CHEMISTRY FORM TWO NOTES

Chapter One: Atomic Structure

Chemistry for Secondary Schools - Form Two

Introduction

Substances are made up of tiny particles called atoms. Atoms are the smallest particles of matter that carry an element's properties. In this chapter, you will learn about the concept of atomic structure and the determination of the atomic number and mass number of an element. The competencies developed will enable you to analyse the composition, behaviour, and properties of different chemical substances, enhancing your understanding of their interactions and applications in real-life situations.

Think

Atomic structure is the foundation of modern life

Concept of Atomic Structure

Task 1.1

Use an interactive simulation, video, or any other reliable resource to visualise the structure of an atom and examine its sub-atomic particles.

Atoms are fundamental units that make up all matter, forming everything in the surroundings, including the air, materials, and objects used daily. Understanding atomic structure helps scientists predict the behaviour of substances and design new materials and cleaner energy. For example, understanding the behaviour of atoms helps in food preservation, water purification, making batteries and medicines. This leads to technological and scientific developments that shape the world.

An atom is composed of smaller particles called sub-atomic particles. These include protons, neutrons, and electrons. The arrangement of these particles within an atom is referred to as the atomic structure. The structures of atoms are understood through the atomic theory, which has been developed over time through experiments and scientific discoveries.

Task 1.2

Use reliable resources to analyse various atomic theory discoveries and come up with an idea about the structure of the atom.

About the year 400 BC, a Greek philosopher known as Democritus was the first to consider the idea that matter is made up of particles. Such an idea was not accepted because there was no experimental evidence to support it. About 2000 years later, an English man called John Dalton revived the discussion. He used experimental evidence to convince people that matter is made up of particles called atoms. Through that experiment, he deduced Dalton's spherical model of an atom shown in Figure 1.1. Dalton's discovery helped scientists understand chemical reactions and how substances combine.

Figure 1.1: Dalton's spherical model of an atom

Dalton Atomic Theory

In 1803, Dalton developed the theory about the atom. The four main points (assumptions) of Dalton's Atomic Theory are summarised as follows:

1. Matter is made up of tiny particles called atoms. (The word atom means 'indivisible' in Greek).
2. Atoms can neither be created nor destroyed.
3. Atoms of the same element are identical and have the same mass and properties. Atoms of a given element are different from those of any other element. Atoms of different elements can be distinguished from one another by their respective relative weights.
4. Compounds are formed by a combination of two or more different kinds of atoms. Atoms always combine in simple whole number ratios.

Dalton never imagined that anyone would ever be able to see an atom. However, modern technology has provided direct evidence that shows the positions and patterns of individual atoms. The use of modern technology has enabled scientists to carry out experiments on the atom that Dalton could not. This has led to slight modifications to Dalton's Atomic Theory and thus formulated the so called modern concepts of Dalton's Atomic Theory.

Modern Modifications to Dalton's Atomic Theory

1. Atoms can be either created or destroyed by means of nuclear reactions. The atom can change form through special processes such as nuclear fusion (combining the atomic nuclei) or nuclear fission (splitting the atomic nucleus). For example, an atom of uranium-235 can be split into two separate atoms.
2. Some elements have atoms of more than one kind which differ slightly in mass. Such atoms are called isotopes. For example, carbon has three isotopes known as carbon-12, carbon-13, and carbon-14.
3. An atom is made up of smaller sub-atomic particles called protons, neutrons, and electrons.
4. Atoms of different elements may chemically combine in many different ratios to form compounds.

The modern atomic theory builds on Dalton's original ideas by recognising sub-atomic particles, isotopes, and nuclear reactions while retaining the ideas of chemical combinations and reactions. Dalton's discovery thus helped scientists understand chemical reactions and how substances combine.

Sub-atomic Particles

In 1897, J.J. Thomson carried out experiments and described an atom as a sphere of positive charge, with negative particles called electrons spread throughout the sphere. This model of the atom was referred to as plum pudding model, as shown in Figure 1.2.

Figure 1.2: Thomson's plum pudding model of the atom

Thomson, therefore, managed to discover the electron among the three sub-atomic particles. His discovery led to inventions of electronic devices such as televisions, radios and computers. However, another scientist called Ernest Rutherford reasoned that if Thomson's model was correct, then the mass of the atom was evenly spread throughout the atom. He carried out experiments and discovered that most of the mass of an atom is concentrated in the nucleus (central core) of the atom. Within the nucleus, there are positively charged particles called protons. This was the second sub-atomic particle to be discovered. Rutherford's discovery helped scientists discover nuclear energy used to produce electricity and in radiotherapy.

Rutherford's Findings

1. Protons, the positively charged particles of an atom, are located in the nucleus.
2. Most of the mass of the atom is located in the nucleus.
3. The nucleus has a relatively smaller volume compared to the whole atom.
4. Electrons have very small masses compared to the protons.
5. Most of the space in an atom is empty.
6. Electrons are negatively charged particles in an atom. They move around the nucleus in orbits.

Rutherford thus developed the planetary model of the atom shown in Figure 1.3.

Figure 1.3: Rutherford's planetary model of the atom

In 1932, a scientist named James Chadwick discovered the neutrons, which also forms part of the nucleus. Figure 1.4 shows the locations of neutrons and other sub-atomic particles in an atom. Neutrons have the same mass as protons but no charge. They are located in the nucleus of an atom. They were the third sub-atomic particles to be discovered. The Chadwick discoveries have made nuclear power possible, helping to produce electricity and develop medical treatments such as cancer therapy.

Figure 1.4: Location of sub-atomic particles in the atom

Properties of Neutrons

1. They have no charge (are neutral).
2. They have nearly the same mass as the corresponding protons.
3. They have a mass nearly 1840 times the mass of an electron.

Properties of Sub-atomic Particles

Sub-atomic particle	Symbol	Location	Charge	Real mass (g)	Relative mass
Proton	p+	In the nucleus	+1	1.6726 × 10-24	1
Neutron	n	In the nucleus	0	1.6750 × 10-24	1
Electron	e-	Outside the nucleus	-1	9.109 × 10-28	1/1840

Activity: To build a 3D atomic model of carbon

Requirements: Six medium-sized red beads and white beads, small-sized black beads, cardboard, clay or glue and string or wire

Procedure:

1. Build the nucleus by randomly arranging six red beads (protons) and six white beads (neutrons). Then, stick them together using glue to form a tight cluster at the centre.
2. Create electron shells using cardboard. Make two circular shells of different sizes.
3. Attach six black beads (electrons) to the shells - two on the inner shell and four on the outer shell.
4. Use strings or wires to connect the shells to the nucleus, creating a 3D model.

Determination of Atomic Number and Mass Number

Atomic Number

The atomic number (Z) of an element is the number of protons in the nucleus of an atom of that element. In a neutral atom, the number of protons equals the number of electrons. Therefore, the atomic number also represents the number of electrons in a neutral atom.

Mass Number

The mass number (A) of an atom is the total number of protons and neutrons in the nucleus of that atom. It can be calculated using the formula:

Mass Number Formula

Mass number (A) = Number of protons + Number of neutrons

It is also possible to calculate the number of neutrons and number of electrons of an atom if its mass number and atomic number are given.

Example 1.1

Atom Q has a mass number of 49 and an atomic number of 24. Calculate the number of neutrons and the number of electrons in atom Q.

Solution

Mass number = 49; atomic number = 24

(a) Neutron number = mass number – atomic number = 49 – 24 = 25

(b) Number of electrons = number of protons = atomic number = 24

Note

For the mass number with fractions, such as chlorine (35.5), calculating the number of neutrons and electrons involves only a whole number. In this case, for chlorine, 35 is used.

Nuclide Notation

Atoms of different elements can be represented by chemical symbols that indicate their respective atomic numbers and mass numbers. Using an arbitrary element X, the mass number (A) is placed on its upper left end, while its atomic number (Z) is placed on the lower left end. Thus, element X is represented as ^AX_Z. This is known as the nuclide notation. The following are examples of nuclide representations of different atoms:

• Hydrogen ¹H₁
• Boron ¹¹B₅
• Nitrogen ¹⁴N₇
• Oxygen ¹⁶O₈

With this information, it is possible to deduce the number of neutrons and electrons in the atom, and to write the electronic configuration. For example, in the oxygen atom, the mass number is 16 and the atomic number is 8. Therefore, the number of electrons is 8 and the number of neutrons is 16 – 8 = 8. The nucleus of the oxygen atom can therefore be represented as shown in Figure 1.7.

Figure 1.7: The nucleus of the oxygen atom

Isotopes

Task 1.6

Watch educational video on isotopes and explain the uses of isotopes in carbon dating, medicine and agriculture.

Atoms of the same element have the same number of protons. However, the number of neutrons in the atoms of the same element may vary. This means that the atomic number of an element does not vary but the mass number can vary. Such atoms of an element are called isotopes. Isotopes are atoms of the same element with the same number of protons but different number of neutrons. Such an existence of the element is called isotopy. Isotopy is the existence of atoms of the same element having the same atomic number but different mass numbers. It is also possible to get the number of sub-atomic particles in a given isotope.

