Chemistry

Key unit competency

To be able to:

• Prepare solutions, measure their pH, and calculate the pH of acidic and alkaline solutions.

• Explain the concept of buffer solution, hydrolysis of salts and discuss its applications in manufacturing industry and biological processes.

Learning objectives

At the end of this unit , students will be able to:

• Define the degree of ionization (α);

• Define the terms Ka, pH, pKa, Kb, Pkb and Kw;

• Write equations for salt hydrolysis reactions and the expression for the hydrolysis constant;

• Explain how buffer solution control pH;

• Explain the buffer capacity in relation to buffer range;

• Describe the applications of buffer solution in domains such as biological processes, agriculture, natural system (e.g. lakes) and industrial manufacture of cosmetic and drugs.

• Perform calculations involving pH, Ka, pKa, Kb, pKb and Kw;

• Interpert th values of Ka and Kb in relation to the strength of acids and bases;

• Prepare different solutions and appropriately use pH-meter to measure their pH;

• Compare the strength of acids and bases on the same concentration using the values of Ka and Kb;

• Relate the values of pH and pOH;

• Calculate the pH and the hydrolysis constant of aqueous solutions of salts;

• Perform experiments to show that hydrolysis of some salts results in neutral, acidic and alkaline solutions;

• Prepare buffer solutions of different pH values;

• Derive Henderson-Hasselbalch relation and use it to calculate the pH of buffer solution.

Introductory activity

Observe the above representations a, b, c, d, e and f which represent some useful products made of acids, bases and salts and answer to the following questions.

1. Identify one chemical substance found in each product above.

2. State whether the chemical substance identified above is an acid, a base or a salt.

3. What do you comment about the pH of the chemical substance in (1) 

Many fruit juices and other soft drinks contain acids and this can be identified by their sour or sharp taste. They contain many dissolved compounds of which acids are among. Some acids are chemicals that can be harmful or dangerous to human and should never be tested by drinking them (i.e. sulphuric acid used in car batteries). 

Bases are products commonly used in our daily life such as soap, toothpaste, magnesium syrup, baking soda. Solutions of bases are slippery on touch because they attack oils on the skin and convert them into soaps. This makes them to be good cleaning agents. Bases dissolve in water to form solutions called alkaline solutions. 

Acids can react with bases to form compounds known as salts which are of great importance in nature. For example ammonium chloride is used as an electrolyte in dry cells, calcium carbonate is used to manufacture cement, potassium nitrate is used as a fertilizer, magnesium sulphate and iron (II) sulphate used to manufacture of drugs.

9.1. Degree of ionization in relation to strength of acids and bases

Activity 9.1

(a) Explain the following terms;

 (i) an acid

 (ii) a base

(b) Give an example of an acid and a base which you encounter in your daily life.

(c) Explain the difference between a strong acid and a strong base.

(d) How is a weak acid different from a weak base?

(e) Give an example of weak acids and weak bases.

(f) Write down chemical equations to show ionization of a weak acid and weak base water.

9.1.1. Acids and bases

In the previous years of your study, an acid has been defined as a substance that donates hydrogen ions while a base was defined as a substance that accepts hydrogen ions. Acidity and alkalinity are measured with a logarithmic scale called pH. 

• Strong and weak acids 

When an acid is dissolved in water, there are more hydrogen ions than hydroxide ions in the solution meaning that the solution is acidic. We can distinguish strong and weak acids. 

A strong acid is ionized completely when dissolved in water. For example hydrochloric acid ionizes completely in water to form hydrogen or hydroxonium ions and chloride ions as illustrated below. The direct arrow indicates that the compound is completely ionized.

When a weak acid is dissolved in water, only a small proportion of it ionizes to form ions. For example ethanoic acid ionizes partially in water to form ethanoate ions and hydrogen or hydroxonium ions. The following illustration shows that the ethanoic acid as a weak acid dissolves partially in water, the direct and indirect arrows symbolize the partial ionization or an equilbrium dissocition reaction

Other weak acids include sulphurous acid (H2 SO3 ), carbonic acid (H2 CO3 ), phosphoric acid (H3 PO4 ), methanoic acid (HCOOH) etc. In general the organic acids are known to be weak acids.

