Chemistry Form 1 Notes : ACIDS, BASES AND INDICATORS

Acids, Bases and Indicators

Introduction to Acids, Bases and Indicators

1. In a school laboratory:

• An acid may be defined as a substance that turns litmus red.
• A base may be defined as a substance that turns litmus blue.

Litmus is a lichen found mainly in West Africa. It changes its colour depending on whether the solution it is in is basic/alkaline or acidic. It is thus able to identify whether another substance is an acid, base, or neutral.

2. Common naturally occurring acids include:

3. Most commonly used acids found in a school laboratory are not naturally occurring. They are manufactured and called mineral acids.

Common mineral acids include:

4. Mineral acids are manufactured to very high concentration. They are corrosive (cause painful wounds on contact with the skin) and react with garments, clothes, and metals.

In a school laboratory, they are mainly used when diluted with a lot of water. This is called diluting. Diluting ensures the concentration of the acid is safely low.

5. Bases are opposite of acids. Most bases do not dissolve in water.

Bases which dissolve in water are called alkalis.

Common alkalis include:

Common bases (which are not alkalis) include:

6. Indicators are useful in identifying substances which look alike.

An acid-base indicator is a substance used to identify whether another substance is alkaline or acidic.

An acid-base indicator works by changing to different colors in neutral, acidic, and alkaline solutions/dissolved in water.

Experiment: To prepare simple acid-base indicator

Procedure

(a) Place some flower petals in a mortar. Crush them using a pestle. Add a little sand to assist in crushing.

Add about 5 cm³ of propanone/ethanol and carefully continue grinding.

Add more 5 cm³ of propanone/ethanol and continue until there is enough extract in the mortar.

Filter the extract into a clean 100 cm³ beaker.

(b) Place 5 cm³ of filtered wood ash, soap solution, ammonia solution, sodium hydroxide, hydrochloric acid, distilled water, sulphuric (VI) acid, sour milk, sodium chloride, toothpaste, and calcium hydroxide into separate test tubes.

(c) Put about three drops of the extract in (a) to each test tube in (b). Record the observations made in each case.

Sample observations

The plant extract is able to differentiate between solutions by their nature. It changes to a similar colour for similar solutions.

• Since lemon juice is a known acid, then sulphuric (VI) and hydrochloric acids are similar in nature with lemon juice because the indicator shows similar colors. They are acidic in nature.
• Since sodium hydroxide is a known base/alkali, then the green colour of indicator shows an alkaline/basic solution.
• Since pure water is neutral, then the orange colour of indicator shows neutral solutions.

7. In a school laboratory, commercial indicators are used. A commercial indicator is cheap, readily available, and easy to store. Common indicators include: Litmus, phenolphthalein, methyl orange, screened methyl orange, bromothymol blue.

Experiment: Using commercial indicators to determine acidic, basic/alkaline and neutral solutions

Procedure

Place 5 cm³ of the solutions in the table below. Add three drops of litmus solution to each solution.

Repeat with phenolphthalein indicator, methyl orange, screened methyl orange, and bromothymol blue.

Sample results

From the table above, the colour of indicators in different solutions can be summarized.

The Universal Indicator

The universal indicator is a mixture of other indicator dyes. It uses the pH scale, which shows the strength of bases and acids. The pH scale ranges from 1 to 14. These numbers are called pH values:

• pH values 1, 2, 3 show a substance is strongly acidic.
• pH values 4, 5, 6 show a substance is weakly acidic.
• pH value 7 shows a substance is neutral.
• pH values 8, 9, 10, 11 show a substance is a weak base/alkali.
• pH values 12, 13, 14 show a substance is a strong base/alkali.

The pH values are determined from a pH chart. The pH chart is a multicolored paper with each colour corresponding to a pH value, i.e.:

• Red corresponds to pH 1, 2, 3 showing strongly acidic solutions.
• Orange/yellow correspond to pH 4, 5, 6 showing weakly acidic solutions.
• Green corresponds to pH 7 showing neutral solutions.
• Blue corresponds to pH 8, 9, 10, 11 showing weakly alkaline solutions.
• Purple/dark blue correspond to pH 12, 13, 14 showing strong alkalis.

