Science and Materials in Nature

In this section, you will learn about acids, bases, and how they interact with water.

Acids are substances that release hydrogen ions (H+) and usually taste sour, like lemon juice. Bases accept hydrogen ions or release hydroxide ions (OH–) and can feel slippery, like soap. Water is neutral and can act as both an acid and a base.

We use the pH scale to measure how acidic or basic a solution is, ranging from 0 (strong acid) to 14 (strong base), with 7 being neutral. You’ll also learn how acids and bases react to form salts, which are important in everyday products like

table salt and fertilisers. Finally, we’ll discuss how to measure the concentration of solutions and why controlling pH is important in fields such as agriculture, medicine, and manufacturing.

At the end of this section, you should be able to:

• differentiate among acids, bases and water.

• apply the knowledge of acids and bases in analysing the formation of salts and their uses.

• describe how to measure the concentration of solutions and how to use the pH scale to identify the concentration of acids and bases/alkalis.

KEY IDEAS

• Acids release hydrogen ions (H+) in water, taste sour, turn blue litmus red whilst Bases accept H+ ions or produce hydroxide ions (OH–) in water.

• Neutralisation involves the reaction between acid and base to form salt and water only.

• Molarity (M) measures how strong a solution is (moles per litre), important for controlling acid/base strength. In determining how concentrated

solution, consider purpose of solution, types of solution, properties and dilution requirement

• pH Scale helps determine hydrogen ion concentration and ensures product safety and effectiveness.

• Acid-base Indicators are substances used to determine is acidic, basic, or neutral state of a substance

• Acids and bases are used in cleaning products, food preparation, medicine, farming, water treatment, and chemical manufacturing.

Acids Let us do this activity to explain what an acid is.

Activity 1.1 Explanation of Acid

Materials needed: lemon juice or vinegar (an acid), water, beaker or clear glass or plastic cup, a spoon and drinking straw(s).

Steps:

1. Half-fill the beaker, clear glass, or plastic cup with water and add one spoonful of lemon juice or vinegar.

2. Dip the drinking straw in the mixture (water and lemon juice or vinegar) and have a taste.

3. If using litmus paper, dip it in the mixture (water and lemon juice or vinegar) to see the colour change.

4. What did you observe?

a. The taste of the water becomes sour

b. Blue litmus paper turns red.

5. Draw your conclusion: example, acids release hydrogen ions (H+) when they dissolve in water. This is the reason the taste of the water changed and turned blue litmus paper red.

Safety Precautions

1. Do not drink the water.

2. If using a drinking straw, each of you must have one drinking straw.

Acids release hydrogen ions (H+) when they dissolve in water due to the way their molecules interact with water molecules.

• Acids are made up of molecules that contain hydrogen atoms. For

example, hydrochloric acid (HCl) has hydrogen (H) and chlorine (Cl) in its molecule, and vinegar has acetic acid (CH₃COOH).

• When you add an acid to water, the acid molecules start to separate or dissociate. For example, when you put hydrochloric acid (HCl) in water, it breaks apart into hydrogen ions (H+) and chloride ions (Cl–): HCl → H+ + Cl–

• The hydrogen ions (H+) are what make the solution acidic.

• These H+ ions are very small and have a positive charge. When there are a lot of them in the water, the solution becomes more acidic. The more hydrogen ions (H+) that are released, the more acidic the solution gets.

• Acids release these H+ ions easily, which is why they taste sour, can corrode metals, and can cause a burning feeling on the skin when they are concentrated.

• Acids can be classified into organic and inorganic depending on their composition and origin (source).

Organic And Inorganic Acid

Organic Acid: The word “organic” means natural or from living things. Organic acids are acids that come from living things (plants and animals) and contain carbon in their chemical structure. Organic acids are usually found in fruits, vegetables, and other natural sources. — See Table 1.1 and Figure 1.1

Table 1.1: Some organic acids and their sources Organic acid Source Formic acid Bees (sting) Lactic acid Milk/ Yoghurt Citric acid Unripe lemon and grapefruit Palmitic acid Palm oil Amino acid Eggs, meat Organic acid Source Acetic acid Vinegar Salicylic acid Aspirin Tartaric acid Grapes Ascorbic acid Citrus fruit Stearic acid Fats; shea butter, cocoa butter

Figure 1.1: Sources of Organic Acids

Inorganic Acid: The word “inorganic” means not from living things. Inorganic acids are acids that do not come from living things. They are usually made from minerals and do not always contain carbon. These acids are often used in laboratories, factories, and even in car batteries. — See Table 1.2 and Figure 1.2

Table 1.2: Inorganic acids and their formulae Mineral / inorganic acid Source Hydrochloric acid (HCl) Gastric juice in the stomach Sulphuric acid (H₂SO₄) Sulphur dioxide released by volcanoes and industrial processes Nitric acid (HNO₃) Forms in the atmosphere during thunderstorms and is a component of acid rain Phosphoric acid (H₃PO₄) Present in rainwater, carbonated drinks such as Coca-Cola, Fanta, Bell-cola

Figure 1.2: Sources of inorganic acids

Activity 1.2 Try Work on Organic and Inorganic Acids Instruction: Write down your answer in your notebook for future reference.

1. Classify the following substances into the table provided below:

Lemon slices, orange slices, parazone, diluted sulphuric acid solution, vinegar (acetic acid), carbonated drinks, yoghurt, diluted HCl solution.

Organic acids Inorganic acids

2. What are the reasons for the classification in the above table?

Properties of Acids

Acids have certain characteristics, also known as properties, which make them different from other substances. These properties can be physical and chemical.

Let us find out how acid tastes through an activity!

Activity 1.3 Physical Properties of Acid

Materials needed: five beakers or cups, lemon juice or fruit, vinegar, orange juice or fruit, tomato juice or fruit and pineapple fruit or juice.

Steps:

1. Pour each substance into each beaker and label them clearly. If using the fruit, squeeze the juice into each beaker or cup.

2. Taste a drop of each substance in each beaker or cup.

3. Dip litmus paper in each substance of each beaker or cup.

4. Add water to each substance in each beaker or cup.

5. Record your findings in your notebook for future reference.

6. Now, answer these questions:

a. Describe the taste of each substance.

b. Copy and complete the table below:

Substance Effect of substance on litmus paper (colour change) Blue litmus paper Red litmus paper Lemon juice Tomato juice Orange juice Vinegar Pineapple juice

c. How does the mixture of water and each substance taste?

d. Give reason to support your answer in question (c) above.

e. From the activity list two physical properties of acid.

Safety Precautions

1. Label the beakers clearly to avoid mixing the substances up.

2. Wash hands with soap and water after the experiment.

Summary of the physical properties of acid you should know:

1. Acid solutions that are diluted taste sour.

2. Acids have a pH level less than 7.

3. Many acids are corrosive.

4. Acids change blue litmus paper to red Chemical Properties of Acids Imagine you have a stomach-ache because your stomach has too much acid in it, which makes you feel a burning sensation. If you take an antacid tablet, which is a base, the base in the tablet will react with the acid in your stomach. The acid and the base cancel each other out, forming salt and water, which makes your stomach feel better. This is called neutralisation reaction. Let us explore another activity!

Activity 1.4 Neutralisation reaction Materials needed: beaker, test tube, vinegar and baking soda.