The strength of an acid can be influenced by four factors namely: 

(a)Bond strength: When the bond holding a hydrogen atom is strong, the compound formed is a weak acid since the bond cannot be easily broken to release hydrogen ions in solution. As an example, consider the halogen acids HF, HCl, HBr and HI. Their bond strength decreases with a decrease in electronegativity of the halogen atoms down the group 17. HI is a stronger acid than HCl or HBr because of a weak H – I bond due to its lower electronegativity of iodine atom than Cl or Br. This bond is easily broken down to release hydrogen ions in solution. HF is a weak acid due to its strong H – F bond caused by a highly electronegativity of fluorine atom. This bond is not easily broken down to release hydrogen ions in solution. 

(b)Nature of the solvent: The more basic the solvent in which the acid is dissolved, the stronger the acid. This is because the proton released by the acid is easily accepted by the basic solvent. This enables the acid to continue ionizing in the solvent to produce hydrogen ions e.g. Benzoic acid is a stronger acid in aqueous ammonia than in water. 

(c) Presence of a halogen atom in organic acids: The halogen atom increases acidity because, being more electronegative than carbon it tends to withdraw electrons towards itself (negative inductive effect). This reduces the electron density in the O – H bond thus weakening it. In solution, the O – H bond is easily broken down to release hydrogen ions.This explains why chloroethanoic acid is a stronger acid than ethanoic acid. 

(d) Number of carbon atoms in organic acids: As the number of carbon atoms increases, the acidity of the organic acid decreases. This is because a long chain alkyl group pushes electrons towards the carboxyl group (positive inductive effect) increasing the electron density on the O – H bond. This makes the O – H bond stronger and not be easily broken down to release hydrogen ions in solution.This explains why methanoic acid is a stronger acid than ethanoic acid.

• Strong and weak bases 

When a base is dissolved in water, there are more hydroxide ions in the solution than hydrogen ions. This kind of solution is called alkaline. Strong and weak bases can be distinguished based on their ionization character. 

When a strong base is dissolved in water, it ionizes completely. For example sodium hydroxide (NaOH) ionizes in water to form sodium ions and hydroxyl ions. Similarly, potassium hydroxide (KOH) is a strong base because it dissociates completely when dissolved in water as illustrated below:

For weak base, they dissolve partially in water. For example ammonia solution ionizes partially in water to form ammonium ions and hydroxyl ions..

Phenylamine (C6 H5 NH2 ) and Hydroxylamine (NH2 OH) are other examples of weak bases 

9.1.2. Degree of ionization (α) 

An acid or a base when dissolved in water, it ionizes completely or partially depending on its characteristic as strong or weak. 

The degree of ionization can be defined as the ratio of the number of ionized molecules to the total number of molecules dissolved in water.It is symbolized by αa in the case of acid and αb in the case of base.

When a weak acid or weak base is dissolved in water, partial ionization occurs such that equilibrium is set up between the undissolved molecules and the ions formed in water. This equilibrium is known as ionic equilibrium of an acid or base. The degree of ionization is a fraction ranging from 0 to 1. a is equal to zero for the insoluble substances or non-electrolytes (i.e. substances which do not ionize in water). a is equal to 1 for strong electrolytes such as strong acids or bases because they ionize completely in water. a is less than 1 in the case of weak electrolytes such as weak acids or bases because they ionize partially in water. The degree of ionization can also be expressed as percentage as shown by the relation below.

The greater the degree of ionization, the stronger the acid or base. For example, the concentration of an acid that undergoes ionization is equal to the concentration of hydrogen ions formed as shown below.

Example 

What is the degree of dissociation of a weak acid in a 0.25M solution, given the concentration of H+ as 0.001 mol dm-3? 

Solution

Checking up 9.1 

1. A 0.1moldm-3 methanoic acid solution contains 0.0042 moldm-3 of hydrogen ions. Calculate the percentage of the acid that is ionized. 

2. A solution of 0.035M nitrous acid contains 0.0037M of hydrogen ions. Calculate the percentage ionization of the acid.

9.2. Explanation of acid and base dissociation constants (Kand K)

Activity 9.2

1. Explain your understanding of the following expressions

 (a) Acid dissociation constant.