Example 1.3

State the number of protons, neutrons, and electrons in the following isotopes:

(a) ¹²C and ¹⁴C

(b) ¹H, ²H and ³H

Solution

(a) ¹²C, Mass number = 12

Number of protons = atomic number = 6

Number of electrons = number of protons = 6

Number of neutrons = 12 – 6 = 6

¹⁴C, Mass number = 14

Number of protons = 6

Number of electrons = 6

Number of neutrons = 14 – 6 = 8

(b) ¹H, Mass number = 1

Number of protons = 1

Number of electron = 1

Number of neutron = 1 – 1 = 0

Examples of Isotopes and Their Abundances

Element	Chemical symbol	Atomic number	Isotopes	Abundance
Hydrogen	H	1	1H (protium or hydrogen)	99.99%
			2H (deuterium)	0.01%
			3H (tritium)	Trace
Carbon	C	6	12C	98.9%
			14C	1.1%
Chlorine	Cl	17	35Cl	75%
			37Cl	25%

Applications of Isotopes

Isotopes are special forms of elements that can be used in various areas, including research, medicine, industry, and agriculture. Some isotopes are radioactive. Radioactive isotopes are special types of atoms that give off energy called radiation. Even though the emitted radiations are dangerous if not handled properly, these isotopes are very useful in many ways. Scientists use these isotopes in various applications, including determining the age of ancient objects using carbon-14, treating diseases in hospitals, tracking how plants absorb nutrients in the soil, and producing electricity for everyday uses.

Carbon Dating: Finding the Age of Ancient Items

Scientists use carbon-14 to determine the age of ancient objects. Carbon-14 is a radioactive isotope of carbon that slowly breaks down over time. When a plant or animal dies, it stops taking in carbon-14 from the air. Scientists measure how much carbon-14 is left in bones, wood, or fossils to estimate the number of years that have passed since the organism died. Carbon dating helps archaeologists and historians learn about the past, including the age of ancient human tools, animal fossils, and historical items.

Tracers in Medicine, Industry, and Agriculture

Some radioactive isotopes are used as tracers that help scientists track movements or processes inside the body, in the environment, or in industrial systems.

(a) Medicine: Diagnosing and treating diseases

Iodine-131 is used in hospitals to check how the thyroid gland works. When a small dose of iodine-131 is administered to a patient, the radiation emitted by the iodine helps to create images of the thyroid. Special machines detect the radiation to help diagnose thyroid problems. The radiation also helps to shrink or destroy damaged thyroids cells leading to the treatment of thyroid related diseases.

(b) Agriculture: Studying how plants absorb nutrients

Phosphorus-32 is used to track how plants take in nutrients from the soil. Scientists use it to improve fertilisers and help farmers grow healthier crops.

(c) Industrial and environmental applications

Some isotopes, such as tritium (³H) tracks how water moves underground. This allows scientists to understand the sources of drinking water. Chlorine isotopes are used to study the movements of chlorine in water sources such as rivers, lakes, and underground water. Sodium-24 is used to detect leaks in underground pipes. This isotope is made artificially.

Relative Atomic Mass

Task 1.7

Consider a library search to find the differences between atomic mass and relative atomic mass.

An atom is very small and it would be difficult to measure its actual mass. To overcome this difficulty, chemists developed a simpler way to express the mass of an atom. This involved expressing the mass of an atom in relation to a chosen standard atomic mass. The carbon atom was chosen as the standard atom (reference atom) and its mass was arbitrarily chosen as 12 units (not actual value). Then, using an instrument called a mass spectrometer, all the other atoms were compared to this standard atom. This reference is called the Carbon-12 scale. For example, it was found that the:

• magnesium atom was twice as heavy as the reference atom, so its mass was put at 24.
• hydrogen atom was 1/12 as heavy as the reference atom, so its mass was put at 1.
• helium atom was 1/3 as heavy as the reference atom, so its mass was put at 4.

The mass of an atom obtained by comparing it with the arbitrary mass of a carbon-12 atom is called its relative atomic mass (R.A.M. or A_r). The relative atomic mass of an element is the average mass of one atom of the element relative to 1/12th the mass of one atom of carbon-12. Therefore, R.A.M. may not necessarily be a whole number.

Task 1.8

Average mass of atom of an element (∑ = Summation)

For isotopic elements, the relative atomic mass (R.A.M.) can be calculated using the following formula:

Relative Atomic Mass Formula

Relative atomic mass (R.A.M.) = ∑ (isotopic mass × percentage abundance)

Note: ∑ = Summation

Chemistry for Secondary Schools - Form Two

© MITIHANI POPOTE 2025

Chapter Two: Periodic Classification

Chemistry for Secondary Schools - Form Two

Introduction

The periodic table is a systematic arrangement of elements that helps in understanding patterns in their physical and chemical properties. Elements in the same row (period) and the same column (group) exhibit trends in melting points, boiling points, density, electronegativity, ionisation energy, atomic size, and reactivity. Recognising these trends is essential for applications in various settings, including home settings and industries such as material manufacturing, energy storage, and healthcare.

Think

How does the periodic table help us predict the behavior of elements?

Development of the Periodic Table

The periodic table has evolved over time as scientists discovered more elements and recognized patterns in their properties. The modern periodic table is based on the periodic law, which states that the properties of elements are periodic functions of their atomic numbers.

Mendeleev's Periodic Table

Dmitri Mendeleev, a Russian chemist, is credited with creating the first widely recognized periodic table in 1869. He arranged elements in order of increasing atomic mass and noticed that elements with similar properties occurred at regular intervals.

Modern Periodic Table

The modern periodic table is arranged in order of increasing atomic number rather than atomic mass. This arrangement better reflects the periodic nature of element properties.

Modern Periodic Table

Electronic Configuration and Element Positioning

The position of an element in the periodic table is determined by its electronic configuration. Elements with the same number of electrons in their outermost shells belong to the same group, while elements with the same number of electron shells belong to the same period.

Trends in Physical and Chemical Properties

Task 2.5

Use an interactive simulation or any reliable resources to explore the trends in physical properties across periods of the periodic table. Analyze these trends and explain their practical applications in real-life scenarios.

Trends in Physical Properties Across Periods

(a) Atomic Size Decreases Across a Period

Atomic radius decreases from left to right across a period due to increasing nuclear charge, which pulls electrons closer to the nucleus.

(b) Increase in Melting and Boiling Points

Melting and boiling points generally increase across a period for metals due to stronger metallic bonding, reaching a peak around Groups IV-VI, then decrease for non-metals.

Example: Elements found in Groups V to VIII, such as nitrogen (N₂), oxygen (O₂), and neon (Ne), exist as simple molecules with low melting and boiling points. These are commonly used in gas form; for example, oxygen is used in hospitals for respiration, nitrogen in food preservation, and neon in lighting systems.

(c) Increase in Ionisation Energy

Ionisation energy increases from left to right across a period due to stronger nuclear attraction, making it harder to remove electrons from an atom. Elements with high ionisation energy, such as noble gases, are used in lighting and insulation due to their chemical stability.

Trends in Chemical Properties Across Periods

(a) Increase in Electronegativity

Electronegativity increases from left to right across a period, with non-metals attracting electrons more strongly. This is due to the stronger nuclear attraction resulting from the increasing number of protons in the nucleus. This trend is particularly crucial in various applications, such as semiconductor technology, where elements like silicon and germanium are widely used in electronic devices.

(b) Decrease in Metallic Character

Metallic characters decrease, and non-metallic characters increase across a period. Metals tend to lose electrons, while non-metals gain electrons. This trend is important in various applications such as battery production, where metals act as electron donors.

(c) Variations in Chemical Reactivity

Metals on the left are highly reactive and lose electrons easily, while non-metals on the right become more reactive in gaining electrons. For metals, reactivity decreases from left to right, while for non-metals, reactivity increases from left to right. This trend plays crucial roles in different activities such as drug formulation and material design, ensuring the stability and effectiveness of compounds in pharmaceuticals and engineering.

Trends in Physical Properties Down a Group

(a) Atomic Size Increases Down a Group

Atomic size increases down a group as more electron shells are added. This expansion causes atoms to become larger, affecting their physical behavior. Larger atomic size influences material performance under high-pressure environments, making certain elements suitable for deep-sea applications and industrial machinery.

(b) Density Increases Down a Group

Density also increases down a group as atomic mass increases more significantly than the atomic volume. Heavier elements tend to have stronger structural properties, making them valuable in industries that require durability and strength. These elements play crucial roles in construction and aerospace engineering, where materials must withstand extreme conditions.

(c) Melting Points Generally Decrease Down a Group

Melting points for metals generally decrease down a group due to weaker metallic bonding. As atomic sizes increase, the attraction between metal atoms weakens, reducing the energy required to melt the substances. This trend is significant in the design of alloys for safety devices such as fuses, which need to melt easily to prevent electrical hazards.

(d) Ionisation Energy Decreases Down a Group

Ionisation energy decreases down a group in the periodic table. This is because as atomic sizes increase, the outermost electrons are farther from the nucleus. As a result, the attractions between the nucleus and outer electrons become weaker, making it easier for metals to lose electrons. The trend in ionisation energy influences an element's reactivity, the types of compounds it forms, its electrical conductivity, and its potential biological and industrial processes such as metallurgy, battery design, and semiconductor manufacturing.

