Acids, Bases, Buffers and pH

Introduction

Acids and Bases Arrhenius proposed the first theory of acids and bases. According to Arrhenius, pure water dissociates to some extent to produce hydrogen ions (H+) and hydroxide ions (OH-). When this occurs, equal amounts of H+ and OH- ions are produced:

H2O(l) H+(aq) + OH-(aq) (1)

It is now known that the hydrogen ion does not exist as such in water; rather, it is bonded to water and exists as the hydronium ion (H30+). Therefore the ionization of water may be better represented by equation (2):

2 H2O(l) H3O+(aq) + OH-(aq) (2)

An Arrhenius acid is a substance that increases the hydronium concentration in water. This causes the concentration of H30+ ions to be higher than that of OH- ions, and the solution is said to be acidic. The following shows the ionization of nitric acid, HN03, in water, and its behavior as an Arrhenius acid:

HNO3(aq) + H2O(l) H3O+(aq) + NO3-(aq) (3)

Similarly, an Arrhenius base is a substance that increases the hydroxide ion concentration when dissolved in water. The resulting solution has a higher concentration of OH- ions than H3O+ ions and is said to be basic or alkaline. Sodium hydroxide is a base. It ionizes to form hydroxide ion when dissolved in water:

NaOH(s) Na+(aq) + OH-(aq) (4) Both acids and bases can be classified as either strong or weak, terms that describe the extent to which they produce either H3O+ or OH- ions, respectively, in solution. A strong acid is one that is essentially 100% dissociated when dissolved in water. Hydrochloric acid, HCI, and nitric acid, HNO3, are strong acids, and both will produce an H3O+ concentration equal to the concentration of the acid. For example, a 2.0 M HNO3 will dissociate to 2.0 M H3O+. There are six strong acids: HCl, HBr, HI, HNO3, H2SO4 and HClO4. A weak acid is one that undergoes only

Purpose: To examine properties of acids and bases, buffers, and salt solutions, and use different methods to measure

the pH of those solutions.

partial dissociation in water. Any acid that is not one of the six strong acids must be a weak acid. Acetic acid, CH3COOH, is an example of a weak acid and will produce an H30+ concentration that is much lower than the concentration of the acid. Only a small fraction of the molecules of a weak acid are ionized in solution. Most of the acid is present as molecules. For example, in a 2.0 M acetic acid solution, only 13 of every 1000 acid molecules are ionized (1.3% dissociation).

CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq) (5) The double arrow indicates that a state of dynamic chemical equilibrium exists. Here the two opposing reactions occur simultaneously at the same rate. At equilibrium the concentrations of the reactants and products have stabilized at some constant value and it appears that the reaction has ceased. The state of equilibrium is described by an equilibrium constant, K. The general expression for K is given in equation (6):

[product]

K = [reactant]

The brackets, [ ], denote molar concentration. (6)

The equilibrim constant for the dissociation of an acid is called the acid dissociation constant, Ka. For example Ka for acetic acid is given by the following equation:

- +

3 3a

3

[CH COO ][H O ]K =

[CH COOH] (7)

Because there is always a very large quantity water (the solvent), the concentration of H2O is essentially a constant and is therefore not included in the equilibrium expression. The higher the value of Ka , the higher is the degree of dissociation of the acid. Thus an acid with a Ka value of 2.6 x 10-2 is stronger than an acid with a Ka of 8.5 x 10-8.

Sodium hydroxide, NaOH, and calcium hydroxide, Ca(OH)2, are strong bases. In solution, these compounds dissociate completely into ions. All group 1 hydroxides and Group 2 hydroxides beginning with calcium and continuing down the periodic table are considered strong bases. Aqueous ammonia, NH3(aq) (or ammonium hydroxide, NH4OH), and methylamine, CH3NH2, are weak bases. In solution, they produce a hydroxide ion concentration that is very small compared to that of the base.

NH3(aq) + H2O(l) NH4

+(aq) + OH-(aq) (8)

The strength of a base is expressed in terms of a base dissociation constant, Kb. For example Kb for ammonia is given in equation (9):

+ -

4b

3

[NH ][OH ]K =

[NH ] (9)

A base with a Kb value of 9.2 x 10-3 is stronger than a base with a Kb of 4.3 x 10-6.