 (b) Base dissociation constant.

2. (a) Draw an equation to show ionization of ethanoic acid in water.

 (b) Derive an expression for acid dissociation constant, Ka of ethanoic acid in water.

3. (a) Express an equation for the ionization of methylamine in water.

 (b) Write the expression for the base dissociation constant, Kb of methylamine.

Weak electrolytes such as weak acids or weak bases are partially ionized in water thus their dissociation is reversible and can reach equilibrium. Therefore the equilibrium constants can be used to explain the strength of acids and bases based on their ionization reactions. 

9.2.1. Acid dissociation constant, Ka 

The acid dissociation constant also known as acidity constant or acid ionization constant is the equilibrium constant for the ionization reaction of an acid. It is denoted by Ka . 

This constant is a quantitative measure of the strength of the acid in solution in units of moldm-3. Acid dissociation constants are mostly associated with weak acids because strong acids are completely ionized in the aqueous solution and their Ka values are extremely large or infinity. 

Consider a weak acid HA, its ionization is represented as follows:

The equilibrium constant for the dissociation of HA is as follows:

9.2.2. Base dissociation constant, Kb 

The base dissociation constant also known as base ionization constant is the equilibrium constant for the ionization reaction of a weak base denoted by Kb . The equilibrium constant Kb measures the strength of the base in solution in units of moldm-3. Base dissociation constants are mostly associated with weak bases because strong bases are completely ionized in the aqueous solution and their Kb values are extremely large. 

Consider a weak base, BOH undergoing ionization which is represented as follows:

The base dissociation constant Kb can be expressed by applying the equilibrium law:

Note: Kb provides a measure of the strength of the base. When Kb is large, the base is highly dissociated, and the base is strong. And when Kb is small, very little of the base is dissociated and the base is weak. 

9.2.3 Relationship between equilibrium (dissociation) constant and the degree of ionization α; Ostwald’s dilution law

 The relationship between the equilibrium constant and the degree of dissociation of weak electrolytes has been introduced by Ostwald in 1888. 

The Ostwald’s dilution law states that “the degree of ionization of a weak electrolyte is inversely proportional to the square root of the molar concentration of the electrolyte”.

Where C is the molar concentration and Ka is the equilibrium constant.

 Consider the dissociation equilibrium of formic acid HCOOH which is a weak electrolyte in water.

Where α is the degree of ionization and represents the fraction of the total concentration of HCOOH that exists in the completely ionized state. Further (1 – a) is the fraction of the total concentration of HCOOH in the unionized state. 

Relationship between Ka and a. 

If the total concentration of the acid is C at the initial step and its degree of ionization is aC; at the equilibrium Ca, Ca and C (1 – a) represent the concentration of H+, and HCOOH- respectively as shown below. 

If a is too small, Ka = a2 C

The higher the Ka value, the greater the degree of ionization and the stronger the acid.

The higher the Kb value, the greater the degree of ionization and the stronger the base. 

Worked examples 

1. (a) A solution of a weak acid, HA contains 0.25M, given the concentration of H+ as 0.001 moldm-3.Calculate the dissociation constant of HA. (b) Determine the degree of ionization of this acid in 1M solution? 

Solution:

(b) Degree of dissociation in a 1 M solution:

2. Calculate the degree of ionization of 0.1M acetic acid, CH3 COOH if its ionization constant Ka = 1.8 x10-5 moldm-3.

Checking up 9.2

1. Calculate the degree of ionization of a 0.04M ethanoic acid solution at 25 0C given that its Ka is 1.3 x 10-5 mol dm-3.

2. The acid dissociation constant of a monobasic acid is 4.39 x 10-5 mol dm-3 at 25 0C.

 Calculate the degree of ionization of a 0.01M solution of the acid.

3. (a) Write an equation for the ionization of methylamine in water.

 (b) Express the dissociation constant Kb for methylamine.

 (c) The hydroxyl ion concentration of a 1M methylamine solution is 0.04 moldm-3.

 Calculate the Kb for methylamine.