Trends in Chemical Properties Down a Group

(a) Electronegativity Decreases Down a Group

Electronegativity decreases down a group. This is because, as atomic size increases, the outer electrons are farther from the nucleus. This condition reduces the attractions between the nucleus and electrons, making atoms less able to attract electrons in a chemical bond.

(b) Metallic Character Increases Down a Group

Metallic character increases down a group. These properties are crucial in selecting metals for use in catalysis and chemical processing industries. Reactivity trends differ between metals and non-metals as you move down a group. Metals become more reactive because they are more likely to lose electrons, which is advantageous in processes such as metal extraction from ores. In contrast, non-metals become less reactive because their ability to attract electrons weakens. This behavior is applied in industries where reactive metals facilitate the development of substances as in cleaning agents, while non-metals help stabilize compounds to prevent undesired reactions.

Trends in Selected Groups

Groups I, II, and VII elements exhibit distinct physical and chemical properties due to their unique positions in the periodic table. Group I (alkali metals) and Group II (alkaline earth metals) are highly reactive metals, while Group VII (halogens) consists of reactive non-metals. Their physical properties, such as atomic radius, ionisation energy, density, and melting points, exhibit clear trends within each group.

Group I: Alkali Metals

Group I consists of metals such as lithium (Li), sodium (Na), potassium (K), rubidium (Rb), and caesium (Cs). Each of these elements has one electron in its outermost shell. Lithium, sodium and potassium react very readily with water or air and are stored in oil.

Name	Atomic Number	Electronic Configuration	Atomic Radius (pm)	Melting Point (°C)	Density (g/cm³)	Electronegativity
Lithium	3	2,1	152	180	0.53	1.0
Sodium	11	2,8,1	186	98	0.97	0.9
Potassium	19	2,8,8,1	231	63	0.86	0.8

Note

Francium (Fr) is also an alkali metal but is rarely discussed in experiments involving Group I elements due to its radioactive nature. It is also among the rarest naturally occurring elements.

Properties of Group I Elements

Physical Properties:

• They are good conductors of heat and electricity.
• They are soft metals.
• They have low density.
• They have shiny surfaces when freshly cut.

Chemical Properties:

• They burn in oxygen or air with a characteristic flame colour to form white solid oxides. These oxides dissolve in water to form alkaline solutions of the metal hydroxides.
• They react vigorously with water to give alkaline solutions and hydrogen gas.

Activity 2.1: To investigate the reactivity of Group I elements (alkali metals) down the group

Requirements: Small pieces of lithium (Li), sodium (Na), and potassium (K), three beakers, water, phenolphthalein or universal indicator, forceps or tongs, watch glass, dropper, safety goggles and gloves

Caution: Handle Group I elements cautiously by using small quantities, wearing protective gear, working in a controlled environment, avoiding direct contact, and keeping a fire extinguisher nearby to prevent accidents during their vigorous reactions with water.

Procedure:

1. Fill three beakers halfway with water.
2. Add a few drops of phenolphthalein or universal indicator to each beaker.
3. Using forceps, carefully drop a small piece of lithium into the first beaker and observe the reaction.
4. Repeat Step 3 for sodium and potassium in separate beakers.
5. Record observations for each metal, focusing on the rate of reaction, colour change (due to the indicator), production of gas and sound.
6. Clean up all the apparatus safely under the supervision of the teacher.

Questions:

1. Which element was the fastest to react with water?
2. How did the colour of water change and what did it indicate?
3. What trend did you observe in the reactivity of those elements down the group?
4. How does this experiment support the idea of periodic classification?

Group II: Alkaline Earth Metals

Group II consists of elements such as beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). These elements have two electrons in their outermost shells.

Name	Atomic Number	Electronic Configuration	Atomic Radius (pm)	Melting Point (°C)	Density (g/cm³)	Electronegativity
Beryllium	4	2,2	112	1287	1.85	1.5
Magnesium	12	2,8,2	160	650	1.74	1.2
Calcium	20	2,8,8,2	197	842	1.55	1.0

Properties of Group II Elements

Physical Properties:

• They are harder metals than those in Group I.
• They are silvery grey in colour when pure and clean. However, they tarnish quickly when left in the air due to the formation of the respective metal oxides.
• They are good conductors of heat and electricity.

Chemical Properties:

• They burn in oxygen or air with a characteristic flame colour to form a solid white oxide.
• They react with water but much less vigorously than the elements in Group I.
• The reactivity of metals increases down the group. For example, the reaction of calcium with water is vigorous, while that of magnesium with water is relatively slow.

Activity 2.2: To demonstrate the differences in reactivity of calcium and magnesium with water

Requirements: Two test tubes, magnesium ribbon, calcium, distilled water, and measuring cylinder

Procedure:

1. Transfer about 5 cm³ of distilled water in a test tube.
2. Add a small amount of calcium (half spatulaful) from the container in which it is stored. Record your observations.
3. Repeat Steps 1 and 2 using a clean piece of magnesium ribbon. Record your observations.

Questions:

1. What happens when the powder of calcium and magnesium ribbon are dropped in water?
2. What is the difference in the reactivities of calcium and magnesium?

Group VII: Halogens

Task 2.7

Use online sources or any reliable resources to study trends in physical and chemical properties in Group VII.

Group VII elements include fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). These are highly reactive non-metals.

Properties of Group VII Elements

Physical Properties:

• They exist in different physical states: fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid.
• Their densities increase down the group.
• Their melting and boiling points increase down the group.

Chemical Properties:

• Their reactivity decreases down the group.
• They react with metals to form salts.
• They react with hydrogen to form acids (hydrogen halides).

Chapter Summary

1. The periodic table of elements is a method of displaying chemical elements in a table format. It was developed after several modifications to the Mendeleev's periodic table.
2. Mendeleev's Periodic Law states that the properties of elements are periodic functions of their relative atomic masses.
3. The Modern Periodic Law states that the properties of elements change systematically according to their atomic numbers.
4. Periodicity refers to the regular periodic changes of properties of elements due to changes in atomic numbers.
5. Elements with the same number of electrons in their outermost shells belong to the same group.
6. The group number signifies the number of electrons in the outermost shell of an element.
7. Elements with the same number of shells belong to the same period.
8. The electronic configuration of an element entails the electronic arrangements in the shell(s), number of shells, and the group to which the element belongs.

Review Exercise 2

Choose the correct answer for Questions 1-7. For other questions, provide the answers as per the demands indicated.

1. Non-metals are generally better ______ than metals.

(a) conductors of electricity

(b) brittle materials

(c) malleable materials

(d) conductors of heat

2. The electronic arrangement of an element is 2; 3. This element is in ______ of the periodic table.

(a) Group II

(b) Group VIII

(c) Period 3

(d) Period 2

3. The following statements describe the alkaline earth metals except they ______ burn in oxygen to form solid white oxides.

(a) become less reactive down the group.

(b) are good conductors of heat and electricity.

(c) react with water.

Answer the following questions:

11. (a) Given the elements calcium, sulfur, chlorine, helium and neon, write their:

(i) period numbers.

(ii) group numbers.

(iii) atomic numbers.

(iv) number of electrons in their atoms.

(v) electronic configurations.

(b) Which of the above elements would you expect to have similar properties? Give reasons.

12. During a science lesson, students added a small piece of sodium (Na) into water and observed fizzing and heat. They repeated the same with magnesium (Mg), and there was very little reaction. Using periodic classification, explain why the reaction of sodium with water was vigorous than that of magnesium.

13. Greenish Secondary School is building a simple application (App) that can help Form Two students to identify if an element is a metal, non-metal, or metalloid. What features of periodic classification should be included in the App to help students make the correct decision?

Chemistry for Secondary Schools - Form Two

© MITIHANI POPOTE 2025

Chapter Three: Chemical Bonding, Formula and Nomenclature

Chemistry for Secondary Schools - Form Two

Introduction

Chemical bonding involves holding atoms together to form molecules or compounds. A chemical formula refers to symbols and numbers that represent the composition of a particular chemical substance. Nomenclature means naming. In this chapter, you will learn about the concept of chemical bonding, the concept of chemical formulas, the determination of empirical and molecular formulas of common compounds and the nomenclature of binary inorganic compounds using the IUPAC system. The competencies developed will enable you to explore the relationships between chemical substances and apply the acquired knowledge in understanding properties of different materials used in daily life.

Think

Contribution of chemical bonding, formulas, and nomenclature to global processes and activities

Concept of Chemical Bonding

Task 3.1

Use videos, interactive simulations, or other reliable resources to explore the concept of chemical bonding.

A bond is anything that holds two or more substances together. Many things used in our daily lives are constructed using different materials joined together by some bonds. For example, in a brick wall, each brick is joined to the other by a bond made of mortar. Similarly, chemical substances are made of atoms that are held together by chemical bonds.

Figure 3.1: Bricks bonded using mortar to form wall

A chemical bond is a force of attraction that holds atoms or ions together to form molecules or compounds. The bond may result from forces of attraction between oppositely charged ions or through the sharing of electrons.

