Carbon and Its Allotropes – at a glance

 

Carbon is found in Group IV, Period II of the periodic table. Its electronic configuration is 1s² 2s² 2p². It occurs naturally in different forms called allotropes.

Allotropy

Allotropy is the ability of an element to exist in two or more different forms but in the same physical state.

I. Crystalline allotropes: Diamond, Graphite, Fullerenes

Ii. Amorphous forms: Coal, Charcoal, Coke, Soot, 

Uses: cutting tools, drilling, jewelry, precision instruments.

Crystalline Allotropes of carbon 

1. Diamond

Each carbon atom is bonded to 4 other carbon atoms in a tetrahedral structure. Making diamond a giant molecule with an octahedral shape (Forms a rigid 3D network.)

Properties of diamond 

i. It is the Hardest natural substance known 

ii. High melting point

iii. It does not conduct electricity 

iv. It is Transparent and shiny

Uses:

I. It is used for making Cutting tools

II. It is used as Jewelry

III. It is used for making Industrial drilling machines 

2. Graphite

Graphite has flat layers of carbon atoms with free electrons. Each carbon atom in graphite bonds with 3 others forming hexagonal layers.

Each Layers are weakly held together by weak van der walls forces of attraction which makes it easy for one layer to easily slide over another (causing graphite to flake easily)

Properties of graphite 

i. Soft and slippery

ii. Good Conductors of electricity (due to free electrons)

iii. It is Black and opaque

iv. It has a high melting point

Uses:

i. It is used in making lead Pencil 

ii. It is used as a dry Lubricants

iii. It is used as Electrodes in batteries

iv. It is used for lining the inside of crucibles 

3.  Fullerenes (Modern Allotrope)

Fullerenes (e.g. C₆₀) are spherical shape (like a football) carbon molecules called buckyballs (Buckminsterfullerene) (C₆₀) or C₇₀ and others – elongated shapes

 The carbon atoms are arranged in closed hollow structures such as spheres, ellipsoids, or tubes.

Structure

 i. Carbon atoms are arranged in hexagons and pentagons

ii. The most common fullerene is Buckminsterfullerene (C₆₀), which has a spherical shape like a football

iii. Each carbon atom forms three covalent bonds (sp² hybridization)

Properties

I. They have light weight but are very strong.

ii. They are Good electrical conductivity

iii. They have High stability and 

iv. They can act as antioxidants

Uses

i.  Nanotechnology and electronics      

ii. Drug delivery in medicine

iii. They are used as Lubricants

iv. They are used as Superconductors (in some modified forms)

Carbon Nanotubes   (cylindrical fullerenes)

Cylindrical tubes made of graphene

Properties 

i. They are Very strong

ii. The have Good electrical conductivity

Uses:

I. They are used in Nanotechnology

II. Used in Medicine (for drug delivery)

III. Used in making Electronics

Graphene

Structure: This is a single layer of graphite (one atom thick)

Properties of graphene 

i. It is Extremely strong

ii. It is an Excellent conductor of heat and electricity

iii. It is Flexible and lightweight

Uses:

i. It is used in Electronics

ii. It is used in Sensors

iii. It is used in Advanced materials

 Amorphous Carbon

i. Charcoal – 

a. Wood charcoal absorbs 

b.  sugar charcoal

c. animal charcoal 

Uses 

i. absorbs gases and colours

ii. as fuel 

ii. Carbon black & lampblack – used in tyres, inks and polish

iii. Coal – used mainly as fuel

 Types of Coal (check post on coal): there are four stages of coal

• Peat – about 60% carbon
• Lignite – about 67% carbon
• Bituminous – about 88% carbon
• Anthracite – about 94% carbon (hardest and purest)

 Destructive Distillation of Coal Heating:     

This is when coal is heated to a very high temperature in the absence of air. 