The universal indicator is available as:

• Universal indicator paper/pH paper
• Universal indicator solution

When determining the pH of an unknown solution using:

• pH paper: dip the pH paper into the unknown solution. It changes to a certain colour. The new colour is matched to its corresponding one on the pH chart to get the pH value.
• Universal indicator solution: add about 3 drops of the universal indicator solution into about 5 cm³ of the unknown solution in a test tube. It changes to a certain colour. The new colour is matched to its corresponding one on the pH chart to get the pH value.

Experiment: To determine the pH value of some solutions

1. Place 5 cm³ of filtered wood ash, soap solution, ammonia solution, sodium hydroxide, hydrochloric acid, distilled water, sulphuric (VI) acid, sour milk, sodium chloride, toothpaste, and calcium hydroxide into separate test tubes.
2. Put about three drops of universal indicator solution or dip a portion of a piece of pH paper into each. Record the observations made in each case.
3. Compare the colour in each solution with the colors on the pH chart provided. Determine the pH value of each solution.

Sample observations

Note

1. All the mineral acids hydrochloric, sulphuric (VI), and nitric (V) acids are strong acids.
2. Two alkalis/soluble bases, sodium hydroxide and potassium hydroxide, are strong bases/alkalis. Ammonia solution is a weak base/alkali. All other bases are weakly alkaline.
3. Pure/deionized water is a neutral solution.
4. Common salt/sodium chloride is a neutral salt.
5. When an acid and an alkali/base are mixed, the final product has pH 7 and is neutral.

Properties of acids

(a) Physical properties of acids

1. Acids have a characteristic sour taste.
2. Most acids are colourless liquids.
3. Mineral acids are odorless. Organic acids have characteristic smells.
4. All acids have pH less than 7.
5. All acids turn blue litmus paper red, methyl orange red, and phenolphthalein colourless.
6. All acids dissolve in water to form an acidic solution. Most do not dissolve in organic solvents like propanone, kerosene, tetrachloromethane, petrol.

(b) Chemical properties of acids

1. Reaction with metals

All acids react with reactive metals to form a salt and produce hydrogen gas.

Metal + Acid → Salt + Hydrogen gas

Experiment: Reaction of metals with mineral acids

1. Place 5 cm³ of dilute hydrochloric acid in a small test tube. Add 1 cm length of polished magnesium ribbon. Stopper the test tube using a thumb. Light a wooden splint. Place the burning splint on top of the stoppered test tube. Release the thumb stopper. Record the observations made.
2. Repeat the procedure in (1) using zinc granules, iron filings, copper turnings, aluminum foil in place of magnesium ribbon.
3. Repeat the procedure in (1) and (2) using dilute sulphuric (VI) acid in place of dilute hydrochloric acid.

Sample observations

• Effervescence/bubbles produced/fizzing in all cases except when using copper.
• Colourless gas produced in all cases except when using copper.
• Gas produced extinguishes a burning wooden splint with an explosion/pop sound.

Explanation

Some metals react with dilute acids, while others do not. Metals which react with acids produce bubbles of hydrogen gas. Hydrogen gas is a colourless gas that extinguishes a burning splint with a pop sound. This shows acids contain hydrogen gas.

This hydrogen is displaced/removed from the acids by some metals like magnesium, zinc, aluminium, iron, and sodium.

Some other metals like copper, silver, gold, platinum, and mercury are not reactive enough to displace/remove the hydrogen from dilute acids.