Bronsted and Lowry proposed a more general theory of acids and bases. A Bronsted-Lowry acid is a substance that donates H+ (protons) and is called a proton donor, while a base is a proton acceptor. Notice with this definition, no water is needed, and so formation of H3O+ or OH- is not required. The H+ can be donated to any other molecule (base) that can accept it. It is also useful to consider the relationship between related pairs of structures with either one less or one more H+. For example, consider CH3COOH/ CH3COO- . These are examples of conjugate acid-base pairs. For any conjugate acid-base pair, the acid has one more proton than its conjugate base. Any Bronsted-Lowry acid-base reaction involves 2 conjugate acid-base pairs, as illustrated in the examples below:

CH3COOH(aq) + OH-(aq) CH3COO-(aq) + H2O(l) (10)

acid base conj. base conj. acid

NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) (11)

base acid conj. acid conj. base

Two general statements apply to acid-base equilibria:

• The stronger an acid, the weaker is its conjugate base. Also, the stronger a base, the

weaker is its conjugate acid.

• Equilibrium favors the formation of the weaker acid and the weaker base. HNO3 is a strong acid and gives up protons readily; this means that its conjugate base, NO3

-, is not going to accept protons readily. Thus, NO3

- is a weak Bronsted base. On the other hand, CH3COOH is a weak acid and does not readily give up protons; its conjugate base, CH3COO-, has a tendency to accept protons and is therefore a significant Bronsted base. In general, anions of weak acids are relatively stronger bases, and cations of weak bases are relatively stronger acids. The relative strength of acids and bases can be used to predict whether the products or reactants will be favored (present in larger amounts) at equilibrium. Consider the following reaction. CH3COOH + NO3

- CH3COO- + HNO3 (12) Compare the acids on each side. At equilibrium, we can see that the reactants are present in a greater amount, as the reaction will NOT proceed as written to form the stronger acid, (HNO3).

pH Scale

The pH of a solution is a measure of its hydronium ion concentration. The pH scale was developed to express low concentrations of H3O+ ion without the inconvenience of negative exponents. The defining equation of pH is:

pH = -log[H3O+] (13)

In a similar manner, a measure of the concentration of the hydroxide ion can be expressed as the pOH of the solution:

pOH = - log [OH-] (14) In any aqueous solution, the product of the hydronium ion concentration and the hydroxide ion concentration equals a constant, called the ion product constant, Kw:

K w = [H30+][OH-] = 1.0 X 10-14 at 25 °C (15)

Because of this relationship between the hydronium and hydroxide ion concentrations in aqueous solutions, the pH and pOH of a solution are also related:

pH + pOH = 14 (16) If the pH of a solution is 5, the pOH must equal 9. If the pH of a solution increases, the pOH decreases, and vice versa. If [H3O+] is greater than [OH-], the solution is considered to be acidic; acidic solutions have a pH less than 7. On the other hand, if [OH-] is greater than [H3O+], the solution is basic; basic solutions have a pH greater than 7. A pH of exactly 7 represents a neutral solution, one in which the [H3O+] = [OH-] = 1.0 X 10-7 M.

There are two common ways to measure the pH of a solution. The easiest method involves using wide-range pH paper (a paper saturated with indicator dyes that change color in response to the pH of a solution ) or universal indicator solution. A more accurate and precise pH value requires the use of a pH meter. This is an electronic device with a special electrode that generates a small voltage proportional to the hydronium ion concentration of the solution in which it is placed. The electrical signal is amplified, converted to pH, and displayed. Although pH paper is less sensitive than the pH meter, it is satisfactory in many applications.

Hydrolysis A salt, to a chemist, is any ionic compound. When a salt is dissolved in water, it dissociates into its component ions. Some ions react with water to a small degree to produce solutions that are either acidic or basic. Any chemical reaction in which water is a reactant is called a hydrolysis reaction. Therefore, the pH of a salt can be neutral, acidic, or basic depending upon how the ions react with water.