The relationship between Ka and Kb is derived from the dissociation of conjugate acid and conjugate base. For example considering NH4 + and NH3 ; NH4 + is a conjugate acid and NH3 is a conjugate base. 

The ionization of a conjugate acid in water is shown below:

Ka and Kb are related each other through the ionization constant.

When we multiply by–log10 to both sides, we get:

Checking up 9.3

1. (a) Ethyl ammonium ion,CH3CH2-NH3 + is a conjugate acid.Write an equation for the ionization of this ion in water.

 (b) Write the expression of acid dissociation constant, Ka ,of ethyl ammonium ions water.

2. (a) Ethylamine, CH3CH2-NH2

 is a conjugate base. Write the chemical equation for ionization of ethylamine in water.

 (b) Write an expression for the base dissociation constant, Kb of ethylamine in water.

3. From the reactions above, derive an expression to show the relationship between Ka and Kb

9.4. Use of Ka or pKa and Kb or pKb to explain the strength of acids and bases 

Activity 9.4 1. (a) Ethanoic acid has pKa of 4.77 at 25 0 C. What is meant by pKa of the acid? 

(b) Given the following acids and their corresponding Ka values in the table below,

State the strongest and the weakest acid and justify your answer. 

2. The ionization of ethanoic acid in water is shown by the following equation :

Use the above equation to write to relate the acid dissociation constant, Ka and the degree of ionization, a

The acidic and basic dissociation constants are respectively represented by Ka and Kb . Based on the negative logarithm, pKa and pKb are derived from–log10Ka and –log10Kb respectively. Therefore, Ka , pKa , Kb and pKb are more helpful in predicting whether a species donates or accepts a proton.

Checking up 9.4

9.5. Explanation of ionic product of water
Activity 9.5

Water is a neutral liquid at pH = 7

a. Explain why at 25 °C water has a pH equals to 7.

b. (Demonstrate how the ionization of water molecules is made.

c. What do you understand by “Ionic product of water ?

Water is a weak electrolyte and neutral in nature. Pure water is commonly known as universal solvent and water takes part in many equilibrium reactions. Water ionizes as shown by the reaction equation as follows:

The ionic product of water increases with temperature (Table 9.3). However at all temperatures, the concentration of H+ ions remains equal to the concentration of OH- ions in pure water.

9.6. Definition and calculations of pH and pOH of acidic and alkaline solutions 
Activity 9.6

The pH is a scale commonly used to measure the degree of acidity or alkalinity of a solution.It is measured on a scale of 0 to 14. The term pH is derived from “p,” which is a mathematical symbol expressing the negative logarithm, and “H,” the chemical symbol for Hydrogen. In general pH can be defined as the negative logarithm of Hydrogen ion activity or pH = -log [H+] 

Both pH and pOH are two methods that are used to describe the strength of acids, alkali, or ionizable salts (i.e. salt of a weak acid and strong base, salt of weak base and a strong acid or salt of weak acid and weak base).The pH is a measure of the hydrogen ion concentration of a solution.

Within this section, the following terms will be commonly used such as [H+] = hydrogen ions concentration; [OH- ] = hydroxide ions concentration. 

The pH can be expressed as the negative logarithm of the concentration of [H+] in a solution. The Solutions having a high concentration of hydrogen ions have a low pH and solutions with a low concentrations of H+ ions have a high pH. On the other hand, “pOH” is the negative of the logarithm of the concentration of hydroxide ions in a solution.It describes how alkaline a solution is; the more alkaline a solution is, the higher the concentration of hydroxide ions in the solution.

Importance of pH and pOH 

Knowing the pH and pOH is important because it helps us to find the concentration of hydrogen or hydroxide ions in the solution; it allows determining the acidity and alkalinity of the solution. While knowing the pH of a solution, the strength of the acid and base in the solution can be determined and the separation of strong and weak acids or strong and weak bases can be done based on the value of pH or pOH. 

Mathematically pH or pOH can be expressed as follows:

The Figure 9.1 shows the location of acidic, neutral and alkaline compounds based on pH scale while the Figure 9.2 shows the pH or pOH scale (from 0 to 14). Figure 9.3 shows examples of compounds corresponding to different values of the pH scale.