Note

A molecule is the smallest particle of an element or compound which can normally exist separately.

Chemical Bonding

Chemical bonding involves electrons in the outermost shells of atoms. When the outermost shell is completely filled with electrons, the atoms are said to be stable, otherwise they are unstable. Only helium, neon, and argon have stable electronic arrangements. Helium has a maximum number of 2 electrons in its outermost shell. Neon and argon have a maximum of 8 electrons in each of their respective outermost shells. These stable atoms are generally unreactive and can exist freely as single atoms. Unstable atoms cannot exist freely as single atoms. For unstable atoms to become stable, they should acquire electronic arrangements similar to those of noble gases. This means that they can either lose, gain or share electrons through chemical bonding.

Types of Chemical Bonds

Ionic Bonding

Formed by transfer of electrons from a metal to a non-metal

NaCl, MgO

Ionic Bond Diagram

Covalent Bonding

Formed by sharing of electrons between non-metals

H₂O, CO₂

Covalent Bond Diagram

Metallic Bonding

Formed by attraction between positive metal ions and delocalized electrons

Fe, Cu, Al

Metallic Bond Diagram

Ionic (Electrovalent) Bonding

Ionic bonding usually occurs between a metal and a non-metal. It involves the transfer of electron(s) from the atoms of the metal to the atoms of the non-metal. This results in the formation of positively charged ions (cations) and negatively charged ions (anions) that are held together by electrostatic forces.

Example: Formation of Sodium Chloride (NaCl)

Sodium atom (Na) loses one electron to form Na⁺ ion

Chlorine atom (Cl) gains one electron to form Cl^- ion

The resulting ions attract each other to form NaCl

Covalent Bonding

Covalent bonding takes place between two or more non-metals. It involves atoms of the non-metals sharing electrons that are in their outermost shells. The shared electrons are counted as part of the outer shell of each atom involved in the bond.

Example: Formation of Water (H₂O)

Oxygen atom shares electrons with two hydrogen atoms

Each bond consists of a shared pair of electrons

This gives each hydrogen atom a stable configuration like helium, and oxygen a stable configuration like neon

Valency

Task 3.3

Use ball and stick models or colored beads to represent electrons in different shells. Build atoms and determine the valency by observing how many electrons are in the outer shell.

Valency refers to the ability of an atom of a given element to combine with other atoms, and is measured by the number of electrons that the atom will donate, receive or share to form a chemical bond. It is the combining power/capacity of an element or a radical. The combining capacity of an atom of a given element is determined by the number of hydrogen atoms it combines with or displaces.

Examples

The valency of chlorine is 1 because one atom of hydrogen combines with one atom of chlorine to form hydrogen chloride (HCl).

The valency of zinc is 2 because two atoms of hydrogen are displaced from dilute acids by one atom of zinc.

It is easy to predict the valencies of elements from the periodic table:

• Group I elements have one electron in their outermost shells, and so, their valency is 1.
• Group II elements have two electrons in their outermost shells, hence their valency is 2.
• For elements with more than four electrons in the outermost shells, the valency number is usually obtained by subtracting the number of electrons from eight.

Example

Sulfur with six electrons in the outermost shell has a valency of 8 - 6 = 2.

Those outermost electrons are called valence electrons. Some elements have more than one valency. For example, iron has valencies of 2 and 3, copper has valencies of 1 and 2, lead has valencies of 2 and 4, and manganese has valencies of 2, 4 and 7.

Valencies of Ionic Elements and Radicals

Category	Valency 1	Valency 2	Valency 3
Metals	Potassium (K+) Sodium (Na+) Silver (Ag+)	Calcium (Ca2+) Magnesium (Mg2+) Zinc (Zn2+)	Aluminium (Al3+) Iron (Fe3+)
Non-metals	Chlorine (Cl-) Fluorine (F-)	Oxygen (O2-) Sulfur (S2-)	Nitrogen (N3-)
Radicals	Ammonium (NH4+) Hydroxide (OH-) Nitrate (NO3-)	Carbonate (CO32-) Sulfate (SO42-)	Phosphate (PO43-)

Radicals

Task 3.4

Use a chemistry simulation or software to identify, classify and write different formulas of radicals.

A radical is a group of atoms which behaves as a single unit and has a positive or negative charge. It contains at least one unpaired electron. Such a group maintains its identity throughout any chemical reaction. Most radicals form the non-metallic part of a compound, so their ions are negatively charged. Examples are CO₃^2- and SO₄^2- ions. An exception is for the ammonium radical, NH₄⁺, which behaves like the metallic part of a compound and forms a positive ion. The valency of the radical is the same as the numerical value that the group acquires when it loses or gains an electron to form an ion.

Note

Ammonium radical (NH₄⁺) has a valency of 1 and can react like metals. Its compounds are similar to those of Group I elements.

Oxidation State

Oxidation state (also called oxidation number) is the total number of electrons that an atom either gains or loses in order to form a chemical bond with another atom. It is the measure of the electron control that an atom has in a compound compared to the atom in the pure element. The neutral atom has no charge.

Rules for Assigning Oxidation States

1. The oxidation number of free elements is zero. For example, all elements in the periodic table have the oxidation number of zero.
2. The sum of the oxidation states of all atoms forming a molecule or ion is the net charge of that species.
3. In simple ions that consist of only one atom, the oxidation number is equal to the charge on the ion.
4. In their compounds, Group I metals have an oxidation number of +1. Group II metals have an oxidation number of +2, while Group III metals have an oxidation number of +3.
5. In their compounds, halogens always have an oxidation number of -1.
6. Hydrogen has an oxidation state of +1 in most compounds. The exception is in hydrides of active metals where the oxidation number is -1.
7. Oxygen has an oxidation state of -2 when present in most compounds, except in peroxides where it is -1, and when bonded with fluorine where it is +2.

Example 3.1

Find the oxidation state of chlorine in KClO₃.

Solution

The oxidation number of potassium is +1

The oxidation number for oxygen is -2

For the three oxygen atoms, the oxidation number is (-2 × 3) = -6

KClO₃ is a neutral compound. Therefore, the oxidation number of the compound is zero.

Therefore, +1 + Cl + (-6) = 0

Cl - 5 = 0

Cl = +5

The oxidation number of chlorine in KClO₃ is +5.

Concept of Chemical Formulas

Task 3.5

Obtain containers such as bottles containing chemicals in the laboratory. Examine them and identify the chemical formulas of the substances on their labels.

A chemical formula is a representation that uses symbols to show the proportions of the elements present in a chemical compound. The number of atoms or groups of atoms are shown by number subscripts. For example, the chemical formula for sodium sulfide is Na₂S, which shows that two atoms of sodium combine with one atom of sulfur to form the molecule of sodium sulfide. For groups of atoms such as radicals, a bracket is used to show that they are being considered as a unit under one valency. For example, in calcium nitrate, Ca(NO₃)₂, the NO₃ radical is in brackets.

Writing Chemical Formulas

There are some points to remember when writing chemical formulas:

1. Positively charged ions (cations) are written before negatively charged ions (anions).
2. A radical must be treated as a unit.
3. Brackets are not used for single elements.
4. The valency 1 is simply assumed and not written in the formula.

The symbols and valencies of the atoms and radicals are important in writing a chemical formula. For example, for arbitrary elements W and X with valencies m and n, respectively, and where X can be a radical or an atom, the following steps can be used to come up with a chemical formula of their compounds:

Step 1: Write the symbols of the elements and radicals; in this case W and X.

Step 2: Write down the ions used, with their valencies as superscripts, that is W^m Xⁿ.

Step 3: Interchange the valencies of W and X and write them as subscripts.

Step 4: The formula of the chemical compound is W_nX_m.

Note

When m and n are equal, there is no need for the exchange and, therefore, are not written since they are in a ratio of 1:1.

Example 3.5

Give the formula of the compound of aluminium and sulfate.

Solution

Step 1: Al SO₄

Step 2: Al³⁺ SO₄^2-

Step 3: Interchange valencies: Al₂ (SO₄)₃

Step 4: The chemical formula is Al₂(SO₄)₃.

Types of Chemical Formulas

Chemical formulas can be divided into three types, namely empirical formula, molecular formula, and structural formula.

Empirical Formula

An empirical formula is the formula which represents the simplest ratio of the atoms or ions in a compound. The simplest formula is usually determined by considering experimental data. That is why it is called "empirical" which means "based on experimentation". For example, CH₂ shows there are twice as many hydrogen atoms as carbon atoms. It does not show the exact number of each atom of the element in the compound.

Molecular Formula

A molecular formula shows the actual number of each atom in a molecule. It is a multiple of the empirical formula. For example, if the empirical formula is CH₂, its molecular formula may be C₂H₄, C₃H₆, C₄H₈, and so on. Therefore, a molecular formula is equal to n multiplied by the empirical formula, where n is a whole number. Note that when n is 1, the empirical formula equals the molecular formula.

Structural Formula

A structural formula is a graphic representation of molecular structure showing how the atoms are arranged. At this level, only the empirical and molecular formulas will be studied.