Steps:

1. Pour some vinegar in the test tube.

2. Add the baking soda to the test tube containing vinegar and mix.

3. Dip litmus paper into the product formed in the test tube after the reaction.

4. Write your observation and conclusion in your notebook.

5. Record your answers to the following questions in your notebook:

a. Name the products formed.

b. What effect does the product formed have on litmus paper?

Safety Precautions

1. Do not eat the product formed.

2. Wash your hands with soap properly after the experiment.

Activity 1.5 Testing of Conductivity of Acidic Solution.

Instruction: This activity should be conducted under the supervision of your teacher or laboratory technician in the laboratory.

Materials needed: battery (6V), light bulb or LED, wires with clips iron nails or metal strips, beaker, distilled water, dilute hydrochloric acid (HCl)

solution or vinegar solution and a switch.

Steps:

1. Connect the battery, light bulb, and iron nails or metal strips in a circuit.

2. Place the iron nails or metal strips in a beaker of distilled water and observe what happens.

3. Add some dilute hydrochloric acid (HCl) to the water to form solution.

4. What can you observe? When you dip the electrodes into the acid

solution, the light bulb will light up.

5. Draw a conclusion for this activity: This shows that the acid solution can conduct electricity because there are now ions (H+ and Cl–) that allow electricity to pass through. — See Figure 1.3 for a set-up.

Figure 1.3: Setup of testing conductivity of acid Summary of some chemical properties of acids:

1. Strong acids can conduct electricity when they are dissolved in water because they release ions that allow electric current to pass through the solution.

2. They also react with bases in a process called neutralisation, where they combine to form salt and water. Example: HCl + NaOH → NaCl + H₂O Acid + Base → Salt + Wate r

3. Dilute acids react with metallic salts (carbonates) to form salts, water, and Carbon dioxide. Example, H₂SO₄+ CuCO₃→ CuSO₄+ H₂O + CO₂ Acid + metallic salt → Carbonate + water + carbon dioxide

6. They react with metals to produce hydrogen gas. Example, Zn + 2HCl → ZnCl₂+ H2 Uses of Acids in Everyday Life

1. Nitric acid and phosphoric acid are used in the production of fertilisers to help crops grow better.

2. In batteries, sulphuric acid is widely utilised.

3. Many soft drinks contain phosphoric acid as the main ingredient.

4. Sinks and sanitary ware are also cleaned with hydrochloric acid.

5. Sulphuric acid and hydrochloric acid are commonly used in laboratories for chemical experiments and to test materials.

6. Acetylsalicylic Acid (Aspirin) is a common pain reliever and anti- inflammatory drug.

Bases For easy understanding of basic substances, let us take up an activity!

Activity 1.6 Explanation of Bases or Alkali

Materials needed: baking soda (sodium bicarbonate), soap or liquid detergent, water, small cups or bowls, stirring sticks or rods, paper towels and cotton swabs (for tasting).

Steps:

1. Set up the samples and test the slippery feel.

a. Pour a small amount of baking soda into a cup and add water to make a solution. Stir well until it dissolves.

b. In another cup, mix a little bit of liquid soap with water.

c. Dip a finger into the soap solution and rub it between their fingers.

2. Conclusion: When water mixes with soap, it reacts with water to create a slippery feel. This is because soap is an alkali or base that produces OH– ions in water.

3. Testing the bitter taste (optional and safe):

a. Dip a cotton swab into the baking soda solution.

b. Gently touch it to the tongue.

4. Write your answers in your notebook.

a. How does the baking soda solution taste?

b. Look at your environment and write down some common household bases.

c. How does the soap solution feel between your fingers?

5. Conclusion: A mild bitter taste indicates that the baking soda solution is a base or basic.

Safety Precautions

1. Do not taste any you come across, because some can be harmful.

2. Each of you must use a paper towel for the tasting test.

3. Wash your hands with water and soap after handling the solutions.

Bases produce hydroxide ions (OH–) when they dissolve in water. These ions make the solution basic or alkaline, just like how acids produce hydrogen ions (H+) to make a solution acidic.

Common household bases: Baking soda (Sodium Bicarbonate), ammonia, washing soda (Sodium Carbonate), soap, detergents, toothpaste, milk of Magnesia, shampoo, chlorine bleach, chalk (Calcium Carbonate).

Organic and Inorganic Bases

Organic bases are obtained naturally from plants and animals. They produce KOH (Potassium hydroxide). The decomposition of organic matter produces NH3 (ammonia). Examples: Petre, wood ash, cocoa peels, wasp stings.

Inorganic bases are bases prepared in the laboratory. Examples: KOH (Potassium hydroxide), NH4OH (ammonium hydroxide), Ca (OH)₂(Calcium hydroxide).

Activity 1.7 Try Work on Organic and Inorganic Base

1. Classify the following substances into the table provided below: Petre, wood ash, Potassium hydroxide, cocoa peels, wasp stings, ammonium hydroxide, Calcium hydroxide.

Organic bases Inorganic bases

2. What are the reasons for the classification in the above table?

3. Write down your answers in your notebook for future reference.

4. Share your result with a friend and educate them on the organic and inorganic bases.

Physical Properties of Bases

Remember in activity 1.6, you explored the taste of basic solutions and how basic substance feels when rubbed between fingers. Let us now look at how bases react with litmus paper as a property.

Activity 1.8 Physical Properties of Base

Materials needed: baking soda (sodium bicarbonate), lime water (calcium hydroxide solution) or household ammonia (diluted), water, three clear cups or beakers and red litmus paper.

Steps:

1. Label three cups as A, B, and C.

2. Pour water into each cup.

3. In Cup A, add a small amount of baking soda.

4. In Cup B, pour some lime water or ammonia.

5. Leave Cup C as just water (control sample).

6. Use stirring sticks to mix the substances in Cups A and B until they dissolve.

7. Dip red litmus paper into Cups A, B, and C.

8. Record your observations in your note for future reference.

9. Now, answer the following questions in your note:

a. Copy and complete the table below:

Substance Effect of substance on red litmus paper (colour change) Baking soda solution Lime water or ammonia Water

b. Give the reason for your answer recorded in the table above.

c. From the activity, list two physical properties of base.

Explanation

• Baking soda and lime water (or ammonia) turned the red litmus paper blue and showed a basic pH on the indicator. This tells us they are bases.

• Water did not change the colour because it is neutral.

Why did this happen?

• When bases dissolve in water, they release hydroxide ions (OH–) into the

solution.

• Hydroxide ions (OH–) are what make the solution basic. The more OH– ions there are, the stronger the base.

• This is why baking soda and lime water can make the red litmus paper turn blue; they produce OH– ions in water, which are basic.

What is the role of water?

Water helps dissolve the base and allows it to release the hydroxide ions.

That’s why these substances only become basic when they are mixed with water.

Some physical properties of base

1. Bases have a bitter taste. For example, baking soda and soap taste bitter.

2. They feel slippery or soapy when touched. This is because bases react with the natural oils on the skin.

3. Many bases, like sodium hydroxide (NaOH), are solid at room temperature, while others, like ammonia, are usually found in a dissolved, liquid form.

4. Bases turn red litmus paper blue.

5. They also turn phenolphthalein (a chemical indicator) pink.

Chemical properties of bases

1. Bases react with acids in the neutralisation reaction of salt and water.

Example.

HCl + NaOH → NaCl + H₂O Acid + Base → Salt + Wate r

2. Bases and alkalis react with ammonium salt on heating to produce ammonium gas.

Example.