The pH of strong acids is ranged from 1-3 and weak acids from 4-6. Strong bases have values of pH ranging from 11 to 14 and weak bases have values of pH ranging from8 to 10.

Note: When a drop of hydrochloric acid is added to water, the concentration of hydrogen ions increases and pH decreases. The pH of an acidic solution is less than 7. On the other hand when sodium hydroxide is added to water, the added hydroxyl ions reduce the concentration of hydrogen ions and pH increases. pH of an alkaline solution is greater than 7. The Table 9.4 shows the variation of pH in function of the hydrogen ion concentrations.

9.6.1. pH of strong acids 

Strong acids like hydrochloric acid has a pH around 0 to 1.The lower the pH, the higher the concentration of hydrogen ions in the solution. A strong acid is one which completely dissociates into its ions in water. This makes calculating the hydrogen ion concentration, which is the basis of pH,easier than for weak acids. The following are examples on how to determine pH of a strong acid.

1. Determine the pH of a 0.025 M solution of hydrobromic acid (HBr) solution 

Hydrobromic acid or HBr, is a strong acidand will ionize completely in water by giving H+ and Br- ions. For every mole of HBr, there will be 1 mole of H+ and the concentration of H+ will be the same as the concentration of HBr. Therefore, [H+] = 0.025 M.

9.7. Salt hydrolysis
Activity 9.7.

1. What do you understand by the term “a salt”?

2. Explain two types of salts. Give examples in each case.

3. Discuss how salts are formed. Explain any four methods of preparing salts.

4. What do you understand by “salt hydrolysis”?

5. Give four examples of salts which undergo hydrolysis.

6. a. (i) Write an equation for the hydrolysis of sodium benzoate in water.

(ii) Write an expression for the hydrolysis constant, Kh of sodium benzoate.

 b. A solution contains 0.2 moles of sodium benzoate per litre at 25°C.  

   Calculate the pH of the solution given that the hydrolysis constant, Kh of sodium benzoate is 1.6x10-10 mol dm-3 at 25°C.

In the previous sections, we have seen that the pH value can be used to determine the strength of acid or a base. When the concentrations of [H+] > [OH- ], the water becomes acidic and when [H+] < [OH- ], the water acquires basic nature. A salt hydrolysis is a phenomenon that is observed when there is change of the concentrations of hydrogen or hydroxyl ions in a solution. 

A salt is a chemical substance formed when either part or all the ionizable hydrogen of an acid have been replaced by a metallic ion or ammonium radical.Salts are strong electrolytes because when dissolved in water they dissociate almost completely into two different ions such as cations and anions.

Salt hydrolysis is defined as a reaction in which the cation or anion or both of a salt react with water to form an acidic or alkaline solution. 

When the cations from the salt are more reactive than anions, they interact with water molecules detaching hydroxyl ions. This leaves hydrogen ions in solution making it acidic.

When the anions from the salt are more reactive than cations, they interact with water removing hydrogen ions and releasing the hydroxyl ions in solution. This results in the alkaline solution.

In general, the process of salt hydrolysis is the reverse of neutralization: Salt + Water ↔ Acid + Base 

If acid is stronger than base, the solution is acidic and in case base is stronger than acid, the solution is alkaline. When both the acid and the base are either strong or weak respectively, the solution is generally neutral in nature. As the nature of the cation or the anion of the salt determines whether its solution will be acidic or basic, it is proper to divide the salts into four categories.

(i) Salt of a strong base and a weak acid.

(ii) Salt of weak base and strong acid.

(iii) Salt of weak base and weak acid.

(iv) Salt of a strong acid and a strong base.

9.7.1. Salt of weak acid and strong base

The solution produced by a strong base and a weak acid is basic. The anion reacts with water to form a weak acid and OH- ions. 