Example 1: When ammonium chloride (NH4Cl) is dissolved in water, the ammonium ion undergoes hydrolysis and the solution becomes slightly acidic. The chloride ion (the conjugate base of a strong acid) does not undergo hydrolysis.

NH4+ + H2O H3O+ + NH3 Yes! This reaction can occur because no strong acid or base

formed.

Cl-(aq) + H2O(l) HCl + OH- No! No reaction because strong acid (HCl) won’t form.

So overall (above), the pH will go down (becomes acidic) when NH4Cl is added to water because of the H3O+ that forms in the first chemical equation above.

Example 2: When sodium acetate (NaCH3CO2) is dissolved in water, only the acetate ion undergoes hydrolysis and the solution becomes slightly basic . The sodium ion does not react with water.

Na+ + H2O NaOH + H+ No! No reaction because strong base (NaOH) won’t form.

CH3COO- + H2O CH3COOH + OH- Yes! This reaction occurs because no strong acid or base formed.

So overall (above), the pH will go up (becomes basic) when sodium acetate is added to water because of the OH- that forms in the second chemical equation shown above.

To determine the effect of a salt, consider each equation and decide on the overall effect as done in the two examples above. In some cases, there will be no reaction for either the cation or the anion (and, therefore, no change in pH). In other cases, both reactions will occur, which means more detailed information would be needed to see which equation dominates (has a larger equilibrium constant).

Buffers

A buffer is a solution that resists large changes in pH when small amounts of acid or base are added. A buffer solution consists of either,

1. a weak acid and the conjugate base of that acid, for example, CH3COOH and CH3COO-(from the salt CH3COONa)—this will generally provide an acidic pH, or

2. a weak base and the conjugate acid of that base, for example, NH3 and NH4+ (from the salt

NH4CI)—this will generally provide a basic pH.

Consider first the acetic acid-sodium acetate buffer. The equilibrium that controls the pH is,

CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq) (17)

If a small amount of acid is added to this buffer solution, it combines with the CH3COO-, forming the compounds shown on the right. The pH remains relatively unchanged as the acid is consumed by the conjugate base, CH3COO-.

CH3COO -(aq) + H3O+(aq) CH3COOH + H2O(l) (18)

If a small amount of base is added to the buffer, it is consumed by the acid to form water and acetate ion. Again, this means the pH remains relatively unchanged as the base is consumed by the weak acid.

CH3COOH(aq) + OH-(aq) CH3COO -(aq) + H20(l) (19)

Consider now the ammonia-ammonium chloride buffer which is used to control the pH of solutions in the mildly basic range of about 8 to 10. The equilibrium that controls the pH is that described by the following equation.

NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) (20)

If a small amount of acid is added to this buffer solution, it combines with ammonia, forming ammonium ion and consuming the acid in the process.

NH3(aq) + H3O+(aq) NH4+(aq) + H2O(l) (21)

If a small amount of base is added, it is consumed by the conjugate acid to form water and ammonia.

N H 4+ (aq) + OH-(aq) NH3(aq) + H2O(l) (22)

These equations show how this buffer system can neutralize both acid and base.

The buffering capacity of a buffer (the amount of acid or base that it can neutralize without a significant change in pH) is limited by the nature and concentration of its components. Higher the concentration of reagents results in a higher buffer capacity. The pH of blood is maintained in the normal range (7.35-7.45) by the buffering action of three buffers: the bicarbonate buffer( H2CO3/HCO3

- ), the phosphate buffer (H2P04-/HP04

-), and the

plasma proteins. The buffer that plays the major role here is the bicarbonate buffer. Acidosis and alkalosis result when the blood pH drops below 7.35 or rises above 7.45 respectively.

Calculations for buffers systems use the Henderson-Hasselbach equation. We showed before that for a weak acid HA + H2O A- + H3O+

Ka =[A-][H3O+]

[HA]

Rearranging this we get [H3O+] =Ka[HA]

[A-]

Taking the negative log of both sides gives

pH = pKa + log[A-]

[HA]

the Henderson-Hasselbach equation (23)

Procedure

SAFETY NOTE

Handle hydrochloric acid (HCl) with care. Avoid breathing its vapors and avoid skin contact. If you do spill some on yourself, wash immediately with cold water. Notify your instructor.

