysis produces a basic solution, as shown in equation 12. 

 

            CH3COOH(aq)  +  NaOH(aq)     ®         NaCH3COO(aq)  +  H2O(l)                     (11)

 

            CH3COO-(aq)  +  H2O(l)                        Û         CH3COOH(aq)  +  OH-(aq)                    (12)

 

 

In a similar manner, salts derived from strong acids and weak bases hydrolyze in water to produce acidic solutions.  Ammonium chloride, NH4Cl, is an example of an acidic salt.  Its hydrolysis is shown in equation (13).  In an aqueous solution, ammonium chloride dissociates into its ions producing NH4+ and Cl-.  The ammonium ion will hydrolyze as shown in equation (13) and the resulting solution will be slightly acidic.  The Cl- is derived from a strong acid (HCl) and will not hydrolyze.  Thus, the net result of the hydrolysis reaction will be an acidic solution. 

            NH4+(aq)  +  H2O(l)      Û         H3O+(aq)  +  NH3(aq)                                       (13)

 

 

Materials and Equipment

pH meter                                            

Dropper                      

Acids:  0.10 M HCl, 0.10 M acetic acid, CH3COOH, 0.10 M benzoic acid, C7H5O2H

Bases:  0.10 M NH3, 0.10 M NaOH, 0.10 M Ba(OH)2

Salts:  0.10 M solutions of each KBr, NH4Cl, Na3PO4, K2SO4, and NaHCO3

Indicators:  phenolphthalein and methyl orange

 

Check also:

http://www.funsci.com/fun3_en/acids/acids.htm

http://www.members.aol.com/logan20/ionic_eq.html

 

 

 


 

Procedure

Use a pH meter to find the pH of all solutions.  The electrodes may be fragile, so be very careful when handling the electrode.  Rinse the electrode well with deionized water after each measurement and when done put it back into its solution.  Never leave it out. 

 

In your notebook: Record all observation directly into your notebook.    

 

I. pH, [H+], and Acid Strength

1.  Determine the pH of 0.10 M HCl.  Measure about 5 mL of HCl into a large test tube.  Immerse the pH electrode completely.  Determine and record the pH and the hydronium ion concentration that corresponds to this pH value. 

2.  Perform a 10-fold dilution, starting with 0.10 M HCl.  In a same way as described under (1), determine and record the pH of this solution and the corresponding [H+].

3.  Repeat this procedure with 0.10 M benzoic acid, C7H5O2H, and then with 0.10 M CH3COOH.  For CH3COOH only, perform a second 10-fold solution starting with 0.010 M, then measure the pH. 

4.  Arrange the three acids in order of increasing acid strength.  Which one is the strongest, which one is the weakest acid?  Look up the Ka of each acid.  Is your order in agreement with the listed Ka?  Please explain. 

5.  Now calculate the % ionization of 0.10 M CH3COOH and of 0.0010 M CH3COOH.  Compare your results to you observations. 

 

 

II. pH, [OH-], and Base Strength

 

1.  Determine the pH of 0.10 M NaOH.  Measure about 5 mL of NaOH into a large test tube.  Immerse the pH electrode completely.  Determine and record the pH and the hydronium ion concentration that corresponds to this pH value. 

2.  Perform a 10-fold dilution, starting with 0.10 M NaOH.  In a same way as described under (1), determine and record the pH of this solution and the corresponding [OH-].

3.  Repeat this procedure with 0.10 M Ba(OH)2, and then with 0.10 M NH3.  For NH3, perform a second 10-fold solution starting with 0.010M NH3

4.  Arrange the three bases in order of increasing base strength.  Which one is the strongest, which one is the weakest base?  Look up the Kb of each base.  Is your order in agreement with the listed Kb?  Please explain. 

 

 

III. pH of Solutions of Acid and Base Mixtures

1.  In a test tube, combine 2.5 mL 0.10 M of each hydrochloric acid and sodium hydroxide.  Measure the pH. 

2.  In second test tube, combine 2.5 mL 0.10 M of each benzoic acid and sodium hydroxide.  Measure the pH. 

3.  In a third test tube, combine 2.5 mL 0.10 M of each acetic acid and sodium hydroxide.  Again, measure the pH. 

 

 

IV. pH of Some Common Solutions

Use a pH meter to determine the pH of each of the following 0.10 M solutions. 