NaOH + (NH4)₂SO₄→ Na₂SO₄+ 2NH₃

3. Bases do not react with metals like acids do

4. Bases conduct electricity when dissolved in water.

Activity 1.9 Testing Conductivity of Basic Solution

Materials needed: two small containers (like beakers or cups), distilled water, baking soda or liquid ammonia (as the base), battery, wires, light bulb or LED, stirring stick and measuring spoon.

Steps:

1. Container one: Fill it with pure water.

2. Container two: Fill it with water, then add a spoonful of baking soda or a few drops of liquid ammonia. Stir until it dissolves.

3. Connect a simple circuit using a battery, wires, and a small light bulb (or LED).

4. Make sure the setup works by testing it with a piece of metal to ensure the bulb lights up.

5. Place the ends of the wires that are connected to the light bulb into the container with pure water.

6. Observe what happens. Example, the light bulb should not light up because pure water does not conduct electricity well.

7. Place the ends of the wires into the container with the baking soda or ammonia solution.

8. Observe what happens. Example, the light bulb should light up, showing that the solution can conduct electricity.

9. Record your observations in your notebook.

10. Research and explain the observation you recorded above: Write down your findings in your notebook Uses of Base in Everyday Life

1. Baking soda (Sodium Bicarbonate) - used in cooking, cleaning, and as an antacid.

2. Ammonia - found in glass and surface cleaners.

3. Washing soda (Sodium Carbonate) - used in laundry detergents.

4. Soap - for washing hands, dishes, and clothes.

5. Detergents - used for cleaning clothes and surfaces.

6. Toothpaste - helps clean and protect teeth.

7. Milk of Magnesia - a medicine that neutralises stomach acid.

8. Shampoo - cleans hair and can have mild basic properties.

9. Chlorine Bleach - disinfects and whitens clothes.

10. Antacids - relieve indigestion and heartburn.

11. Chalk (Calcium Carbonate) - sometimes used for writing or as a mild abrasive.

Water Role of water in the dissociation of acids and bases In our previous activities, six and nine, water played a role in both acidic and basic solutions, making it good electrical conductivity.

Now, let us examine the role of water in the dissociation of the bases and acids!

1. Water plays a key role in the dissociation of acids and bases, which means breaking them apart into smaller parts called ions. When an acid, like hydrochloric acid (HCl), is added to water, it separates into hydrogen ions (H+) and chloride ions (Cl–). These hydrogen ions give acids their sour taste and make them conduct electricity.

Similarly, when a base, like sodium hydroxide (NaOH), dissolves in water, it splits into sodium ions (Na+) and hydroxide ions (OH–). These hydroxide ions make the solution feel slippery and help it conduct electricity. Hence, without water acids and bases would not release these ions, so they would not show their usual acidic or basic properties. This is why water is important;

it allows acids and bases to dissociate, letting them act in their usual ways.

— See Figure 1.4.

Figure 1.4: Dissociation of Acid and Base in Water For example, when a dry strip of blue litmus paper is brought near a test tube containing dry hydrogen chloride (HCl) gas, it won’t change colour. This shows that there are no hydrogen ions (H+) present in the dry gas. However, if you moisten the litmus paper with water and then bring it near the gas, the paper turns red. This happens because HCl gas dissolves in water to form H+ ions, which makes it acidic. Without water, these ions don’t form. The same thing happens with bases. If you take a dry piece of solid sodium hydroxide (NaOH) and bring a dry strip of red litmus paper close, there is no colour change. This is because the sodium hydroxide (NaOH) is a hygroscopic compound, which means it can absorb moisture from the air and become wet. When this happens, OH– ions are released, and the red litmus paper turns blue, showing the basic nature of NaOH. Without water, the OH– ions are stuck in the solid and don’t show any basic properties.

2. Heat-releasing dissolution (Exothermic reaction): When an acid (like sulfuric acid) or a base (like sodium hydroxide) is dissolved in water, the solution becomes hot. This means that the dissolution process releases energy (it is exothermic). A part of this energy is used to break the bonds in the acid or base, which releases H+ or OH– ions into the water, allowing the substance to show its acidic or basic nature. — See Figure 1.5.

Figure 1.5: Exothermic reaction Similarities Between Acids and Bases

1. Conduct Electricity: When dissolved in water, bases and acids both can conduct electricity. The reason for this is that they separate into charged particles called ions, which enable electric current to flow through the

solution. For instance, when in water, sodium hydroxide (NaOH) and hydrochloric acid (HCl) both conduct electricity.

2. Corrosive Nature: Because acids and bases may both be corrosive, they can wear down or harm some things, such as skin and metal. If they come into touch with the skin, powerful bases (like sodium hydroxide) and acids (like sulphuric acid) can burn.

3. Reaction with Water: Both acids and bases dissolve in water to show their properties. Acids release hydrogen ions (H+), and bases release hydroxide ions (OH–) when they dissolve.

4. Ability to change colour of indicators: Acids and bases can change the colour of chemical indicators (like litmus paper or universal indicators).

Acids turn blue litmus paper red, while bases turn red litmus paper blue.

Both can cause colour changes in universal indicators, showing different pH values.

5. React with each other: Acids and bases can react with each other in a process called neutralization, producing salt and water. For example, when hydrochloric acid (HCl) reacts with sodium hydroxide (NaOH), they form sodium chloride (table salt) and water (H₂O).

Salts

1. Which main ingredient brings out the taste in food?

2. What other ways do you use salt in everyday life?

A salt is an ionic compound formed when an acid reacts with a base. In the reaction, the H+ of an acid is replaced by another positively charged particle known as a cation derived from a base while the OH– of a base is replaced by another negatively charged particle known as an anion derived from an acid. Such a reaction is called neutralisation reaction.

Apart from acids and bases, a wide variety of compounds exist as salt (see Figure 1.6). One salt which is mostly used in the kitchen and as table salt is sodium chloride. It is also referred to as common salt. Table 1.3 shows examples of salts and their common names.

Table 1.3: Some salts and their chemical formula Chemical formula of salt Common names NaCl Sodium chloride KNO₃ Potassium nitrate Zn(NO₃) ₂ Zinc nitrate NH₄Cl Ammonium chloride FeSO₄ Iron (II) sulphate NH₄NO₃ Ammonium nitrate Ca (HSO₄) ₂ Calcium Hydrogen sulphate CuSO₄ Copper sulphate CuCl₂ Copper (II) chloride CuCl Copper (I) chloride

Figure 1.6: Samples of Different Types of Salt

Formation of Salts

Salts are formed in reactions involving acids and bases. Such reactions result, typically in the formation of salt and water. Other products such as carbon dioxide may also be formed.

A. Formation of salt by neutralisation: A neutralisation reaction is the reaction between an acid and a base to produce salt and water. The general equation of a neutralisation reaction between the acid HA and base BOH is as follows:

• Acid + Base Salt + Water

• HA + BOH AB + H₂O where AB is the salt formed.

Formation of water in addition to the salt makes the reaction a neutralisation reaction.

The specific reaction of the activity above is as follows:

• HCl_((aq))+ NaOH_((aq)) NaCl_((aq))+ H₂O₍ₗ₎ In terms of ions, the equation is as follows:

• H+ + Cl– + Na+ + OH– Na+ Cl– + H₂O From the above ionic equation, the positive ion or cation of the salt, which is Na+, comes from the base. Therefore, it is called the base radical. The negative ion or anion of the salt, which is Cl–, comes from the acid. Therefore, it is called the acid radical.