Formula Calculations

When the percentage compositions of the elements that make up a compound are known, it is possible to obtain both the empirical and molecular formulas of such a compound. The following are the steps considered when calculating the empirical formula:

1. Obtain the mass of each element in the sample compound. If expressed in percentages, convert the percentage of each of the elements to mass.
2. Divide the mass of each element by its relative atomic mass (R.A.M.).
3. Divide each of the values obtained in Step 2 by the lowest value among them.
4. Convert the ratios in Step 3 to whole numbers. These whole numbers give the ratio of each element in the compound.

Example 3.6

What is the empirical formula for a compound of mass 8.1 g if it consists of 4.9 g of magnesium and 3.2 g of oxygen?

Solution

Step 1: Obtain the mass of each element in the compound. These are already given:

Mass of magnesium = 4.9 g

Mass of oxygen = 3.2 g

Step 2: Divide the mass of each element by its R.A.M.

Magnesium: 4.9 ÷ 24 = 0.204

Oxygen: 3.2 ÷ 16 = 0.2

Step 3: Divide by the lowest quotient.

Mg : O = 0.204 ÷ 0.2 : 0.2 ÷ 0.2 = 1.02 : 1

Step 4: Obtain their whole number ratios directly or by approximation.

Mg : O = 1 : 1

The empirical formula is MgO.

Nomenclature of Binary Inorganic Compounds

Everything in the universe bears a name to differentiate it from others. Chemical substances also bear names that range from those of elements to those of compounds. The name of a substance can originate from some factors such as the place of origin, founder, use, and type or classification. Items or substances that fall under a particular group or classification are named systematically. A systematic way of assigning names to items that belong to a particular group or classification is called nomenclature.

Scientists use IUPAC nomenclature in naming chemical compounds so that experts and other interested persons around the world understand exactly what the substance is. IUPAC stands for International Union of Pure and Applied Chemistry.

Binary Inorganic Compounds

While an inorganic compound is any substance in which two or more chemical elements (usually other than carbon) are combined, always in definite proportions, a binary compound is the one which is formed by two chemical substances. Examples of binary inorganic compounds are CaO, NaCl, and PCl₃.

Inorganic compounds are categorised into ionic and covalent. The nomenclature of ionic compounds differs slightly from that of covalent compounds.

Nomenclature of Binary Ionic Compounds

Ionic compounds are formed when a metal combines with a non-metal. The following are the steps considered when naming binary ionic compounds:

1. Name the metallic ion that appears first in the formula using the name of the element itself.
2. The second part of the formula which is usually an anion in the compound will end with a suffix "ide". For example, oxygen becomes oxide, hydrogen becomes hydride and chlorine becomes chloride.

Note

(a) Some metals always have fixed charges when they form ions, that is, Group I metals have a charge of +1, Group II metals have a charge of +2, Group III metals have a charge of +3, silver (Ag) has a charge of +1, and zinc (Zn) has a charge of +2.

(b) Other metals are multivalent and can thus form more than one ion. For example, iron (Fe) is bivalent; it has valencies of 2 and 3, copper (Cu) is also bivalent; it has valencies of 1 and 2. Compounds formed from these metals must be distinguished by stating which valency has been used in the compound. The valency of the respective metal is indicated by capital Roman numbers in parentheses (brackets), for example cobalt(II) chloride, copper(II) oxide and iron(III) oxide.

Example 3.9

What is the name of the compound with the formula FeCl₃?

Solution

(i) Let x be the charge of Fe

(ii) 1(x) + 3(-1) = 0

(iii) x - 3 = 0

(iv) x = +3

(v) So, the Fe is in the +3 oxidation state. Write the name 'iron' and place III in brackets beside it.

(vi) Use the name 'chlorine' but change the last three letters to "ide". So the name is iron(III) chloride.

Nomenclature of Binary Covalent Compounds

Covalent compounds are formed between two non-metal elements. These compounds are named differently from ionic compounds. The number of atoms are presented by prefixes as shown in the following table:

Number	Prefix	Number	Prefix
1	mono-	6	hexa-
2	di-	7	hepta-
3	tri-	8	octa-
4	tetra-	9	nona-
5	penta-	10	deca-

The following are the steps to consider when naming binary covalent compounds:

1. Give the name of the first element.
2. Give the name of the second element with the ending changed to -ide.
3. If more than one compound can form between two elements, use prefixes to indicate the number of atoms of each element.

Example

Give the name for PCl₃.

Answer

(i) Since there is one phosphorus atom, use it as the first part of the name.

(ii) There are three chlorine atoms, so use tri in front of chlorine; then, drop the -ine in chlorine and replace it with -ide.

The name is phosphorus trichloride.

Chemical Names of Common Substances

Chemical names are typically used to provide precise descriptions of substance compositions, including those encountered in daily life. For example, requesting sodium chloride for use in food is uncommon; instead, the term "common salt" is used. This explains the existence of common names for certain substances. However, it is important to note that some common names are inaccurate and may vary from one place to another. Therefore, they cannot tell the chemical composition of a substance.

Activity 3.3: To identify and classify the chemical compounds in toothpaste

Aim: To identify and classify the chemical compounds in toothpaste

Requirements: Toothpaste tube (with ingredient list), notebook, pen/pencil, and chart or table for recording observations

Procedure

1. Read the ingredients on the toothpaste tube.
2. Write down the chemical compounds listed in the ingredients.
3. Explore the chemical formulas of the compounds and identify whether each of the compound contains a metal and a non-metal or only non-metals.
4. Record your findings in a table.

Chapter Summary

1. Chemical bonding involves electrons in the outermost shell of an atom. When the outermost shell is fully filled, the atom is said to be stable.
2. Ions are formed when an atom gains or loses electron(s). Cations are positively charged ions that result from atoms losing one or more electrons. Anions are negatively charged ions that result from atoms gaining one or more electrons.
3. Ionic (electrovalent) bonding usually occurs between a metal and a non-metal. It involves the transfer of electron(s) from the atoms of the metal to the atoms of the non-metal.
4. Covalent bonding takes place between two or more non-metals. It involves atoms of the non-metals sharing electrons that are in their outermost shells.
5. The ability of an atom to combine with other atoms according to the number of electrons it can give, take or share is known as valency.
6. The oxidation state (oxidation number) of an element is the number of electrons that need to be added, shared or removed by its atom, to make a neutral molecule. The oxidation number is arbitrary and may be positive, negative or zero.
7. A radical is a group of atoms which behaves as a single unit and has an overall positive or negative charge. A radical can also be an atom, molecule or ion that has unpaired valence electron. Such a group maintains its identity throughout any chemical reaction.
8. A chemical formula is a representation that uses chemical symbols to show the proportions of the elements present in a chemical compound.
9. An empirical formula is the simplest way of writing a chemical formula and indicates the ratio of the atoms in a compound.
10. A molecular formula is a chemical formula that shows the total number of atoms of each element in a molecule of a substance.
11. A systematic way of naming items or substances of a particular category is known as nomenclature.

Revision Exercise 3

Choose the correct answer for Questions 1-11. For other questions, provide the answers as per the demands indicated.

1. What is the valency of Group I elements?

(a) 1 (b) 2 (c) 3 (d) 4

2. What name is given to the force of attraction that holds atoms together to form a molecule?

(a) Chemical change

(b) Chemical bond

(c) Friction

(d) Centripetal force

Answer the following questions:

12. Explain the difference between ionic and covalent bonding, giving two examples of each.

13. Calculate the oxidation number of sulfur in H₂SO₄.

14. Write the chemical formula for calcium phosphate.

15. Name the following compounds: (a) Fe₂O₃ (b) P₂O₅ (c) CuSO₄

Chemistry for Secondary Schools - Form Two

© MITIHANI POPOTE 2025

Chapter Four: Chemical Reactions

Chemistry for Secondary Schools - Form Two

Introduction

Chemical reactions are an integral part of daily life, constantly occurring in and around us, often unnoticed. For example, everyday activities such as burning charcoal and wood, cooking food, respiration in living organisms, fuel combustion in engines, rusting of metals, and digestion involve transforming substances into new ones. These transformations result from chemical reactions. In this chapter, you will learn the concept of chemical reactions, including chemical equations, and types of chemical reactions. The competencies developed will enable you to accurately present chemical equations and analyse various chemical reactions that yield products essential to daily life.

Think

The impact of chemical reactions on natural and industrial processes

Concept of Chemical Reactions

Task 4.1

Use library books and reliable online resources to search for any five chemical reactions that occur in daily life. For each response, include the name of the chemical reaction, a brief explanation of how it happens, the chemical equation (if applicable) and an example of where this reaction is frequently observed.

A chemical reaction is a process in which one or more chemical substances are converted to one or more different substances. Chemical reactions take place when bonds between atoms in the reacting substance(s) are broken, atoms rearrange, and new bonds between the atoms are formed to make new substance(s). The chemicals that begin the reaction process are called reactants and the new substances formed are called products. The products have properties that are different from their respective reactants.