The products got are 

Coal → Coal gas + Coal tar + Ammoniacal liquor + Coke 

Gasification of coke:   when coke is heated to red hot and white hot and air and steam blown over it, it produces two types of gases (fuel gases)

 i. Producer gas: produced when air is blown over white hot coke      C_(s) → CO + N₂

ii.Water gas: Produce when steam is blown over red-hot coke           C(s) →    CO + H₂

iii. Synthetic gas – CO + H₂

Why Carbon Forms Many Allotropes

Carbon’s ability to form many allotropes is due to:

i. Catenation (bonding with itself)

ii. Ability to form different bond types (single, double)

iii. Flexibility in bonding arrangements

 Chemical Properties of Carbon

i.  Carbon burns in oxygen to form CO₂ or CO

  C_(s) + O_2(s) → CO_(g)  (limited oxygen)

   C_(s) + O_{2(s),/sub. → CO2(s)   (excess oxygen)}  

ii. Combines with elements like sulphur and hydrogen

 

C_(s) + S_(s) → CS₂

C_(s) + H_{2(g) →}   CH₄

iii. Acts as a reducing agent in metal extraction

 iv. Fe₂O₃ + C_{(s) →}  Fe_(s) + CO_2(g)

Is oxidized by strong acids to form CO₂

2H₂SO_4(aq) + C(s) → 2SO_2(g) + 2H₂O_(g) + CO_2(s)

Objective Questions (Carbon & Its Allotropes)

1. Carbon belongs to which group in the periodic table?
   A. Group I
   B. Group II
   C. Group IV
   D. Group VI

2. The atomic number of carbon is:
   A. 4
   B. 6
   C. 12
   D. 14

3. The ability of carbon to form long chains is called:
   A. Isomerism
   B. Catenation
   C. Polymerization
   D. Hybridization

4. Which of the following is an amorphous form of carbon?
   A. Diamond
   B. Graphite
   C. Charcoal
   D. Fullerene

5. Diamond is hard because:
   A. It contains free electrons
   B. It has strong covalent bonds in a 3D network
   C. It is metallic
   D. It contains impurities

6. Which allotrope of carbon conducts electricity?
   A. Diamond
   B. Graphite
   C. Charcoal
   D. Coke

7. In graphite, each carbon atom is bonded to how many other carbon atoms?
   A. 2
   B. 3
   C. 4
   D. 6

8. The structure of diamond is:
   A. Layered
   B. Planar
   C. Tetrahedral
   D. Linear

9. Which allotrope of carbon is used as a lubricant?
   A. Diamond
   B. Graphite
   C. Coal
   D. Coke

10. The black soot obtained from incomplete combustion is called:
    A. Coke
    B. Charcoal
    C. Lampblack
    D. Coal tar

11. Which allotrope of carbon is the hardest natural substance?
    A. Graphite
    B. Diamond
    C. Coke
    D. Charcoal

12. Fullerenes are composed of carbon atoms arranged in:
    A. Chains
    B. Sheets
    C. Spherical shapes
    D. Cubes

13. Which of the following is NOT an allotrope of carbon?
    A. Graphene
    B. Diamond
    C. Silicon
    D. Fullerene

14. The valency of carbon is:
    A. 2
    B. 3
    C. 4
    D. 6

15. Coal is mainly composed of:
    A. Hydrogen
    B. Oxygen
    C. Carbon
    D. Nitrogen

16. Which allotrope has a layered structure with weak forces between layers?
    A. Diamond
    B. Graphite
    C. Fullerene
    D. Charcoal

17. Which of the following is used in cutting tools?
    A. Graphite
    B. Diamond
    C. Coke
    D. Coal

18. The presence of free electrons in graphite makes it:
    A. Hard
    B. Transparent
    C. Conductive
    D. Brittle

19. Which form of carbon is used in water purification?
    A. Activated charcoal
    B. Diamond
    C. Graphite
    D. Fullerene

20. The difference between diamond and graphite is mainly due to:
    A. Atomic number
    B. Number of electrons
    C. Arrangement of atoms
    D. Chemical composition

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NATURE OF MATER

What is Matter: -
 Mater is defined as anything that has mass and occupies space.

Composition of matter: - matter consist of any one of the following particles. To

1. Atoms.  

2. Molecules.  Or

 3. Ions

1. An Atom is the smallest particle of an element that can take part in a chemical reaction.

2. A molecule is the smallest particle of a substance that exist alone and still possess the properties of the substance. E.g H2O, O

3. An Ion is a charged particle; it is formed when an atom loses or gains electrons.  E.g Na+, Cl- 

Radicals are group of atoms with a single charge.  E.g NO-, SO42-, OH-

Matter generally is made up of any one or more of the particles mentioned above.

States of matter

matter can exist in three states.

i. the solid

ii. the liquid and

iii. the gaseous state.

SOLID: - In solids the particles of matter are densely packed and are held by strong forces of attraction (force of cohesion).