How do you think the process of neutralisation can be used to help the human body?

Antacids These are chemicals which are basic and are used to relieve the body of congestion caused by accumulation of acid in the stomach. The congestion is normally caused by indigestion or eating late evening or at night. The basic chemical, when drank, neutralises the acid to produce a salt. This brings relief to the body.

Antacids are sold with commercial names such as Magacid, Gastracid, Andrews Liver Salt, Gastrone etc.

Figure 1.7: Samples of Antacids

Activity 1.10 Formation of Salt by Acid and Base Materials needed: vinegar (acetic acid solution), baking soda (Sodium bicarbonate), two clear glass or plastic cups, stirring rod or spoon and pH paper or universal indicator paper (optional, for pH testing) Steps:

1. Pour a small amount of vinegar (acetic acid solution) in one cup and label A. This will act as your acid solution.

2. Add a small baking soda (Sodium bicarbonate) to the other cup and label B. This will act as your base.

3. Slowly pour the base (baking soda) into the acid solution (vinegar).

4. Stir the mixture gently with a stirring rod or spoon.

5. Briefly explain why precautions must be observed when performing a science experiment.

What to expect/observe

1. Bubbling or fizzing of the reaction mixture indicates the release of carbon dioxide gas. This is a characteristic neutralisation reaction between an acid and a carbonate or hydrogen carbonate.

2. The mixture feels slightly warmer since neutralisation reactions release a little heat into their surroundings. After mixing thoroughly, the final

solution contains sodium acetate (salt), water, and carbon dioxide gas:

Test for the salt

1. Test for the pH of the vinegar and baking soda solution by dipping a pH paper or universal indicator into each of them.

2. Test for the pH of the final solution Hint After neutralisation, the pH should be closer to neutral (around 7), indicating the formation of a salt solution. For the pH to be exactly 7, combine equal molar amounts of the acid and base.

Safety Precautions

1. Even though the vinegar and baking soda are relatively safe household chemicals, protective clothing must be worn to protect the body from their possible spillage.

2. Any spilled vinegar or baking soda must be cleaned immediately.

3. Since the experiment will produce carbon dioxide gas, it must be performed in a well-ventilated area to ensure the presence of enough oxygen for respiration.

B. Formation of salt by reaction of acids with metals: In a reaction between a dilute acid and some metals, a salt is produced along with hydrogen gas.

• Metal + Acid → Salt + Hydrogen

• Zn(s) + H₂SO4(aq) → ZnSO4(aq) + H₂(g)

Activity 1.11 Reaction Between Dilute H₂SO₄and Zinc Metal.

Instruction: Perform this activity under supervision of a teacher in a school laboratory.

Materials needed: test tube, zinc granules, dilute H₂SO₄, candle or match box, test tube rack or retort stand with clamp, rubber bung with delivery tube, soap solution in a basin.

Steps:

1. Place a test tube in its rack or in a clamp mounted on a retort stand.

2. Put a few zinc granules in the test tube.

3. Add dilute H₂SO₄carefully along the sides of the test tube.

4. Plug the mouth of the test tube with a rubber bung fitted with a delivery tube.

5. Dip the end of the delivery tube in a soap solution placed in a shallow basin.

6. Bring a burning candle or matchstick near the mouth of the basin.

Figure 1.8: Reaction between dilute H2SO4 and zinc What to see/observe:

i. Bubbles form in the soap solution, indicating that a gas has been released.

ii. When a burning candle or matchstick is brought near the mouth of the basin, the gas in the soap bubbles burns with a ‘pop’ sound. This confirms that the gas that evolved is hydrogen gas.

Conclusion: The equation of the reaction is H₂SO₄(aq) + Zn(s) ZnSO₄(aq) + H₂(g) The salt formed is zinc sulphate.

Safety Precautions

1. Wear hand gloves and eye goggles to protect the hand and eyes.

2. Wear protective clothes.

3. Clean any spillage of acid.

C. Formation of salt by reaction of acids with metal carbonates and hydrogen carbonates: Reactions between acids, metal carbonates, or metal hydrogen carbonates (bicarbonates) produce salt together with water and carbon dioxide.

• Metal carbonate or Metal hydrogen carbonate + Acid Salt + water + carbon dioxide

• CaCO₃(s) + 2HCl(aq) CaCl₂(aq) + H₂O(l) + CO₂(g)

• NaHCO₃(aq) + HCl(aq) NaCl(aq) + H₂O(l) + CO₂(g)

Activity 1.12 Acids React with Carbonates, Releasing CO₂

Materials needed: test tube, boiling tube fitted with cork, thistle funnel, delivery tube, sodium carbonate, sodium hydrogen carbonate, dilute HCl and freshly prepared lime water.

Steps:

1. Place about 0.8 g sodium carbonate in a test tube.

2. Obtain about 12ml of a freshly prepared dilute HCl.

3. Mix them by stirring until they are thoroughly mixed.

4. Plug the test tube with a rubber bung fitted with a thistle funnel and delivery tube.

5. Pour about 2 ml of freshly prepared lime water in a test tube.

6. Dip the delivery tube into the lime water.

Figure 1.9: Reaction between HCl and Na₂CO3 What to see:

1. There is formation of bubbles in the lime water, indicating the release of a gas.

2. The lime water turns milky, confirming the release of CO₂.

The equation of the reaction is as follows:

• 2HCl(aq) + Na₂CO₃(s) 2NaCl(aq) + H₂O(l) + CO₂(g)

• The salt formed is sodium chloride.

Conclusion: To prove that sodium chloride is the salt formed, litmus paper or a universal indicator can be used to test for the pH of the resultant reaction

solution.

Sodium chloride solution is neutral to indicators, hence there will be no colour change of the indicator.

When sodium hydrogen carbonate is used, the same observations will be recorded. However, the equation of the reaction will be as follows:

• HCl(aq) + NaHCO₃(s) NaCl(aq) + H₂O + CO₂(g) Types of Salt

Table 1.5: Types of salt Types of salt Description Normal salt A salt which is neutral to an indicator. That is to say it neither shows acidic nor basic properties. It is formed from the reaction between a strong acid and a strong base, resulting in a pH of 7. They include chlorides, nitrates, sulphates of metals.

Examples are NaCl, K₂SO₄, CaSO₄, Mg₃(NO3)₂, MgCl₂ Acidic salt Salt which shows acidic properties with an indicator. It is formed from the reaction between a strong acid and a weak base, resulting in a pH lower than 7. Examples of acid salts are NH₄Cl, NaHSO₄, CuSO₄, FeCl₂etc Basic salts Salt which shows basic properties with an indicator. It is formed from the reaction between a weak acid and a strong base, resulting in a pH higher than 7. Examples are NaHCO₃, KHCO₃, CaCO₃, Na₂CO₃etc.

Why is a normal salt neutral to an indicator? Record your findings in your notebook.