Chemical equations

1. Magnesium + Hydrochloric acid → Magnesium chloride + Hydrogen

Mg(s) + 2HCl (aq) → MgCl₂ (aq) + H₂(g)

2. Zinc + Hydrochloric acid → Zinc chloride + Hydrogen

Zn(s) + 2HCl (aq) → ZnCl₂ (aq) + H₂(g)

3. Iron + Hydrochloric acid → Iron (II) chloride + Hydrogen

Fe(s) + 2HCl (aq) → FeCl₂ (aq) + H₂(g)

4. Aluminium + Hydrochloric acid → Aluminium chloride + Hydrogen

2Al(s) + 6HCl (aq) → 2AlCl₃ (aq) + 3H₂(g)

5. Magnesium + Sulphuric (VI) acid → Magnesium sulphate + Hydrogen

Mg(s) + H₂SO₄ (aq) → MgSO₄ (aq) + H₂(g)

6. Zinc + Sulphuric (VI) acid → Zinc sulphate + Hydrogen

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

7. Iron + Sulphuric (VI) acid → Iron (II) sulphate + Hydrogen

Fe(s) + H₂SO₄ (aq) → FeSO₄ (aq) + H₂(g)

8. Aluminium + Sulphuric (VI) acid → Aluminium sulphate + Hydrogen

2Al(s) + 3H₂SO₄ (aq) → Al₂(SO₄)₃ (aq) + 3H₂(g)

2. Reaction of metal carbonates and hydrogen carbonates with mineral acids

All acids react with carbonates and hydrogen carbonates to form salt, water, and produce carbon (IV) oxide gas.

Metal carbonate + Acid → Salt + Water + Carbon (IV) oxide gas

Metal hydrogen carbonate + Acid → Salt + Water + Carbon (IV) oxide gas

Experiment: Reaction of metal carbonates and hydrogen carbonates with mineral acids

1. Place 5 cm³ of dilute hydrochloric acid in a small test tube. Add half a spatula full of sodium carbonate. Stopper the test tube using a cork with delivery tube directed into lime water. Record the observations made. Test the gas also with a burning splint.
2. Repeat the procedure in (1) using zinc carbonate, calcium carbonate, copper carbonate, sodium hydrogen carbonate, potassium hydrogen carbonate in place of sodium carbonate.
3. Repeat the procedure in (1) and (2) using dilute sulphuric (VI) acid in place of dilute hydrochloric acid.

Set up of apparatus

Sample observations

• Effervescence/bubbles produced/fizzing in all cases.
• Colourless gas produced in all cases.
• Gas produced forms a white precipitate with lime water.

Explanation

All metal carbonate/hydrogen carbonate reacts with dilute acids to produce bubbles of carbon (IV) oxide gas. Carbon (IV) oxide gas is a colourless gas that extinguishes a burning splint. When carbon (IV) oxide gas is bubbled in lime water, a white precipitate is formed.

Chemical equations

1. Sodium carbonate + Hydrochloric acid → Sodium chloride + Carbon (IV) oxide + Water

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

2. Calcium carbonate + Hydrochloric acid → Calcium chloride + Carbon (IV) oxide + Water

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

3. Magnesium carbonate + Hydrochloric acid → Magnesium chloride + Carbon (IV) oxide + Water

MgCO₃(s) + 2HCl (aq) → MgCl₂ (aq) + H₂O (l) + CO₂ (g)

4. Copper carbonate + Hydrochloric acid → Copper (II) chloride + Carbon (IV) oxide + Water

CuCO₃(s) + 2HCl (aq) → CuCl₂ (aq) + H₂O (l) + CO₂ (g)

5. Copper carbonate + Sulphuric (VI) acid → Copper (II) sulphate + Carbon (IV) oxide + Water

CuCO₃(s) + H₂SO₄ (aq) → CuSO₄ (aq) + H₂O (l) + CO₂ (g)

6. Zinc carbonate + Sulphuric (VI) acid → Zinc sulphate + Carbon (IV) oxide + Water

ZnCO₃(s) + H₂SO₄ (aq) → ZnSO₄ (aq) + H₂O (l) + CO₂ (g)

7. Sodium hydrogen carbonate + Sulphuric (VI) acid → Sodium sulphate + Carbon (IV) oxide + Water

NaHCO₃(s) + H₂SO₄ (aq) → Na₂SO₄ (aq) + H₂O (l) + CO₂ (g)