Properties of solids 

i., solids have fixed, or definite volumes

ii. solids also have fixed shapes.

iii solids cannot be compressed

                                      Solid state (particles are tightly packed and are held by strong forces) 

LIQUIDS: -In liquids the molecules (particles) are close together in an orderly manner with little freedom of movement.  Molecules in a liquid are close together but are not held so

rigidly in position and can move past one another.

Properties of Liquids

i.  a liquid no fixed or definite shape but it takes the shape of its container, 

ii. liquids have a fixed or definite volume 

iii. liquids cannot be compressed.

                                              

                                              Liquid state (particles are not held tightly together)

GAS: - In a gas, the particles/molecules are separated by distances that are large compared with the size of the molecules.

Properties of gases

i.; gasses have no fixed or definite volume (will occupy entire volume of its container) 

ii. gases have no fixed shape.

iii.   gases can be compressed

 Gases differ from liquids and solids in the distances between their individual particles.

 Gases (particles of gases are wide apart)

The three states of matter can be inter-convertible without changing the composition of the substance. Upon heating, a solid it will change into a liquid form at a particular temperature.

The temperature at which this transition occurs is called the melting point of the solid. Further heating will convert the liquid into a gas at a particular temperature called the boiling point of the liquid. Cooling a gas on the other hand, will cause it to change into a liquid, this process is called Condensation. When the liquid is cooled further, it will change into the solid form, and the process is known as Freezing. 

Some solids change directly to the gaseous without passing through the liquid state and vice Versal. These processes are both known Sublimation and Deposition.

          Changes in the state of Matter

   

  Now remember we defined chemistry as the study of the composition, properties and the uses of mater as well as the principle that underlies the changes that matter undergoes, and so far, we are told that mater consist of one or any two of atoms, molecules or ions. now the properties of matter can be divided into two, that is, physical properties and chemical properties.

PROPERTIES OF MATTER 

    The properties of matter can be grouped into two

i. Physical properties are those properties that a substance (matter) shows by itself (i.e those properties that you see when you look at, touch or smell the substance) these include Temperature, Pressure, Colour, Smell, Density, Solubility, Melting and boiling points.

ii. Chemical properties on the other hand are those properties that matter exhibits when it comes in contact with other substances like acids, bases, even water to form new substances.

Physical and Chemical change

A physical change is one that is easily reversible and in which no new substance is formed.

example of physical changes is.

i. Dissolving sodium chloride in water

ii. Changing water to ice

iii. Evaporation of liquids

iv. Melting of candle wax

A chemical change is one which is not easily reversible and in which a new substance is always formed.

example of changes is.

i.   burning of wood, paper, or clothe.

ii.  rusting of iron when exposed to air and moisture.

iii. decomposition of organic matter.

iv. dissolving a metal like calcium or magnesium in dilute acids.

v.   the hardening of cement   when exposed to air.

OBJECTIVE QUESTIONS

1. Which of the following is a physical change 

a). Dissolving sugar on water

b). Addition of acid to base 

c). Burñing of wood

d). Rusting of iron

2. 

THEORY QUESTIONS

1(i). State two differences between the properties of solids and gasses

 (ii). What process does each of X, Y and Z represent in the changes shown below?

Great 👍 — I’ve improved the blog note, added SEO elements, examples, and practice questions so it will be more useful for your readers (especially students).

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Nature of Matter

Introduction

The nature of matter is a fundamental concept in chemistry that explains what matter is made of and how it behaves. Everything around us — from the air we breathe to the food we eat — is composed of matter. Understanding the nature of matter helps scientists explain physical and chemical changes in substances.

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What is Matter?

Matter is anything that has mass and occupies space.

Mass refers to the amount of substance present, while space occupied by matter is known as volume.

Examples of Matter

• Water in a bottle
• Air inside a balloon
• Sand on a beach
• A wooden table

All these substances have mass and occupy space, therefore they are matter.

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Characteristics of Matter

Matter possesses several important characteristics:

1. Matter Has Mass

Mass is the quantity of matter in a body. It is measured using a balance and expressed in grams (g) or kilograms (kg).

2. Matter Occupies Space

The space that matter occupies is called volume. For example, water poured into a container fills the available space.

3. Matter Is Made of Tiny Particles

Matter is made up of extremely small particles called:

• Atoms
• Molecules
• Ions

These particles are too small to be seen with the naked eye.

4. Particles of Matter Are Always Moving

The particles that make up matter are in constant motion. The speed of this movement increases when the temperature rises.