Ethanoate ions disturb the ionic equilibrium of water where they accept hydrogen ions from water. This makes the resultant solution alkaline due to excess hydroxyl ions present. Such a solution has pH value greater than 7. The ions produced CH3 COO¯ in turn react with water to form a weak acid, CH3 COOH and OH¯ ions as follows:

pH of the resultant solution on this hydrolysis is then calculated from, pH = 14 – pOH

9.7.2. Salt of weak base and strong acid

These are salts which undergo cation hydrolysis to form acidic solutions. Examples of such salts include;

Ammonium ions disturb the ionic equilibrium of water by accepting hydroxyl ions from water. This makes the resultant solution acidic due to excess hydrogen ions from water thus a pH value less than 7.

9.7.3. Salt of weak base and weak acid 

These are salts which undergo both cation and anion hydrolysis when dissolved in water. The nature of the resultant solution depends on the relative strength of the weak base and the weak acid. It may finally be acidic, alkaline or neutral. 

E.g. if Kb of the anion is greater than Ka for the cation, the solution formed is alkaline because the anion is greatly hydrolyzed to produce more hydroxyl ions in solution. 

When the Kb of the anion is less than Ka the cation, the resultant solution is acidic because the cation will be hydrolyzed to a greater extent producing excess hydrogen ions in solution. And if Kb is approximately equal to Ka , the resultant solution is neutral. 

Examples of such salts include:

9.7.4. Salt of strong acid and strong base

Salts of strong acids and strong bases dissolve in water to give neutral solutions.

For example; 

Aqueous solutions of the salts consist of ions with very little affinity for hydrogen ions or hydroxyl ions in water. The ions of such salts are only hydrated (i.e. surrounded by water molecules). For example; Sodium chloride dissociates in water to give the anion Cl¯ .

Since hydrochloric acid is a strong acid, Cl¯  is very weak base. Cl¯ is unable to accept a proton (H+) from an acid, particularly water. That is why Cl¯  does not hydrolyze. The pH of sodium chloride solution remains unaffected. 

Hydrolysis constant, Kh and the degree of hydrolysis 

Consider a salt, CH3 COONa whose concentration is Cs . The salt undergoes anion hydrolysis as follows:

Dissociation of the acid:

9.8. Buffer solution
Activity 9.8

1. Explain the term buffer solution.

2. Analyze and differentiate the types of buffer solutions.

3. Discuss how pH and pOH of a buffer solution can be calculated.

9.8.1. Definition of buffer solution

Buffer solution is a solution that resists a pH change when a small amount of a base or acid is added to it.

A buffer solution consists of a solution of a weak acid and its salt with a strong base. This is called an acidic buffer. Examples of acidic buffers include: Ethanoic acid and  sodium ethanoate solution, carbonic acid and sodium hydrogen carbonate. 

A buffer solution may also consist of a solution of a weak base and its salt of strong acid. This is called a basic or alkaline buffer. Examples of basic buffers include: Aqueous ammonia and ammonium chloride solution, aqueous ammonia and ammonium sulphate or ammonium nitrate.

 In general a buffer solution is an aqueous solution which involves the mixture of a weak acid and its conjugate base, or vice versa. 

9.8.2. pH of buffer solution 

The pH of a buffer solution or the concentration of the acid and base can be calculated using the Henderson-Hasselbalch equation. Henderson-Hasselbalch equation was given by Lawrence Joseph Henderson  (1878-1942) and Karl Albert Hasselbalch (1874-1962). 

Taking negative logarithms to base 10 on both sides;

This equation is known as the Henderson-Hasselbalch equation. Therefore, the pH of an acidic buffer depends on the relative concentration of the salt and the acid in the mixture.

9.9. Preparation of buffer solutions of different pH
Activity 9.9

1. Write an expression to show Henderson-Hasselbalch equation.

2. Explain how Henderson-Hasselbalch equation is used for the determination of the mass of salt to be dissolved when preparing an acidic buffer solution.

3. State two methods of preparing buffer solutions of different pH.

A buffer solution is a solution that resists a change in pH, because it contains species in solution able to react with any added acid or base.

In general, preparing a buffer solution requires either a weak acid and a salt of the acid’s conjugate base or a weak base and a salt of the base’s conjugate acid.