Properties of Salts

1. Soluble salts dissolve in water to form solutions.

NaCl(s) + H₂O NaCl(aq)

2. Carbonates and hydrogen carbonates react with acids to form salt, water and CO₂.

CaCO₃(s) + 2HCl(aq) CaCl₂(aq) + H₂O(l) + CO₂(g)

3. Bases react with ammonium salts to form another salt, water and ammonia.

NH4Cl(s) + NaOH(aq) NaCl(aq) + H₂O(l) + NH₃(g)

4. Salts decompose on heating to release oxygen gas.

2NaNO₃ 2NaNO₂+ O₂

5. Neutral/normal salts do not affect an indicator such as litmus paper, litmus

solution or universal indicator.

Uses of salts

Activity 1.13 Uses of salt Write down at least five uses of salt in everyday life activities in your notebook.

Table 1.6: Uses of salts Use Explanation Seasoning and flavouring.

Enhances the taste of food.

Preservation. Draws water out of food items and inhibits bacterial growth on them.

Industrial uses. As a raw material to produce chlorine and sodium hydroxide, in the manufacture of soap and other chemical processes.

Treatment of water. To soften hard water to make it suitable for use.

Textile industry. As a fixative to help dyes adhere to fibres to produce desirable colour patterns.

Fertilisers. As an essential component of fertilizers to produce macro and micronutrients.

Use Explanation

De-icing. For removing snow from roads, sidewalks and other places where snow may cause slipping and other forms of inconvenience.

Personal care. To make skin care products due to its cleansing properties.

Household cleanup To clean stains, sanitise refrigerators, erase watermarks and brighten brass and copper ware.

In the garden To eliminate the activities of ants, slugs, worms, ivy plants and to control weeds.

Figure 1.10: Some uses of salt

Activity 1.14 Exploring the Types of Solution

The diagrams below show a mixture of solute and solvent to form a

solution. Observe them carefully and discuss the questions with your friends.

Figure 1.11: Formation of solution

1. Describe your observation from Fig. 1.6 above

2. What happens to the solution’s appearance if you add more solute?

3. How will you differentiate between solute and solvent to create a uniform

solution? Give one example from the diagram

4. Using the diagrams provided, list one example of each type of solution (gaseous, liquid, and solid) Is there any difference between solid solution, liquid solution and gaseous

solution? Let us find out!

• A solid solution is a homogeneous mixture in which the solvent is solid, and the solute is either gas, liquid, or solid. The components are dispersed uniformly at the molecular level. A solid solution is an alloy, such as bronze, in which copper acts as the solvent and tin as the solute.

• A liquid solution is made up of a liquid solvent and a solute, which can be gas, liquid, or solid. These solutions are widespread in everyday situations.

For example, saltwater is a liquid solution in which the solvent is water, and the solute is salt.

• A gaseous solution is generated when gases combine with one another to form a homogeneous mixture. In this scenario, the solvent is a gas, while the solute can be gas, liquid, or solid. Air, for example, is a mixture of nitrogen as a solvent and oxygen as a solute, among other gases.

Activity 1.15 Identifying Acid-Base Concentration Using the pH Scale Materials needed: Use the figure 1.12 to perform this activity

Figure 1.12: Acidity or alkalinity of some household items Steps:

1. Explain the pH scale (0-14) and its significance in identifying acidity (below 7 is acidic, 7 is neutral, above 7 is basic).

2. Predict which household liquids are acidic based on flavour and prior knowledge.

3. In small groups, test the pH of various liquids using the indicator solution or test strips,

4. Create different concentrations of a base (e.g., baking soda).

5. Measure pH and discuss how concentration affects pH values

6. Record your observations.

7. Share your findings and discuss the strongest acid identified and the pH values recorded.

Questions: Record your answers in your notebook for future reference.

Use the pH scale in figure 1.12 to help you answer these questions:

1. What does the pH scale tell us about acids and bases?

2. Which household liquids do you think are acidic, and why?

3. What pH results did you find for each liquid, and how did they match your predictions?

4. How does changing the concentration of baking soda affect its pH level?

5. Why is it useful to know the pH of common household items?

6. What other indicators can be used to identify the strength of an acidic or alkaline solution?

Hint Here is a quick overview of some key points to keep in mind from the activity pH scale ranges from 0 to 14; where below 7 is acidic, 7 is neutral, and above 7 is basic. Vinegar, lemon juice, and soda are all examples of sour foods.

Higher quantities basic substances increase the pH, making it more basic.

Knowing PH of substances helps with the safe usage and reactions of household products in everyday life.

Activity 1.16 Preparation of Sugar Solution

Materials needed: granulated sugar (sucrose, C₁₂H₂₂O₁₁), distilled water, balance/ scales, volumetric flask or beaker and stirring rod.

Steps:

1. Calculate the Required Mass of Solute: Decide on the solution’s molarity (M) and volume (V).

• Use the molar mass of sucrose (342.3 g/mol) to calculate the mass

• Amount substance in moles (n) = M × V

• Mass = Moles × Molar Mass

• For example, to prepare 1 dm³

• of a 0.5 mol/dm³

• sucrose solution:

• Moles = 0.5 mol/dm³× 1 dm³= 0.5 mole

• Mass = 0.5 moles × 342.3 g/mol = 171.15 g

2. Measure 171.15 g of sucrose.

3. Add the sugar to a volumetric flask or beaker. Add distilled water gradually while stirring until the total volume reaches 1 litre.

4. Stir until all the sugar is completely dissolved.

5. Label the container with the concentration and date.

Questions: Put your ideas in your notebook.

1. How can you tell when the sugar has fully dissolved?

2. How does the concentration of sugar solution affect its taste and properties?

3. Why is it important to accurately measure the amount of sugar when preparing a solution?

Activity 1.17 Preparation of a Solution

Materials needed:

• Solute (the substance being dissolved)

• solvent (the liquid in which the solute dissolves)

• a balance/ scales (to measure the mass of the solute accurately)

• volumetric flask (a flask with a marking that indicates the final volume of the solution)

• graduated cylinder (to measure the solvent accurately)

• A stirring rod or magnetic stirrer (to facilitate dissolving the solute)

• distilled or deionized water (recommended as solvent for most chemical solutions)

• safety glasses and gloves (for handling chemicals) Steps:

1. Calculate the amount of solute needed and the volume of solution you want to prepare (in litres, L).

2. Use the following formula to calculate the mass (m) of solute required:

m = C × V × M Where:

m = Mass of solute in grams (g) C = concentration of the desired solution (mol/dm³) V = Volume of the solution in dm³M = molar mass

3. Carefully weigh out the calculated mass of the solute using a balance.

4. Rinse the volumetric flask with a small amount of solvent to remove impurities.

5. Using a funnel, carefully transfer the weighed solute to a volumetric flask and add a little amount of solvent, gently swirling to aid in dissolving.

6. Rinse the flask with solvent to ensure that all the solute has been transferred, then gradually add distilled water while swirling until it reaches the graduation mark.

7. Using a dropper, adjust the solvent level drop by drop until the bottom of the meniscus matches the mark, then close the flask with a stopper and invert several times to ensure a complete mix.

8. Record your observation into your science notebook for discussion with your friends Safety Precautions

1. Always use safety goggles, gloves, to protect yourself.

2. Read labels and use funnels or pipettes to avoid spills.

3. Familiarize yourself with the location of safety equipment and protocols in the science laboratory.

Questions: Record your suggested answer in notebook

1. What is the purpose of rinsing the weighing boat or container after transferring the solute to the volumetric flask?

2. Why is it important to add distilled water gradually while swirling the flask?

3. How do you ensure that the bottom of the meniscus is at the correct graduation mark when regulating the volume?

4. What would happen if you filled the volumetric flask beyond the graduation mark?

5. Why is it necessary to reverse the flask several times after sealing it?

How Will You Calculate the Molar Mass of a Solute?