Features Indicating a Chemical Reaction Has Occurred

• Evolution of a gas
• Formation of precipitates
• Change in colour
• Change in temperature
• Change in state

A chemical reaction is either reversible or irreversible. A reversible reaction proceeds in both forward and backward directions. In the forward reaction, the reactants are converted into products, whereas in the backward reaction, the products become the reactants. An irreversible reaction proceeds in one direction from reactants to products.

Chemical Equations

Task 4.2

1. Use library books and reliable online resources to explore various molecular equations.

2. Use chemistry software step by step to represent the molecular equations identified in 1.

A chemical reaction is expressed in the form of a chemical equation. A chemical equation is a symbolic or words representation of a chemical reaction. A chemical equation written in words is referred to as a word equation. For example, when calcium metal reacts with chlorine gas to form solid calcium chloride, the word equation is written as:

Calcium + Chlorine gas → Calcium chloride

A chemical equation written in symbols is referred to as a formula equation. For example, the formula equation for a reaction between calcium metal and chlorine gas is:

Ca(s) + Cl₂(g) → CaCl₂(s)

The formula equation is more useful than the word equation. However, it is advantageous to write the equation in word form first. Formula equations provide useful information, including compositions, amounts, formulas, and the physical states of substances that are involved. Formula equations may also state the conditions for the reactions to take place.

States of Matter in Chemical Equations

The reactants and products in the formula equations are either solids, liquids or gases. These states of matter are represented in a chemical equation by using symbols:

• (s) for solid
• (l) for liquid
• (g) for gas
• (aq) for aqueous (dissolved in water)

These symbols should be included when writing the chemical equations and are stated in parentheses after the chemical symbol.

Direction of Reaction

In a chemical equation, a headed arrow is used to separate the reactants and products.

A full-headed arrow is used to separate reactants and products for an irreversible reaction:

Reactants → Products

Double half-headed arrows pointing in the opposite directions are used to separate reactants and products for a reversible reaction:

Reactants ⇌ Products

Key Characteristics of Chemical Equations

• Reactants and products: It lists the substances involved in the reaction, with reactants on the left side and products on the right side.
• Chemical formulas: Each substance is represented using its chemical formula.
• Direction of reaction: An arrow points from the reactants to the products, indicating the direction of the reaction.
• Balanced equation: A chemical equation follows the law of conservation of mass, ensuring that the number of atoms for each element is the same on both sides of the equation.
• States of matter: The physical state of each substance is often indicated.
• Energy changes: If applicable, energy changes such as heat or light may be noted.
• Reaction conditions: Temperature, pressure or catalysts may be noted above or below the arrow.

Molecular Equations

A molecular equation is an equation representing a reaction showing the reactants and products in undissociated form. In molecular equations, reactants and products are considered neutral regardless of their exact physical states.

Examples of Molecular Equations

PbO₂(s) + SO₂(g) → PbSO₄(s)

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

In principle, when writing molecular equations, the reactants and products should be balanced.

Balancing Chemical Equations

All chemical equations must be written in accordance with the law of conservation of mass. This law states that, in a chemical reaction, the total mass of the products equals the total mass of the reactants. This means, when balancing a chemical equation, the number of each atom on both sides of the equation must be equal because atoms do not vanish during a reaction, but are reorganised.

Steps for Balancing Chemical Equations

Step 1: Write the equation in a word form.

Step 2: Write the unbalanced equation including correct chemical formulas for reactants and products.

Step 3: List the number of atoms of each element on both sides of the equation.

Step 4: Balance one element at a time. Usually start with metals or more complex elements, and complete by balancing hydrogen and oxygen if present.

Step 5: Use coefficients (whole numbers) to balance atoms. Never change subscripts in a chemical formula.

Step 6: Count atoms of all elements on both sides to make sure they are equal.

Step 7: Simplify coefficients if necessary. The final equation should use the smallest whole-number coefficients and should include the state symbols.

Example 4.1

Hydrogen chloride gas is formed when hydrogen gas burns in chlorine gas. Write a balanced chemical equation for the reaction.

Solution

Step 1: Write the equation in word form

Hydrogen gas + Chlorine gas → Hydrogen chloride gas

Step 2: Write the unbalanced chemical equation using symbols

H₂ + Cl₂ → HCl

Step 3: List the number of atoms of each element on both sides

Left side: H = 2, Cl = 2

Right side: H = 1, Cl = 1

Step 4: Balance the atoms

We need 2 HCl molecules to balance hydrogen and chlorine

Step 5: Use coefficients to balance

H₂ + Cl₂ → 2HCl

Step 6: Count atoms to verify balance

Left side: H = 2, Cl = 2

Right side: H = 2, Cl = 2

Step 7: Include state symbols

H₂(g) + Cl₂(g) → 2HCl(g)

Ionic Equations

An ionic equation is a chemical equation in which compounds in aqueous solutions or in molten states are written in terms of dissociated ions. Spectator ions are ions that do not change their valence states in a reaction. The net ionic equation is the result of an ionic equation without the spectator ions.

Example 4.5

Write a net ionic equation for the reaction between dilute hydrochloric acid and an aqueous sodium hydroxide.

Solution

Step 1: Write the balanced formula equation

HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

Step 2: Write the complete ionic equation

H⁺(aq) + Cl^-(aq) + Na⁺(aq) + OH^-(aq) → Na⁺(aq) + Cl^-(aq) + H₂O(l)

Step 3: Identify and remove spectator ions

Spectator ions: Na⁺ and Cl^-

Step 4: Write the net ionic equation

H⁺(aq) + OH^-(aq) → H₂O(l)

Types of Chemical Reactions

Table 4.3

Use library books and reliable online resources to explore various chemical reactions which occur in everyday life, and classify each one based on its reaction type.

Chemical reactions drive numerous everyday processes, such as cooking, energy production, plant growth, and digestion. These reactions convert raw materials into essential products that support life and modern living.

Combination (Synthesis) Reactions

A combination reaction is a chemical reaction in which two or more chemical species combine to form a single product.

A + B → AB

Example: Burning of magnesium in oxygen

2Mg(s) + O₂(g) → 2MgO(s)

Decomposition Reactions

A decomposition reaction is a chemical reaction in which a compound breaks down into its components.

AB → A + B

Example: Electrolysis of water

2H₂O(l) → 2H₂(g) + O₂(g)

Displacement Reactions

A displacement reaction is a chemical reaction in which a more reactive element displaces a less reactive element from its compound.

A + BC → AC + B

Example: Zinc with hydrochloric acid

Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)

Precipitation Reactions

A precipitation reaction is a chemical reaction in which two soluble substances combine to give an insoluble substance (precipitate).

AB(aq) + CD(aq) → AD(s) + CB(aq)

Example: Silver nitrate with sodium chloride

AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq)

Redox Reactions

A redox reaction is a chemical reaction in which the oxidation number of the participating chemical species changes.

Example: Copper(II) oxide with hydrogen

CuO(s) + H₂(g) → Cu(s) + H₂O(g)

Neutralisation Reactions

A neutralisation reaction is a chemical reaction between an acid and a base to give salt and water.

Acid + Base → Salt + Water

Example: Hydrochloric acid with sodium hydroxide

HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

Combination Reactions in Detail

There are three types of combination reactions:

1. Combination between two elements

This reaction occurs when two elements combine to give a single compound.

2Mg(s) + O₂(g) → 2MgO(s)

Activity 4.2: To burn magnesium in oxygen

Aim: To burn magnesium in oxygen

Requirements: A pair of tongs, heat source (Bunsen burner), gas jar, lighter, magnesium ribbon, oxygen source, and steel wool

Procedure

1. Clean about 0.1 g of a piece of magnesium ribbon by using steel wool.
2. Hold the ribbon by using a pair of tongs and heat it over a Bunsen burner or any heat source flame.
3. When it starts to burn, lower it into the gas jar of oxygen. Do not drop it into the jar.

Questions

1. What is the colour of flame when magnesium is burned?
2. What is the balanced chemical equation associated with this experiment?

2. Combination between elements and compounds

This occurs when an element reacts with a compound to form another compound.

2CO(g) + O₂(g) → 2CO₂(g)

3. Combination between two compounds

Two compounds may react with each other to form a new compound.

CaO(s) + CO₂(g) → CaCO₃(s)

Real-life Applications of Combination Reactions

• Formation of water: Hydrogen and oxygen gases combine to form water
• Rusting of iron: Iron reacts with oxygen and water to form hydrated iron(III) oxide (rust)
• Formation of calcium hydroxide: Calcium oxide reacts with water to form calcium hydroxide used in construction
• Photosynthesis: Carbon dioxide and water combine in plants to form glucose and oxygen
• Synthesis of ammonia: Nitrogen and hydrogen gases combine to form ammonia

Decomposition Reactions in Detail

Decomposition reactions are classified into three main types:

• Catalytic decomposition: Uses a catalyst to alter the rate of reaction
• Electrolytic decomposition: Achieved by exposing to electric current
• Thermal decomposition: Occurs when a compound is exposed to direct heat

Examples of Decomposition Reactions

Catalytic: Potassium chlorate with manganese(IV) oxide catalyst

2KClO₃(s) → 2KCl(s) + 3O₂(g)

Electrolytic: Electrolysis of water

2H₂O(l) → 2H₂(g) + O₂(g)

Thermal: Heating lead(II) nitrate

2Pb(NO₃)₂(s) → 2PbO(s) + 4NO₂(g) + O₂(g)

Real-life Applications of Decomposition Reactions

• Digestion of food: Complex food molecules break down into simpler substances
• Spoilage of food: Organic matter decomposes due to microorganisms
• Thermal decomposition of limestone: Used in cement industry

Combustion Reactions

A combustion reaction is a chemical process in which a substance reacts with oxygen to produce heat and light.