Handle sodium hydroxide (NaOH) with care. Avoid skin contact. If you do spill some on yourself, wash immediately with cold water. Notify your instructor.

! Use caution with all acid and base reagents.

•Note: You will need freshly boiled deionized water for portions of this experiment. Pour about 80 mL of deionized water into a clean 100 mL beaker and boil for about 5 minutes to remove any dissolved CO2, which affects the pH of water. Set this aside to cool.

• Use of the pH Meter

Follow the lab instructor's directions for the use of the pH meter. Handle the meter with great care as the electrode is very fragile. Calibrate using a one point calibration before taking any pH measurements. Before each measurement, rinse the electrode with deionized water and gently wipe it dry with a Kimwipe. At the end of the experiment, rinse the electrode thoroughly, dry it with a Kimwipe, switch off the power, and cap.

• Use of Wide-Range pH Paper and Universal Indicator Solution

To use pH paper, tear the strips of paper into pieces about 2 cm long and lay them on a clean glass plate or a sheet of white paper. Then, to test the pH of a solution, obtain a droplet of the solution on the end of a clean stirring rod or Pasteur pipet and transfer it to a piece of pH paper. Compare the color produced with the color chart on the pH paper vial. Always rinse the stirring rod or pipet with deionized water before testing the next sample.

To use universal indicator solution, transfer about 1 mL of the solution to be tested to a test tube, add 1 or 2 drops of universal indicator, and compare the color produced to the colors of the standards for Universal Indicator.

• Waste disposal: all the test solutions may be flushed down the drain.

Begin by boiling water (see previous page) and then cooling to room temperature for

future use.

Part A. Measuring the pH of Common Acids and Bases

1. Pour about 25 mL of the following solutions into clean 30 mL plastic beakers. Measure

the pH of each solution with the pH meter. Record the pH values. 2. Repeat the pH measurement of these same solutions using pH paper and universal

indicator solution. Record the colors produced and pH values. 3. Calculate the hydronium ion concentration for each solution using the pH values from the

pH meter.

a. 0.1 M NH3 b. 0.1 M NaOH c. 0.1 M H3BO3 d. 0.1 MHCl

Part B. Hydrolysis: Measuring the pH of Salt Solutions

1. Pour about 25 mL of the following solutions into clean 30 mL plastic beakers. Measure

the pH of each solution with the pH meter. Record the pH values.

2. Calculate the hydronium ion concentration for each solution using the pH values from the

pH meter.

a. 0.1 M Na2CO3 b. 0.1 M (NH4)2SO4 c. 0.1 M NaCl

Part C. Buffer Action

Use only the pH meter for the following pH measurements.

1. Measure 25.0 mL of 0.1M CH3COOH using a graduated cylinder, and transfer the solution to a clean 50 mL beaker. Measure the pH.

2. Measure 25.0 mL of 0.1M CH3COONa using a graduated cylinder, and transfer the solution to a clean 50 mL beaker. Measure the pH.

3. Mix the two solutions and measure the pH of the resulting buffer. Record the pH value. Divide the buffer solution into 2 portions in 30 mL plastic beakers.

4. To one portion of the buffer (25 mL), add 5 drops of 1.0 M HCl one drop at a time. Swirl the solution after each addition, and measure the pH.

5. To the other portion of the buffer, add 5 drops of 1.0 M NaOH, one drop at a time. Swirl the solution after each addition, and measure the pH.

6. Repeat steps #4 and 5, using 25 mL of freshly boiled deionized water instead of the

buffer solution.

7. Calculate the volume of 0.1M CH3COOH and 0.1M CH3COONa needed to prepare a

buffer with a pH of 4.0. Make the buffer and measure the pH.