Molar mass is the mass of one mole of a substance, naturally expressed in grams per mole (g/mol). To calculate the molar mass of a solute, you need to sum the atomic masses of all the atoms in its chemical formula. Use the steps below to calculate the molar mass of a compound:

1. Identify and write down the chemical formula of the compound (e.g., H₂O for water)

2. Break down the compound into its constituent elements. For H₂O, the elements are hydrogen (H) and oxygen (O).

3. Atomic mass of each element can be obtained from the periodic table. See

figure 1.13

Figure 1.13: Periodic table Atomic mass of Hydrogen (H) = 1.01 g/mol Atomic mass of Oxygen (O) = 16.00 g/mol

4. Multiply the atomic mass of each element by the number of times it appears in the formula. Finally, add the total masses of all the elements in the compound.

Worked example 1

Calculate the molar mass of water

• Chemical Formula H₂O:

• Identify the elements in the compound (e.g., H₂O = H and O).

• For each element, multiply its atomic mass by the number of atoms present in the formula. H: 1.01 g/mol × 2 = 2.02 g/m and O: 16.00 g/mol × 1 = 16.00 g/mol

• Total molar mass of H₂O = 2.02 g/mol + 16.00 g/mol = 18.02 g/mol

Worked example 2

Calculate the molar mass of glucose (C₆H₁₂O₆)

• Chemical Formula: C₆H₁₂O₆

• Find the Atomic Masses:

Carbon (C): 12.0 g/mol Hydrogen (H): 1.0 g/mol Oxygen (O): 16.0 g/mol

• Multiply by the Number of Atoms:

= Molar Mass of C₆H₁₂O₆ = (72.06 x 6) + (12.12 x 12) + (96.00 × 6) = 180.18 g/mol

Worked example 3

Calculate the molar mass of calcium chloride (CaCl₂), follow these steps:

• Identify the Elements: Calcium (Ca) and Chlorine (Cl)

• Find atomic masses:

Calcium (Ca) = 40.08 g/mol Chlorine (Cl) = 35.45 g/mol

• Multiply by number of atoms Calcium: 40.08 g/mol × 1 = 40.08 g/mol Chlorine: 35.45 g/mol × 2 = 70.90 g/mol

• Total molar mass of CaCl₂= 40.08 g/mol + 70.90 g/mol = 110.98 g/mol Practice the following to check your understanding!

Calculate the molar masses of the following:

a. glucose (C₆H₁₂O₆)

b. sodium chloride (NaCl).

Next, you will examine how to prepare dilute solutions and concentrated solutions. To help you with your activity, know that concentration tells us how much solute is in a solution and Molarity (M) = moles of solute over litres of

solution.

Activity 1.18 Concentration of Solutions

Materials needed: distilled water, solute (like salt or sugar), volumetric flasks, beakers, graduated cylinders, pipettes, analytical balance and stirring rods Steps:

1. Preparing a dilute solution:

a. Calculate how much solute you need (e.g., for 1 L of 0.1 M salt).

b. Weigh the amount using an analytical balance (e.g., 5.844 g for NaCl).

c. Dissolve it in some distilled water in a volumetric flask.

d. Add more distilled water until you reach the 1 L mark on the flask.

e. Label the flask with concentration and date.

2. Preparing a concentrated solution:

a. Calculate how much solute you need (e.g., for 500 mL of 2M salt).

b. Weigh out the required amount (e.g., 58.44 g for NaCl).

c. Dissolve it in less water than needed at first, then transfer to a volumetric flask and add water up to the 500 mL mark.

Safety Precautions

1. Wear gloves, and goggles.

2. Know where safety equipment is located

3. Label appropriately Imagine after your activity, you are to observe the diagram in figure 1.14 and then put your ideas down as you read

Figure 1.14: Dilute and concentrated solutions Use figure 1.14 to explore the following questions:

1. Which is a concentrated solution and why?

2. Which is a diluted solution and why?

3. In your own words, how would you explain the terms dilute and concentrated solutions to someone who has never learned about solutions?

4. Why it’s important to know the exact concentration of a solution?

5. What differences did you notice between the dilute, medium, and concentrated solutions you prepared? How did these differences help you understand the concept of concentration?

6. Why do you think it is important to quantify the concentration of a

solution?

7. If you were to create a new solution, what factors would you consider when determining how concentrated it should be?

What are the Key Characteristics You Should Understand About Concentration?

Concentration (C) is the amount of a substance in a certain volume of solution.

We can describe the concentration of the solution in different ways.

Let us investigate the different ways of describing concentration of solution with worked examples and practice the questions to check your understanding!

A. Mass percentage (w/w): The mass percentage of a component of a solution is defined as: Mass % of a component = Mass of the component solution (g)_________________________ Total mass of the solution (g) × 100 Question: Choose the correct answer What is the mass percentage of a component in a solution if the mass of the component is 25 grams and the total mass of the solution is 200 grams?

A. 10% B. 12.5% C. 12.5% D. 15% B. Volume percentage (v/v): The volume percentage is defined as:

= Volume of solute____________________ Total volume of the solution × 100 Question: What is the volume % of rose extract in a solution prepared by dissolving 14.0 cm³rose extract in a solvent to make 200 cm³of solution?

Solution:

Volume percentage = Volume of solute____________________ Total volume of the solution × 100 = 14.0cm³________ 200.0 cm³× 100 = 7.0 % Rose solution (v/v) Choose the correct answer

1. An SHS General science learner prepares a solution by mixing 60 mL of acetic acid with enough water to make a total volume of 500 ml. The value of the volume percent (v/v%) concentration of acetic acid in the

solution is A. 10% B. 12% C. 15% D. 20% C. Mass by volume percentage (w/v): The mass-by-volume percentage is another unit commonly used in medicine and pharmacy.

Formula = Mass of Solute (g)____________________ Volume of the solution(mL) × 100

Worked Example

What is the mass/volume % of glucose solution prepared by dissolving 50 g glucose in enough water to make 1000 cm³of solution?

Solution:

= Mass of Solute (g)_____________________ Volume of the solution(cm³) × 100 = 50.0g/1000 cm³× 100 = 5.0 % glucose (w/v) Question: Write your answers in your notebook.

Determine the mass/volume % of glucose solution prepared by dissolving 30 g glucose in enough water to make 900 cm³of solution?

D. Molarity (concentration in moles per dm³): Concentration (C) expresses the moles of solute in dm³of solution. The most common solution concentration unit used in chemistry is concentration (C).

If you have the mass of the solute, you can calculate the moles using the molar mass of the solute. The formula is:

number of moles of a solute (n) = mass of solute(n)__________________ molar mass of solute( M) Measure the solution’s volume in dm³. If you have the volume in cm³, convert it to dm³by dividing it by 1000.

Calculate the concentration:

Use the formula to find the molarity.

Concentration (C) = amount of solute (n)_____________________ Volume of solution in dm³(V)

Worked example

Calculate the concentration of a solution made by dissolving 0.0974 moles of NaCl in 1.5 dm³of water.