5. Particles of Matter Attract Each Other

There are forces of attraction between particles of matter which hold them together.

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Particle Nature of Matter

The particle theory of matter states that:

1. Matter is made up of tiny particles.
2. The particles have spaces between them.
3. The particles are constantly moving.
4. The particles attract one another.

This theory helps explain processes such as diffusion, dissolving, and changes of state.

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States of Matter

Matter exists mainly in three states.

1. Solid

Solids have definite shape and definite volume. The particles are tightly packed and can only vibrate in fixed positions.

Examples

• Salt
• Iron
• Wood
• Stone

Properties of Solids

• Fixed shape
• Fixed volume
• High density
• Particles closely packed

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2. Liquid

Liquids have definite volume but no definite shape. They take the shape of the container in which they are placed.

Examples

• Water
• Oil
• Alcohol
• Kerosene

Properties of Liquids

• Definite volume
• No fixed shape
• Flow easily
• Particles loosely packed

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3. Gas

Gases have no definite shape and no definite volume. Their particles are far apart and move freely.

Examples

• Oxygen
• Carbon dioxide
• Nitrogen
• Air

Properties of Gases

• No fixed shape
• No fixed volume
• Easily compressed
• Particles move rapidly

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Changes in the State of Matter

Matter can change from one state to another when heat energy is added or removed.

Change	Description	Example
Melting	Solid → Liquid	Ice turning to water
Freezing	Liquid → Solid	Water forming ice
Evaporation	Liquid → Gas	Water forming vapor
Condensation	Gas → Liquid	Steam forming water droplets
Sublimation	Solid → Gas directly	Camphor or naphthalene

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Importance of Studying the Nature of Matter

Understanding the nature of matter helps us to:

• Understand chemical reactions
• Study properties of substances
• Explain changes of state
• Develop new materials and medicines
• Improve processes in industry and technology

It is also important in fields like medicine, environmental science, engineering, and materials science.

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Conclusion

The nature of matter explains the structure, properties, and behavior of substances. Since matter is made up of tiny particles that are constantly moving and attracting each other, many natural phenomena such as diffusion, evaporation, and chemical reactions can be explained through this concept.

Objective Questions

1. Matter is anything that has ______ and occupies space.
2. The space occupied by matter is called ______.
3. The smallest particles that make up matter are called ______.
4. The state of matter with definite shape and volume is ______.
5. The process by which a liquid changes to gas is called ______.

Theory Questions

1. Define matter.
2. State four characteristics of matter.
3. Explain the particle nature of matter.
4. Describe the three states of matter.
5. Explain two changes of state with examples.

 

FATS AND OILS – COMPLETE STUDENT NOTE

Introduction

Fats and oils are lipids, a class of organic compounds that are insoluble in water but soluble in organic solvents like alcohol or ether.

They are important macronutrients in food, providing energy, insulation, and cell structure.

Composition of Fats and Oils

Fats and oils are composed of:

1. Glycerol (C₃H₈O₃) – a 3-carbon alcohol

2. Fatty acids – long chains of carboxylic acids (R-COOH)

General formula of a triglyceride (fat/oil):

Glycerol+ 3 Fatty acids → Triglyceride + 3H₂O

This is formed through a condensation reaction (esterification) where water is released.

Hence Fats and Oils are triglycerides (esters) of long-chain carboxylic acids (fatty acids and propan-1,2,3-triol (glycerol)

Types of Fats and Oils

1. Saturated Fats

• No double bonds between carbon atoms in fatty acid chains

• Usually solid at room temperature

• Found in: Butter, lard, coconut oil

• Example: Stearic acid (C₁₇H₃₅COOH)

2. Unsaturated Fats

• Have one or more double bonds in the fatty acid chains

• Usually liquid at room temperature (oils)

• Found in: Vegetable oil, olive oil, fish oil

• Example: Oleic acid (C₁₇H₃₃COOH)

Subtypes:

• Monounsaturated – 1 double bond

• Polyunsaturated – 2 or more double bonds

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Structural Formulae

1. Triglyceride (generalized)

       H   H   H
       |   |   |
HO–CH2–CH–CH2–OH (Glycerol backbone)
       |   |   |
       R1  R2  R3 (Fatty acids)

2. Example – Triolein (unsaturated oil)

CH2–O–CO–C17H33
CH–O–CO–C17H33
CH2–O–CO–C17H33

3. Example – Tristearin (saturated fat)

CH2–O–CO–C17H35
CH–O–CO–C17H35
CH2–O–CO–C17H35

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Properties of Fats and Oils

1. they are insoluble in water

2. they are soluble in organic solvents like ether and chloroform

3. they have High calorific value (~9 kcal/g)

4. they are Non-polar molecules

Chemical Properties

1. they reaction with alkali to form soap (saponification reaction)

2.