Acids, Bases, Buffers and pH

Data and Calculations Sheet

Name ___________________________

Part A. Measuring the pH of Common Acids and Bases

Solution pH Meter pH paper Universal Indicator

Reading [H3O+] color pH value color pH value

a. 0.1 M NH3

b. 0.1 M NaOH

c. 0.1 M H3BO3

d. 0.1 M HCI

Part B. Hydrolysis: Measuring the pH of Salt Solutions

Solution pH

meter reading [H3O+]

a. 0.1 M Na2CO3

b. 0.1 M (NH4)2SO4

c. 0.1 M NaCl

Part C. Buffer Action Solution pH Meter

Reading

a. 0.1M CH3COOH

b. 0.1M CH3COONa

c. CH3COOH/CH3COONa buffer

Part C. (continued)

Solution pH Meter Solution pH Meter

buffer Reading

buffer Reading

+ 1 drop of 1M HCl

+ 1 drop of 1M NaOH

+ 2 drops of 1M HCl

+ 2 drops of 1M NaOH

+ 3 drops of 1M HCl

+ 3 drops of 1M NaOH

+ 4 drops of 1M HCl

+ 4 drops of 1M NaOH

+ 5 drops of 1M HCl

+ 5 drops of 1M NaOH

Solution pH

Meter Solution pH Meter

Boiled (but cooled) H2O Reading Boiled (but cooled) H2O Reading

+ 1 drop of 1M HCl

+ 1 drop of 1M NaOH

+ 2 drops of 1M HCl

+ 2 drops of 1M NaOH

+ 3 drops of 1M HCl

+ 3 drops of 1M NaOH

+ 4 drops of 1M HCl

+ 4 drops of 1M NaOH

+ 5 drops of 1M HCl

+ 5 drops of 1M NaOH

pH 4.0 buffer

volume of CH3COOH used volume of CH3COONa used experimental pH of buffer

Post-laboratory Assignment

1. Compare your pH values from the meter, the pH paper and the indicator solution. How well

do they agree? Which do you think is most accurate? And least accurate?

2. In Part A of this experiment, which base had a higher pH? Explain why. And which acid

had a lower pH? Explain why.

3. In Part B, three salts were added to water. Write balanced chemical equations that show why the pH changed or did not change. (Again, see Examples 1 and 2 under hydrolysis in the reading.)

4. In Part C, how did addition of HCl and NaOH affect the pH of the acetic acid/sodium acetate buffer? How did the additions affect the pH of water? What conclusions can you draw from these results?

5. A buffer solution is made by mixing Na2HPO4 with NaH2PO4 . Which of these is the acid and which is the base?

a. Write an equation to show how the buffer neutralizes any added acid (H3O+). How does

this reaction affect the equilibrium?

b. Write an equation to show how the buffer neutralizes any added base(OH-). How does this

reaction affect the equilibrium?

6. Which of these mixtures can serve as a buffer solution? Explain why or why not.

a. CH3NH2 and CH3NH3+CI -

b. NaOH and NaCI

c. HBr and KBr

7. Calculate the pKa and the Ka of CH3COOH using the Henderson-Hasselbach equation.

Compare your results to the literature value and comment on how accurate you were.

8. Calculate the Ka of H3BO3.

Acids, Bases, Buffers and pH

Pre-laboratory Assignment

1. A solution has a pH of 8.7.

a. What is the pOH of this solution?

b. Is the solution acidic, basic or neutral?

2. Solution A has a hydronium ion concentration of 3.8 x 10-8 M. Solution B has a hydronium ion concentration of 2.5 x 10-4 M.

a. Which solution has the higher hydronium ion concentration?

b. Which solution has the higher pH?

c. Calculate the pH of the more acidic solution. Show all work.

3. Calculate the pH of the following buffer solutions.

a. 50.0 mL of 0.10M HF and 50.0 mL of 0.10M NaF. Ka for HF = 7.1 x 10-4

b. 40.0 mL of 0.10 HNO2 and 60.0 mL of 0.10 M NaNO2 Ka for HNO2 = 4.5 x 10-4

Hint: Calculate the weak acid and conjugate base concentrations once the two solutions are combined and use the Henderson-Hasselbach equation.