Concentration (C) = number of moles_____________________ Volume of solution in dm³(V) = 0.0974 moles/1.5 dm³= 0.065 mol / dm³Questions Choose the correct answer as you read the questions

1. You have 5 grams of NaCl (sodium chloride) and the molar mass of NaCl is 58.44 g/mol. What number of moles of NaCl do you have?

A. 0.0855 moles B. 0.0810 moles C. 0.0825 moles D. 0.0785 moles

2. How do you convert volume from cm³ to dm³?

A. Add 1000 B. Divide by 1000

C. Multiply by 1000

D. Subtract 1000

Activity 1.19 Dilution process Materials needed: funnel, beaker, dropping pipette, wash bottle, volumetric flask and water.

Figure 1.15: Set-up for dilution Steps:

1. Weigh the container containing the requisite mass of solid using a precise balance. Transfer the solid to a beaker, reweigh it, and calculate the mass difference.

2. To dissolve the solute entirely, add 100 cm³ of distilled water to the beaker and stir with a glass rod. If required, gradually boil the solution.

3. Fill a 250 cm³ volumetric flask with the solution using a funnel.

4. Rinse the beaker, glass rod, and funnel to remove any residue.

5. Fill the flask to the mark with distilled water, mix thoroughly by inverting the flask several times, and then use a volumetric pipette to measure 25.0 cm³ of the original solution into another 250 cm³ volumetric flask.

Reflect on the following questions from Activity 1.20.

1. What is the purpose of weighing the container before and after adding the solid?

2. How do you ensure the solid is completely dissolved in the distilled water?

3. Why is it important to rinse the beaker, glass rod, and funnel during the transfer process?

4. What steps should be taken to accurately dilute the original solution using a volumetric pipette?

5. How does inverting the flask multiple times help in reaching a homogeneous solution?

Dilution of Solution

Diluting solutions involves raising the volume of a solution while decreasing the concentration without changing the number of moles.

Since the number of moles remains constant, we can use the concentration equation to generate a dilution equation: C₁V₁= C₂V₂ C₁= Original Concentration V₁= Original Volume C₂= New Concentration V₂= New Volume

Worked example

1. 70 cm³of water is added to 240 cm³of a 0.80 mol dm⁻³solution of KOH.

Calculate the concentration of the diluted solution.

• Using C₁V₁= C₂V

• V2 = V1 + volume of water = 250 + 70 = 320 cm³

• Making C2 the subject = C1V1_____ V2 = 0.80mol dm − 3 × 240cm³_____________________ 320cm³= 0.6 mol dmmol /dm³PH MEASUREMENT In our previous lesson, we learnt the various pH levels of acidic and basic solutions.

Try these questions: Write your answer in your notebook

1. Why do you do think it is important to know the pH of acidic and basic

solution?

2. What do the pH levels indicate about the properties of these substances?

3. How might these substances affect the environment or health based on their pH?

Importance of pH Measurement

1. The pH scale is essential for determining whether a solution is acidic, neutral, or basic. This knowledge is vital for conducting experiments and studying chemical reactions in various scientific fields.

2. Monitoring pH levels in natural water bodies, such as rivers and lakes, helps assess their health. Acidic waters can harm aquatic ecosystems, while excessively basic waters can hinder plant growth. Keeping track of pH is crucial for environmental conservation.

3. Soil pH influences the availability of nutrients for plants. Farmers use pH measurements to optimize soil conditions for specific crops. For example, some crops thrive in acidic soil, while others prefer neutral or alkaline conditions.

4. In the medical field, pH measurement is essential for understanding various bodily functions. For instance, human blood must maintain a specific pH range (7.35-7.45), as deviations can indicate health problems. Additionally, antacids are used to neutralize stomach acidity, and their effectiveness can be monitored through pH levels.

5. Many industries, including food production and pharmaceuticals, rely on pH monitoring to ensure product quality and safety. For example, the pH of food can significantly affect its flavour, preservation, and overall safety.

pH Paper: a special type of paper used to test how acidic or basic a solution is. It is treated with certain chemicals that change colour when they come into contact with liquids of different pH levels. It provides a quick way to find out the acidity or basicity of a solution without needing complex equipment. pH paper is commonly used in schools, laboratories, and even at home (like testing soil or pool water). The pH scale ranges from 0 to 14, where:

a. 0 to 6 indicates acidic solutions (like lemon juice).

b. 7 is neutral (like pure water).

c. 8 to 14 indicates basic or alkaline solutions (like baking soda).

Figure 1.16: pH paper

Activity 1.20 Measuring the pH of Substances Using pH Paper Materials needed: vinegar, lemon juice, baking soda solution, tap water and pH paper Steps:

1. Pour the substances each into each beaker and label them A, B, C, and D.

2. Take a strip of pH paper and dip it into each solution one at a time. Make sure to remove the paper after a few seconds.

3. Observe colour changes.

4. Use a pH colour chart (which usually comes with the pH paper) to match the color of the paper to the corresponding pH value.

5. Record your observation in the table below:

Substance Colour change pH value Tap water Lemon juice Baking soda solution Vinegar Questions: Write down your answers in your notepad.

1. From activity 22, list the acidic and basic solutions.

2. Explain your answer in (1) above.

pH Meter: an electronic device used to measure the acidity or basicity of a

solution. It provides a precise pH reading, which indicates how acidic (pH less than 7) or basic (pH greater than 7) a substance is.

Table 1.7: Types of portable pH meter Pen testers

Figure 1.17: Pen tester Pen testers are inexpensive pH meters the size of a pocketbook. Pen testers have many uses in the building, hydroponics, food production, and pool or spa care industries.

Handheld meters

Figure 1.18: Handheld meter Handheld meters often have a more robust build and a slightly larger shape than pen testers.

With this design, the electrode is constructed independently of the meter. Hand-held meters are designed for usage in the field. Environmental officers use them in field research, aquaculture, agriculture, and water treatment.

Benchtop pH meters

Figure 1.19: Benchtop pH mater The largest of the three pH meter categories are benchtop meters. They can be put on a wall or a desk. They are often the most accurate pH meters, making them ideal for laboratory and professional use. Benchtop pH meters are frequently used in laboratories.

Table 1.8: Types of pH meters based on Usage Laboratory pH meter

Figure 1.20: Laboratory pH meter It has a large measuring range, is highly accurate, and is versatile.

Industrial pH meter

Figure 1.21: Industrial pH meter A pH meter is a precise tool that measures how acidic or basic a solution is. It combines analogue and digital features for stable, efficient, and accurate readings. With built- in alarms, it can signal if the pH level goes beyond a set range, making it easier to monitor conditions. Its design helps reduce interference, ensuring accurate results in various environments.

Table 1.9: Types of pH meters based on reading Analog pH meter

Figure 1.22: Analog pH meter An analogue pH meter is the original type of model. A pointer will show the pH level on analogue pH meters. The needle will move toward a number representing the pH level after the measuring electrode has been put into the sample. When using an analogue pH meter, one must be careful to obtain accurate findings. The little pointer is the reason for this.

Digital pH meter

Figure 1.23: Digital pH meter Developed subsequently to analogue meters these have a numerical display or the pH value.

Calibration of pH Meter pH meters need regular calibration to ensure accurate measurements. Calibration involves adjusting the pH meter using standard pH buffer solutions (pH 4.01, pH 7.00, pH 10.01) to set accurate reference points across the pH scale.

Calibration of a pH meter adjusts the device to account for any shifts in its readings, called electrode drift. This process ensures that the pH meter provides accurate measurements for different solutions, making sure the results are reliable.