9.9.1. By mixing weak acid and its corresponding salt or weak base and its corresponding salt

It is very important to prepare buffer solutions of known pH in the laboratory. Using the Henderson-Hasselbalch equation, there are two terms which determine the final pH of the solution that is pKa , whose value is responsible for the ‘coarse selection’ of pH and the ratio [conjugate base]/ [acid] that provides ‘fine tuning’ to the final pH.

So, to prepare an acidic buffer solution of known pH, select an acid whose pKa is within the range of one unit of the desired pH. The ratio of salt to acid concentrations is then adjusted to achieve the desired pH.

Example 

Suppose you want to prepare an acidic buffer with a pH of 4.0. A suitable weak acid would be ethanoic acid CH3 COOH because its pKa is 4.8. The conjugate base is ethanoate ion CH3 COO- , which is provided by the sodium ethanoate salt, CH3 COONa. Ethanoic acid is available as a laboratory bench reagent with concentration of 1.0 mol dm-3. 

The question is, what mass of sodium ethanoate should I add to the ethanoic acid to make this buffer solution? 

Using the Henderson-Hasselbalch equation;

Therefore, an acidic buffer solution of pH 4.0 will be prepared by dissolving 13 g of sodium ethanoate in 1.0 dm3 of 1.0 mol dm-3 ethanoic acid in a volumetric flask. Insert a calibrated pH meter to monitor the pH of the prepared acidic buffer solution. 

Note that it is the ratio of acid to conjugate base that is important in determining the pH of the buffer solution not the concentrations. However, when more concentrated solutions are used, the buffer solution can efficiently react with added acid and base before becoming saturated.

9.9.2. By partial neutralization 

Buffers can also be prepared by the partial neutralization of a weak acid by a strong base through titration process.

This is done by running excess weak acidic solution from the burette into a strong basic solution in the conical flask until half neutralization takes place. A weak acid is partially neutralized by the strong base to form salt of weak acid. This salt is formed together with excess of the weak acidic solution which makes the resultant solution a buffer solution. Do this while measuring the pH of the resultant buffer solutionwith a pH meter.

For example reacting 40 ml of 1.0M propanoic acid (C2H5COOH) solution (pKa is 4.87) with 60.0 ml of 0.10 M sodium hydroxide solution (NaOH). This will produce a buffer solution consisting of propanoic acid and sodium propanoate of pH 4.13.

Checking up 9.9

The following materials and chemicals are provided by the technician laboratory.

• 100ml Beaker,

• 10ml measuring cylinder,

• 50ml measuring cylinder,

• Electronic weighing scale,

• Calibrated pH meter,

• Glass rod,

• 100ml volumetric flask

• Ethanoic acid,

• sodium ethanoate and distilled water 

Procedure

1. Measure exactly 50 ml of distilled water to a 100 ml beaker.

2. Using a measuring cylinder, add 5 ml of 0.3M ethanoic acid to the beaker.

3. Then weigh exactly 0.3g sodium ethanoate (CH3COONa)

4. Add a little of the sodium ethanoate at a time, stirring the mixture with a glass rod to dissolve.

5. Insert the calibrated pH meter into the resultant solution in the beaker and measure its pH.

6. Quantitatively transfer the buffer solution to a 100 ml volumetric flask.Add distilled water up to the mark. Cap and invert the flask twice to mix.

9.10. Explanation of the working of buffer solutions
Activity 9.10

1. a. Explain what is meant by a buffer solution?

 b. How does buffer acts?

2. Explain each of the following terms as used in the study of buffer solutions.

a. Buffer capacity

b. Buffer range

3. Explain two applications of buffer solutions that are used in:

a. Biological processes

b. Manufacturing industry

In a buffer solution, the weak acid and the conjugate base or weak base and its conjugate acid are responsible for controlling pH.

9.10.1 Working process of an acidic buffer

Consider a buffer solution made of ethanoic acid and sodium ethanoate solutions. Ethanoic acid is partially ionized because it is a weak acid whereas sodium ethanoate is fully ionized because it’s a strong electrolyte.

When a small amount of a base is added, the added hydroxyl ions react with ethanoic acid to form water. This prevents an increase in the concentration of hydroxyl ions hence pH is kept constant.