Fuel + O₂(g) → CO₂(g) + H₂O(l) + Energy

Examples:

• Burning of coal: C(s) + O₂(g) → CO₂(g) + Energy
• Cooking with natural gas: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) + Energy
• Candle burning: Hydrocarbon wax + O₂ → CO₂ + H₂O + Energy
• Respiration: C₆H₁₂O₆(aq) + 6O₂(g) → 6CO₂(g) + 6H₂O(l) + Energy

Displacement Reactions

Displacement reactions have various practical applications:

• Extraction of metals: More reactive metals displace less reactive metals from their compounds
• Purification of metals: Less reactive metals are displaced from their solutions by more reactive metals
• Rust prevention: Sacrificial protection using more reactive metals like zinc
• Treatment of wastewater: Metals displace harmful substances or ions from wastewater

Example: Extraction of iron

Fe₂O₃(s) + 2Al(s) → 2Fe(s) + Al₂O₃(s)

Chapter Summary

1. Chemical equations are short forms of chemical reactions. The reactants are placed on the left hand side of the equation, while the products are placed on the right hand side of the equation.
2. Chemical equations are written by using the chemical symbols or formulas. The state of each of the reactants and products is indicated in parentheses and the equations are balanced.
3. An ionic equation is a chemical equation in which compounds in aqueous solutions or in molten states are written in terms of dissociated ions.
4. Spectator ions are ions that do not change their valence states in a reaction.
5. The net ionic equation is the result of an ionic equation without the spectator ions.
6. There are different types of chemical reactions. These include combination, decomposition, displacement, precipitation, neutralisation and redox reactions.

Revision Exercise 4

Choose the correct answer for Questions 1–5. For other questions, provide the answers as per the demands indicated.

1. Which of the following statements describes chemical reactions?

(a) Occur only in living organisms.

(b) Occur in water only.

(c) Produce new substances.

(d) Only occur outside living organisms.

2. A certain metal hydroxide reacts with hydrochloric acid to produce salt and water only. What type of reaction is this?

(a) Precipitation

(b) Displacement

(c) Neutralisation

(d) Combination

Answer the following questions:

6. Differentiate oxidation from reduction in terms of electron transfer and changes in oxidation states.

7. Write a balanced chemical equation for each of the following reactions:

(a) Zinc reacts with silver nitrate solution to produce zinc nitrate and silver.

(b) Aqueous potassium iodide reacts with aqueous solution of lead(II) nitrate to produce potassium nitrate and lead(II) iodide.

8. Identify the type of reaction in each of the following chemical equations. Explain your answer:

(a) CuCO₃(s) → CuO(s) + CO₂(g)

(b) CaO(s) + H₂O(l) → Ca(OH)₂(aq)

(c) Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

Chemistry for Secondary Schools - Form Two

© MITIHANI POPOTE 2025

Chapter Five: Acids, Bases, and Salts

Chemistry for Secondary Schools - Form Two

Introduction

Acids, bases, and salts are fundamental chemical substances that play crucial roles in our daily lives. From the food we eat to the products we use for cleaning, these substances are everywhere. Understanding their properties, behavior, and applications helps us appreciate their importance in biological systems, industrial processes, and environmental management. In this chapter, you will learn about the properties of acids and bases, pH scale, acid-base indicators, neutralization reactions, and the preparation and uses of salts.

Think

How do acids and bases affect our daily lives from digestion to cleaning?

Acids

An acid is a chemical substance which produces hydrogen ions (H⁺) in water as the positively charged ions. Acids can be classified as mineral acids (inorganic acids) or organic acids based on their origin and composition.

Properties of Acids

Physical Properties

• Sour taste
• Corrosive nature
• Conduct electricity in aqueous solutions
• Turn blue litmus paper red

Chemical Properties

• React with metals to produce hydrogen gas
• React with bases to form salt and water
• React with carbonates to produce carbon dioxide
• Have pH values less than 7

Activity 5.1: Reactions of acids with metals

Aim: To investigate the reaction of acids with different metals

Requirements: Test tubes, dilute hydrochloric acid, dilute sulfuric acid, zinc granules, magnesium ribbon, copper turnings, lead shots, test tube rack, and splint

Procedure:

1. Place small pieces of each metal in separate test tubes.
2. Add 5 cm³ of dilute acid to each test tube.
3. Observe any reactions and test any gas produced with a burning splint.
4. Record your observations.

Questions:

1. Which metals reacted with the acids?
2. What gas was produced in the reactions?
3. Write balanced chemical equations for the reactions that occurred.

Strengths of Acids

The strength of an acid is a measure of its ability to ionize (dissociate) in water to produce hydrogen ions (H⁺). The more the acid ionizes in water, the stronger it is.

Strong Acids

Ionize completely in water to give large amounts of H⁺

• Hydrochloric acid (HCl)
• Sulfuric acid (H₂SO₄)
• Nitric acid (HNO₃)

HCl(aq) → H⁺(aq) + Cl^-(aq)

Weak Acids

Ionize partially in water to produce small amounts of H⁺

• Ethanoic acid (CH₃COOH)
• Carbonic acid (H₂CO₃)
• Citric acid (C₆H₈O₇)

CH₃COOH(aq) ⇌ CH₃COO^-(aq) + H⁺(aq)

Note

The pH values of solutions increase with decrease in acidity; therefore, weak acids have higher pH values than strong acids. Solutions of strong acids are good conductors of electricity because they contain more free mobile ions to carry the charges than solutions of weak acids.

Basicity of an Acid

Basicity of an acid is the number of ionizable hydrogen atoms per molecule of the acid that can be displaced by a metal in solution.

Type	Basicity	Examples	Ionization
Monobasic	1	HCl, HNO3, CH3COOH	H+ + A-
Dibasic	2	H2SO4, H2CO3	2H+ + A2-
Tribasic	3	H3PO4	3H+ + A3-

Bases

A base is a substance that neutralizes an acid by reacting with hydrogen ions. Bases include the oxides, hydroxides, and carbonates of metals.

Properties of Bases

Physical Properties

• Bitter taste
• Slippery or soapy feel
• Most are insoluble in water
• Turn red litmus paper blue

Chemical Properties

• Have pH values greater than 7
• React with acids to form salt and water
• React with ammonium salts to produce ammonia gas
• Soluble bases react with most cations to precipitate hydroxides

Alkalis

An alkali is a soluble base which, when dissolved in water, forms hydroxide ions (OH^-). Therefore, all alkalis are bases, but not all bases are alkalis.

Soluble Bases (Alkalis)	Insoluble Bases
Sodium hydroxide (NaOH)	Copper(II) hydroxide (Cu(OH)2)
Potassium hydroxide (KOH)	Lead(II) hydroxide (Pb(OH)2)
Calcium hydroxide (Ca(OH)2)	Copper(II) oxide (CuO)
Ammonium hydroxide (NH4OH)	Iron(III) oxide (Fe2O3)

Strengths of Bases

The strength of a base or alkali is its ability to ionize in aqueous solution to produce hydroxide ions (OH^-). The more the base ionizes in aqueous solution, the stronger it is.

Strong Bases

Ionize completely in aqueous solutions

• Sodium hydroxide (NaOH)
• Potassium hydroxide (KOH)
• Calcium hydroxide (Ca(OH)₂)

NaOH(aq) → Na⁺(aq) + OH^-(aq)

Weak Bases

Ionize partially in aqueous solutions

• Ammonia (NH₃)
• Aluminium hydroxide (Al(OH)₃)
• Magnesium hydroxide (Mg(OH)₂)

NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH^-(aq)

Activity 5.2: To identify acidic and basic substances

Aim: To identify acidic and basic substances

Requirements: Coconut, citrus fruits, ripe tomatoes, vinegar, sour milk, wood ash, pieces of chalk, yoghurt, blender, knife, water, red litmus papers, blue litmus papers, and beakers

Procedure:

1. Prepare juices of coconut, citrus fruits, and ripe tomatoes.
2. Use litmus papers to test the acidity or basic nature of the juices obtained.
3. Mix some wood ash and chalk powder, each with little water, decant to obtain a supernatant and use litmus papers to test.
4. Take small amounts of vinegar, sour milk, and yoghurt into separate beakers and test for acidity or basic nature using litmus papers.

Questions:

1. Which substances are acidic and which ones are basic?
2. How relevant is this activity to your daily life?

Neutralisation Reaction

If an acid and a base are mixed together in correct amounts, a neutral solution is produced. A reaction of this type is called neutralisation. Neutralisation is therefore a reaction between an acid and a base to produce salt and water.