8. Potassium hydrogen carbonate + Sulphuric (VI) acid → Potassium sulphate + Carbon (IV) oxide + Water

KHCO₃(s) + H₂SO₄ (aq) → K₂SO₄ (aq) + H₂O (l) + CO₂ (g)

9. Potassium hydrogen carbonate + Hydrochloric acid → Potassium chloride + Carbon (IV) oxide + Water

KHCO₃(s) + HCl (aq) → KCl (aq) + H₂O (l) + CO₂ (g)

10. Sodium hydrogen carbonate + Hydrochloric acid → Sodium chloride + Carbon (IV) oxide + Water

NaHCO₃(s) + HCl (aq) → NaCl (aq) + H₂O (l) + CO₂ (g)

3. Neutralization by bases/alkalis

All acids react with bases to form a salt and water only. The reaction of an acid with metal oxides/hydroxides (bases) to form salt and water only is called a neutralization reaction.

Since no effervescence/bubbling/fizzing takes place during neutralization:

• The reaction with alkalis requires a suitable indicator. The colour of the indicator changes when all the acid has reacted with the soluble solution of the alkali (metal oxides/hydroxides).
• Excess of the base is added to ensure all the acid reacts. The excess base is then filtered off.

Experiment 1: Reaction of alkali with mineral acids

1. Place about 5 cm³ of dilute hydrochloric acid in a boiling tube. Add one drop of phenolphthalein indicator. Using a dropper/teat pipette, add dilute sodium hydroxide dropwise until there is a colour change.
2. Repeat the procedure with dilute sulphuric (VI) acid instead of hydrochloric acid.
3. Repeat the procedure with potassium hydroxide instead of sodium hydroxide.

Sample observation:

Colour of phenolphthalein changes from colourless to pink in all cases.

Explanation

Bases/alkalis neutralize acids. Acids and bases/alkalis are colourless. A suitable indicator like phenolphthalein changes colour to pink when all the acid has been neutralized by the bases/alkalis. Phenolphthalein changes colour from pink to colourless when all the bases/alkalis have been neutralized by the acid.

Chemical equations

• Sodium oxide + Hydrochloric acid → Sodium chloride + Water
  Na₂O(s) + HCl → NaCl(aq) + H₂O(l)
• Potassium oxide + Hydrochloric acid → Potassium chloride + Water
  K₂O(s) + HCl → KCl(aq) + H₂O(l)
• Sodium hydroxide + Hydrochloric acid → Sodium chloride + Water
  NaOH(s) + HCl → NaCl(aq) + H₂O(l)
• Ammonia solution + Hydrochloric acid → Ammonium chloride + Water
  NH₄OH(s) + HCl → NH₄Cl (aq) + H₂O (l)
• Potassium hydroxide + Hydrochloric acid → Potassium chloride + Water
  KOH(s) + HCl → KCl(aq) + H₂O(l)
• Sodium hydroxide + Sulphuric (VI) acid → Sodium sulphate + Water
  2NaOH(s) + H₂SO₄ → Na₂SO₄ (aq) + 2H₂O (l)
• Potassium hydroxide + Sulphuric (VI) acid → Potassium sulphate + Water
  2KOH(s) + H₂SO₄ → K₂SO₄ (aq) + 2H₂O (l)
• Ammonia solution + Sulphuric (VI) acid → Ammonium sulphate + Water
  2NH₄OH(s) + H₂SO₄ → (NH₄)₂SO₄ (aq) + 2H₂O (l)
• Magnesium hydroxide + Sulphuric (VI) acid → Magnesium sulphate + Water
  Mg(OH)₂(s) + H₂SO₄ → MgSO₄ (aq) + 2H₂O(l)
• Magnesium hydroxide + Hydrochloric acid → Magnesium chloride + Water
  Mg(OH)₂(s) + HCl(aq) → MgCl₂ (aq) + 2H₂O(l)