9.10.2. Working process of a basic buffer 

Consider a buffer solution of aqueous ammonia and ammonium chloride. Aqueous ammonia is partially ionized while ammonium chloride is completely ionized because it is a strong electrolyte.

Such a solution contains a few hydroxyl ions and a large proportion of ammonium ions from the salt. 

When a small amount of a base is added, the added hydroxyl ions react with ammonium ions to form un ionized aqueous ammonia. This prevents any increase in the concentration of hydroxyl ions in the solution hence no change on pH.

Similarly when a small amount of an acid is added, the added hydrogen ions react with un ionized aqueous ammonia to form water. This prevents an increase in the concentration of hydrogen ions thus keeping pH constant.

9.10.3. Definition of buffer capacity and buffer range 

Buffer capacity 

This is a measure of the ability of a buffer solution to resist changes in pH when a base or acid is added. On addition of a base or acid to a buffer system, the effect on pH change can be large or small depending on the initial pH of the solution and it’s ability to resist that pH change. 

Buffer capacity is therefore defined as the number of moles of acid or base which when added to one litre of a buffer solution changes its pH by 1. 

Buffer capacity (β) has no units since it is a ratio of number of moles of acid or base added, to change in pH of buffer solution. 

Mathematically;

Buffer capacity is the efficiency of the buffer solution to control pH. The acid added to one litre of buffer solution changes its pH by only 1 unit.The buffer capacity is maximum at pH=pKa where [acid]=[conjugate base]. 

Buffer range 

Buffer range is the pH range within which the buffer solution is effective. 

Buffer range corresponds to the change in pH of about ±1. For example, acetic acid/ sodium acetate buffer works at optimum pH of 4.8. This means that it would work from pH 3.8 on addition of an acid to pH 5.8 when a base is added. Beyond this pH range, the buffer has no capacity to buffer the solution. 

The effective buffer range of buffer solutions is different and depends on the acid or base dissociation constant.

9.10.4. Applications of buffer solutions 

pH of various solutions has to be controlled either in industries or in cells of living organisms because any slight change may greatly affect the functioning of the whole system.

 In biological processes

• Buffer solutions are used to maintain the pH of human blood constant at 7.4. Intravenous injections must be correctly buffered so that they do not change the pH of blood in humans. 

• Proteins in living organisms act as buffers by controlling pH in body cells. They are composed of amino acids linked together in long chains. Each amino acid has two functional groups; the amino group and acidic carboxyl group. The carboxyl group donates hydrogen ion when pH in body cells is high (alkaline). Amino group accepts hydrogen ion from body cells when pH is low (Acidic). 

• Buffers keep the correct pH for enzymes to work in many organisms. Enzymes work best in a specific range of pH. If this pH is not controlled, the rate of enzyme activity slows down or enzymes stop working due to denaturation. 

• In bacteriological research, culture media are generally buffered to maintain the pH as bacteria required a constant pH all the time to grow. 

In agricultural processes 

• Plants grow well in soils with a narrow pH range. Some soils become acidic due to acidic rains which is a serious problem to plant or crop growth. Organic matter and mineral salts in a fertile soil act as buffers. It is important to maintain

 • pH of the soils for proper growth of vegetation and soil microorganisms.

 In natural systems

 • Water bodies such as lakes, rivers and streams are important habitats for aquatic organisms e.g. fish and young amphibians. They should have a stable pH to ensure survival of these organisms. Otherwise extreme pH may cause physical damage to the gills and fins of fish.

In industries 

• Buffer solutions are used in fermentation to control pH changes such that anaerobic fermentation bacteria are not killed. This prevents the solutions from becoming too acidic and spoiling the product. 

• Buffer solutions are used in manufacture of cosmetics whose pH must be controlled to prevent adverse effects on body cells. 

• They are used in the production of pharmaceutical drugs to prevent deterioration when the drug is administered or stored. Buffers ensure stability and clinical effectiveness of the drugs. 

• Dyes used in textile industries are buffered in order to maintain colour strength in different fabrics after production. 

• Buffers are used in leather industries to control pH during tanning and dyeing. This gives a product of fine texture and colour.

END UNIT ASSESSMENT

SECTION A: Multiple choice questions