Activity 1.21 Demonstration of How to Use pH Meter Materials needed:

• pH meter (including electrodes: glass pH electrode and reference electrode)

• pH buffer solutions (pH 4.01, pH 7.00, pH 10.01)

• distilled water (for rinsing electrodes)

• sample solutions to measure pH, stirring rod (if needed) and cleaning cloth or tissue Steps:

1. Turn on the pH Meter. If not already calibrated, perform this as follows:

a. Obtain pH 4.01, pH 7.00, and pH 10.01 buffer solutions.

b. Check expiration dates and condition of buffer solutions to ensure accuracy.

c. Rinse the pH electrode with distilled water and blot dry it with clean tissue.

d. Immerse the electrode in the pH 7.00 buffer solution.

e. Allow the reading to stabilise (usually indicated when the display stops changing).

f. Adjust the pH meter according to the manufacturer’s instructions to read pH 7.00.

g. If necessary, rinse the electrode with distilled water and repeat the calibration process with pH 4.01 and pH 10.01 buffer solutions.

h. Confirm calibration success by checking readings against buffer

solution values.

2. Rinse the pH electrode with distilled water and blot dry between measurements.

3. Stir the sample gently to ensure homogeneity (if needed).

4. Immerse the cleaned and calibrated pH electrode into the sample solution.

5. Allow the reading to stabilise (again, indicated when the display stops changing).

6. Record the pH reading displayed on the pH meter.

7. Rinse the electrode with distilled water and dry it between measurements of different solutions.

After Use:

1. Rinse the pH electrode with distilled water to remove any residue from the sample.

2. Blot dry with a clean tissue or cloth.

3. Store the pH meter with the electrode in a storage solution recommended by the manufacturer to keep it hydrated and extend electrode life.

Safety Precautions

1. Handle sample solutions and electrodes carefully to prevent spills or damage.

2. Follow laboratory safety protocols when working with chemicals and solutions.

3. Keep electrodes clean and free from debris or chemical residues.

4. pH electrodes are delicate; avoid touching sensitive parts and handle with care.

5. Regular calibration is performed to maintain accuracy, especially if the pH meter has not been used recently or after prolonged use.

Questions

1. How does pH meter differ from pH paper?

2. From activities 22 and 23, mention some advantages both pH paper and meter.

Advantages of Using pH Meter

1. pH meters provide more precise readings than pH paper, allowing for detailed measurements of pH levels.

2. They can be calibrated to ensure consistent and accurate results, even for repeated measurements.

3. pH meters can be used for various types of solutions, including those that are opaque or coloured, where pH paper might not work well.

4. Digital pH meters provide instant readings, making them convenient and efficient for quick measurements.

Disadvantages of Using a pH Meter

1. Regular calibration and cleaning of the electrodes are needed to keep the device functioning accurately.

2. pH meters are generally more expensive than pH paper, making them less accessible for casual use.

3. The electrodes in pH meters are sensitive and can break easily, requiring careful handling.

4. Digital pH meters need a power source, such as batteries, which may need frequent replacements.

Question: Record your answer in your science notebook.

What happens when you put red dye in bowl or beaker containing a clear water?

Measuring pH Using an Indicator Solution A pH indicator is a special dye that changes colour depending on how acidic or basic a solution is. This allows us to determine if a solution is acidic, neutral, or basic just by looking at the colour. These indicators are usually added to the

solution or soaked into test paper. However, they can only be used with clear and colourless samples because the colour change needs to be seen clearly.

There are two main methods for using pH indicators:

1. Colour Comparison: Dip the indicator into the solution and then compare its colour to a standard colour chart that shows different pH levels. This helps estimate the pH value.

2. pH Test Paper: Soak paper in an indicator solution, then dip it into the test liquid. Compare the colour change to a standard colour chart to see the pH.

Although simple, this method may not be very accurate.

Table 1.10: Types of pH Indicators Universal Indicators A pH indicator can show a wide range of pH values, usually from 1 to 14, by displaying different colours. Each colour corresponds to a specific pH level, allowing users to quickly determine whether a solution is acidic, neutral, or basic. For example, red might indicate a strong acid (low pH), green might indicate a neutral substance (pH 7), and blue or purple might show a strong base (high pH).

Specific Indicators

Some pH indicators are designed to change colour only within a specific, narrower pH range. This allows for more precise detection of small changes in acidity or alkalinity. For example, methyl orange changes colour between pH 3.1 to 4.4, and phenolphthalein does so between pH 8.3 to 10.

Several plants and household chemicals can be used as pH indicators, but in a lab setting, the table below shows the most common chemicals used as indicators:

Table 1.11: Effect of pH indicator on substances.

Examples of indicator Acid Colour Base Colour pH Range Thymol blue (first change) red yellow 1.2 - 2.8 Methyl orange red yellow 3.2 - 4.4 Bromocresol green yellow blue 3.8 - 5.4 Methyl red yellow Red 4.8 - 6.0 Bromothymol blue yellow Blue 6.0 - 7.6 Phenol red yellow Red 6.8- 8.4 Thymol blue (second change) yellow Blue 8.0 - 9.6 Phenolphthalein colourless Magenta 8.2 -10.0

Activity 1.22 Testing the pH of Substances with Natural Indicator Materials needed: turmeric powder, water, filter paper and different solutions Steps:

1. Make a paste of turmeric powder with water.

2. Apply the paste on the filter paper and allow it to dry.

3. Remove the dry powder from the filter paper.

4. Cut the filter paper into small strips.

5. Pour different solutions separately on the strips and note the colour changes.

6. Record the observations in tabular form Observation: Turmeric paper remains yellow in acidic and neutral solutions but turns brown in alkaline solution

Activity 1.23 Testing the pH using universal indicator paper Materials needed: universal indicator paper, dilute NaOH, dilute NH4OH, dilute HCl, dilute H2 SO4, vinegar and distilled water Steps:

1. Take 1cm³of dilute HCl, dilute H2SO4, dilute CH3COOH (vinegar), dilute NaOH distilled

2. water in different test tubes.

3. Add 1.5 cm³of distilled water.

4. Dip a separate universal indicator paper in each tube and match the colour with colour given on the strip.

5. Note the observations in a table below:

Sample Colour of universal indicator paper pH of the solution Question: Write your answer in your notepad.

Name the indicator used in both activities.

1. In a tabular form, classify the following into common household items into acidic and basic substances;

Baking soda, pineapple juice, washing soda, soap detergents, vinegar, toothpaste, milk of Magnesia, chalk, wood ash, Petre, lemon juice, tomato juice, orange juice.

2. Explain the following and give at least two examples each:

a. Base

b. Acid.

3. Explain why acidic or basic solution conduct electricity but distilled water cannot.

4. Describe the role of water in the dissociation of acids and bases.

5. How would you use the concept of neutralization to address the problem of indigestion?

6. Briefly explain the principles underlying a neutralization reaction.

7. Identify the various aspects of everyday life where salt can be used.

8. Calculate the molar mass of sodium oxide (Na2O).

9. You are doing a lab experiment with a sodium chloride (NaCl) solution.

Your teacher gives you a solution with a concentration of 0.1 mol/ dm³. You need to measure a certain amount of this solution for your experiment, and you choose to take 200 cm³. How many moles of NaCl are there in the 200 cm³ of solution you measured?