            KBr                  NH4Cl              Na3PO4                       K2SO4 NaHCO3

 

1.  Measure about 5 mL into a small test tube.  Carefully immerse the pH electrode into the solution.  Measure the pH of each solution and record in on your Data Sheet. 

2.  Write net ionic equations for all solutions that explain why the pH you measured is the correct pH. 

 

 

V. Preparation of a Salt Solution of predetermined pH

In this part of the experiment you are asked to prepare a salt solution of predetermined pH.  You will need to use one of the salts from part IV of this experiment, at a particular concentration.  You will need to calculate the concentration necessary to produce this pH, get your Instructor’s approval, and then verify its pH through measurement.  Be sure to calibrate the pH meter for a more accurate reading.

 

 

Waste Disposal 

All solutions should be poured into the proper waste container located in the hood. 

 

 

Prelaboratory Questions: Acids, Bases, and Salts

 

1.  Define the following terms by means of an equation, or a definition. 

a)  Brønsted-Lowry acid

b)  Brønsted-Lowry base

2.  In the opening statement of this lab it is said that “Carbon dioxide, CO2, is the most important acid-producing compound in our atmosphere; rainwater is therefor naturally slightly acidic because of the dissolved carbon dioxide”. 

a)  Show by means of an equation why carbon dioxide produces an acidic solution. 

b)  SO2 and NO2 are also acidic anhydrides.  Write equations showing this behavior. 

3.  Predict the pH of the following 0.10 M solutions (neutral, acidic, or basic)

NaCl, KCN, NH4Br, NaHSO3             

4.  Identify the acid and base used in the preparation of the following salts. 

KCN, SrBr2                 

5.  You are asked to prepare a solution of a particular salt that has a pH 8.3 (your choices are KCN, NH4Cl, KNO2, or BaCl2.).  What molarity of which salt should you choose?  (Hint:  look up the ionization constant)

6.  The pH of a 0.100 M solution of hydrazoic acid, HN3, is 2.75.  What is the value of the acid-ionization constant of HN3

 

 

 

 

Data and AnalysisAcids, Bases, and Salts

 

For each of the measurements that you perform in this lab, prepare a table that contains the following information: acid type, concentration (M), observed pH, calculated pH, and [H3O+] concentration from your observed pH.  Record your observations. 

 

I. pH, [H+], and Acid Strength

1.  Arrange the above acids in order of increasing acid strength (weakest first)

2.  Explain any deviation between the measured and calculated pH.

3.  Calculate the % ionization of the 0.10M, 0.010M, and 0.0010M CH3COOH solutions. 

 

II. pH, [OH-], and Base Strength

1.  Prepare a table like the one above for the three different bases.  Measure the pH for each one. 

2.  Arrange the above acids in order of increasing base strength (weakest first)

 

III. pH of Acid-Base Solutions

1.  Calculate the pH for each acid-base reaction.  Then measure the pH and comment on any discrepancies that you observe. 

 

IV. pH of 0.10 M Solutions

Measure the pH of each salt solution.  Comment on you results and write net ionic equations. 

 

V.  Preparation of a Salt Solution of predetermined pH

  1. Indicate the salt and pH assigned by your instructor.
  2. Show calculations performed to calculate salt concentration needed.
  3. Report all data collected.
  4. Compare experimentally obtained pH with theoretical pH.
  5. Comment on any discrepancies.

 

 

 

Questions

1.  a) Calculate the pH of the 0.10 M NH4Cl solution.   b) Calculate the pH of the 0.10 M Na3PO4 solution.  c) Are the numbers you have calculated in agreement with the ones you have measured?

2.  Distinguish between a strong and a weak acid. 

3.  How is the strength of an acid related to the strength of its conjugate base?  Please use an example. 

4.  Using the pH that you measured for 0.10 M CH3COOH, calculate the Ka for this acid.