Acid + Base → Salt + Water

Neutralisation reactions that involve the reaction of an acid with a carbonate or hydrogencarbonate produce carbon dioxide gas in addition to salt and water.

Acid + Carbonate/Hydrogencarbonate → Salt + Water + Carbon dioxide

Examples of Neutralisation Reactions

Example 1: Acid with alkali

HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

Example 2: Acid with metal oxide

H₂SO₄(aq) + CuO(s) → CuSO₄(aq) + H₂O(l)

Example 3: Acid with carbonate

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

Applications of Neutralisation Reactions

• Treating insect stings and bites: Bee stings (acidic) neutralized with baking soda; wasp stings (alkaline) neutralized with vinegar
• Relieving indigestion: Antacids neutralize excess stomach acid
• Soil treatment: Adding lime to acidic soils or sulfur to alkaline soils
• Preventing formation of acid rain: Using bases in air pollution control devices
• Treating factory wastes: Neutralizing acidic or alkaline industrial effluents
• Manufacturing fertilisers: Production of ammonium salts through neutralisation

Activity 5.3: To investigate the neutralisation reaction

Aim: To investigate the neutralisation reaction between sodium hydroxide and dilute hydrochloric acid

Requirements: Pipette, burette, retort stand and clamp, measuring cylinder, conical flasks, beakers, evaporating dish, white tile, heat source, tripod stand, wire gauze, 0.1 M sodium hydroxide, 0.1 M hydrochloric acid, and phenolphthalein indicator

Procedure:

1. Transfer 25 cm³ of sodium hydroxide solution into a conical flask using a pipette.
2. Add 2 to 3 drops of phenolphthalein indicator to the solution.
3. Fill the burette with dilute hydrochloric acid and record the initial volume.
4. Run the acid into the alkali until the indicator just changes color.
5. Record the volume of acid used.
6. Mix the same volumes of base and acid without indicator.
7. Evaporate the solution to obtain crystals.

Questions:

1. Why is the phenolphthalein indicator important in this experiment?
2. What is the balanced chemical equation for this reaction?
3. What is the ionic equation of the reaction?

Acid-Base Indicators

Acid-base indicators are chemicals that are used to determine the chemical nature of a substance; whether it is acidic or basic. The acid-base indicators are also known as pH indicators because acidity and alkalinity relate to pH range.

pH Scale

pH scale is a scale of numbers, from 0 to 14 which is used to express acidity, neutrality, and alkalinity. The pH 7 indicates neutrality of a substance. Any substance that has a pH below 7 is acidic, while a substance with pH above 7 is basic.

0

1

2

3

4

5

6

7

8

9

10

11

12

13

14

• pH 0-6: Acidic (lower numbers = stronger acids)
• pH 7: Neutral
• pH 8-14: Basic (higher numbers = stronger bases)

Types of Indicators

Indicator	Color in Acid	Color in Base	pH Range
Phenolphthalein	Colorless	Pink	8-10
Methyl orange	Red	Yellow	3-5
Litmus	Red	Blue	5-8
Bromothymol blue	Yellow	Blue	6-8
Universal indicator	Red → Yellow	Green → Violet	1-14

Natural Indicators

Natural pH indicators are substances found naturally that can be used to determine whether a substance is acidic or basic.

• Red cabbage: Pink in acids, green in bases
• Turmeric: Yellow in acids, red-brown in bases
• Hibiscus flowers: Red in acids, greenish-yellow in bases
• Beetroot: Red in acids, yellow in bases
• Onion: Smell persists in acids, diminishes in bases

Activity 5.4: To prepare a pH indicator from hibiscus flowers

Aim: To prepare a pH indicator from hibiscus flowers

Requirements: Hibiscus flowers, water, beakers, droppers, mortar and pestle, knife, 50% ethanol solution, heat source, evaporating dish, dilute hydrochloric acid, and dilute sodium hydroxide

Procedure:

1. Cut the hibiscus flowers into small pieces and crush them using mortar and pestle.
2. Add about 20 cm³ of ethanol solution into the crushed hibiscus flowers and continue crushing.
3. Decant the mixture and keep the solution obtained aside.
4. Boil the solution to evaporate ethanol, then leave the residue to cool.
5. Test the indicator with acids and bases.

Salts

A salt is an ionic compound which is made up of positively charged ions (cations) and negatively charged ions (anions). Salts occur naturally and are also synthesised in the laboratory.

Types of Salts

Neutral Salts

Formed from strong acids and strong bases

pH = 7

• NaCl
• KCl
• K₂SO₄

Acidic Salts

Formed from strong acids and weak bases

pH < 7

• NH₄Cl
• FeCl₃
• NaHSO₄

Basic Salts

Formed from weak acids and strong bases

pH > 7

• CH₃COONa
• Na₂CO₃
• NaHCO₃

Preparation of Salts

1. Reactions of acids with metals (Displacement method)

Metal + Acid → Salt + Hydrogen gas

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

2. Reactions of acids with bases (Neutralisation method)

Acid + Base → Salt + Water

Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

3. Reactions of acids with carbonates

Acid + Carbonate → Salt + Water + Carbon dioxide

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

4. Preparation of insoluble salts (Precipitation method)

Soluble salt + Soluble salt → Insoluble salt + Soluble salt

Example: Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq)

Properties of Salts

• Physical appearance: Crystalline or powder forms with different colors
• Solubility: Varies with temperature and the specific salt
• Action of heat: Different salts decompose at different temperatures
• Deliquescence: Some salts absorb moisture from air to form solutions
• Hygroscopy: Some salts absorb water from air without dissolving
• Efflorescence: Some salts lose water of crystallization when exposed to air

Uses of Salts

Salt	Uses
Sodium chloride (NaCl)	Food seasoning, preservation, manufacturing soaps
Sodium carbonate (Na2CO3)	Water softening, glass manufacturing, cleaning agent
Sodium hydrogen carbonate (NaHCO3)	Baking, antacid, fire extinguishers
Ammonium chloride (NH4Cl)	Dry cell batteries, fertilizer
Copper(II) sulfate (CuSO4)	Fungicide, electroplating, correction of copper deficiency
Calcium carbonate (CaCO3)	Construction material, antacids, manufacturing paints

Activity 5.5: To investigate the preservative effects of table salt

Aim: To investigate the preservative effects of table salt

Requirements: Table salt (sodium chloride), pieces of fresh meat, beaker, petri dishes

Procedure:

1. Take two pieces of fresh meat.
2. Sprinkle a generous amount of salt on one piece and leave the other without salt.
3. Leave them at room temperature for 24 hours.
4. Observe and record any appearance, texture, and odour.

Questions:

1. What happened to the salted and unsalted pieces of meat?
2. What can you conclude on the effects of the salt in the two pieces of meat?

Chapter Summary

1. An acid is a chemical substance which produces hydrogen ions (H⁺) in water as the positively charged ions.
2. A base refers to a chemical substance which, when dissolved in water, produces hydroxide ions (OH^-) as negatively charged ions. An alkali is a soluble base.
3. A strong acid or strong base dissociates completely in water. A weak acid or weak base dissociates partially in water.
4. The basicity of an acid is the number of hydrogen atoms per molecule of the acid that can be displaced by a metal in a solution.
5. A pH indicator is a chemical substance that exhibits different colours in solutions of different acidities or alkalinities.
6. Neutralisation is a reaction between an acid and a base to produce salt and water.
7. A salt is an ionic compound made up of positively charged ions (cations) and negatively charged ions (anions).
8. Salts can be neutral, acidic, or basic depending on the strengths of the parent acid and base.
9. Salts have various uses in daily life including food preservation, medicine, agriculture, and industry.

Revision Exercise 5

Choose the correct answer for Questions 1–6. For other questions, provide the answers as per the demands indicated.

1. Which of the following is a natural acid?

(a) Nitric acid

(b) Phosphoric acid

(c) Citric acid

(d) Sulfuric acid

2. An acid is any substance that:

(a) contains hydrogen in its formula.

(b) dissolves in water to produce hydrogen ions.

(c) can react with NaOH.

(d) contains oxygen in its formula.

Answer the following questions:

7. Match the premises on properties and uses of acids and bases in Column A with their correct responses in Column B.

8. Explain the following terms and illustrate their significance by providing an example for each:

(a) pH indicators

(b) Neutralisation

(c) Basicity of an acid

9. How are salts useful in our life in the following aspects?

(a) Food preservation and preparation

(b) Agricultural practices

(c) Medical applications

Project: Determination of pH values of natural and synthetic substances

Objective: Collect samples of various substances such as tap water, bottled water, river water, fruits juices, and soaps and detergents. Determine the pH values using natural and synthetic acid-base indicators, litmus papers, or pH meter.

Tasks:

1. Write a detailed report outlining the methodology, materials used, observations, results, discussions, and conclusions regarding the pH properties of the substances.
2. Prepare a brief presentation summarising the experiments, results and key learnings.
3. Enhance the presentation with visuals such as photos of the tests or a pH chart derived from the collected data.

Expected Outcomes:

• Understanding of pH values of common substances
• Ability to use different types of indicators
• Appreciation of the importance of pH in daily life

Chemistry for Secondary Schools - Form Two

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