Functions of Fats and Oils

1. Energy source – They provide twice as much energy as carbohydrates per gram

2. Insulation – they maintain body temperature

3. Protection – they cushion internal organs

4. Cell structure – phospholipids in cell membranes

5. Vitamin absorption – they help in absorption of fat-soluble vitamins (A, D, E, K)

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Examples of Fats and Oils

Type	Example	Source
Saturated	Stearic acid	Animal fat, butter
Saturated	Palmitic acid	Palm oil, meat fat
Unsaturated	Oleic acid	Olive oil, groundnut oil
Unsaturated	Linoleic acid	Sunflower oil, corn oil

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Tests for Fats and Oils

1. Solubility test:

   • Insoluble in water, soluble in alcohol

2. Emulsion test:

   • Shake sample with ethanol, then add water → milky emulsion indicates fat

3. Saponification test:

   • Heat fat with NaOH → soap + glycerol

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Summary

• Fats and oils = triglycerides (glycerol + 3 fatty acids)

• Saturated fats → solid, no double bonds

• Unsaturated fats → liquid, one or more double bonds

• Functions: energy, insulation, protection, vitamin absorption

• Can be tested using solubility, emulsion, and saponification tests

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Objective Questions

1. Fats and oils are classified as:
A. Carbohydrates
B. Proteins
C. Lipids
D. Nucleic acids

2. The monomer unit of fats and oils is:
A. Glucose
B. Glycerol + fatty acids
C. Amino acid
D. Nucleotide

3. The reaction by which glycerol reacts with fatty acids to form a triglyceride is called:
A. Hydrolysis
B. Condensation / Esterification
C. Oxidation
D. Saponification

4. Which of the following fats is saturated?
A. Olive oil
B. Fish oil
C. Butter
D. Sunflower oil

5. Which of the following is an unsaturated fatty acid?
A. Stearic acid
B. Palmitic acid
C. Oleic acid
D. Butyric acid

6. Fats and oils are:
A. Soluble in water
B. Insoluble in water but soluble in organic solvents
C. Ionic compounds
D. Proteins

7. Which test is used to detect fats and oils?
A. Benedict’s test
B. Biuret test
C. Emulsion test
D. Iodine test

8. A fat that is solid at room temperature is usually:
A. Unsaturated
B. Saturated
C. Polyunsaturated
D. Monounsaturated

9. Which products are formed when fats undergo saponification?
A. Soap + Glycerol
B. Water + Glycerol
C. Soap + Fatty acids
D. Alcohol + Soap

10. The function of fats and oils in the body includes:
A. Energy storage
B. Insulation and protection
C. Vitamin absorption
D. All of the above

Theory Questions

Short Answer

1. Define fats and oils.

2. State the difference between saturated and unsaturated fats.

3. Write the general formula for a triglyceride.

4. Name two sources of saturated fats and two sources of unsaturated fats.

5. Explain why fats and oils are insoluble in water.

Structured / Calculation-Oriented

6. Draw the general structure of a triglyceride.

7. Describe the formation of a triglyceride from glycerol and three fatty acids.

8. Explain the emulsion test for fats and oils.

9. State the products of saponification of fats using NaOH.

10. Compare the physical state at room temperature of saturated and unsaturated fats.

Higher-Level / Application

11. Explain the importance of fats and oils in the human diet.

12. A student observes that a sample of fat is solid at room temperature. Identify the type of fat and give one example.

13. Why are unsaturated fats generally considered healthier than saturated fats?

14. Describe how triglycerides can be hydrolyzed in the body to release energy.

15. Compare the chemical structure of stearic acid and oleic acid, indicating the difference in saturation.

PROTEINS – COMPLETE NOTE

Introduction

Proteins are essential biological molecules made of amino acids linked together by peptide bonds.
They are one of the macronutrients needed for growth, repair, and maintenance of the body.

Proteins play important roles in living organisms, such as 

• Building and repairing tissues

• Acting as enzymes and hormones

• Supporting the immune system

• Providing energy when necessary

Structure of Proteins

Proteins have four levels of structure:

1. Primary Structure

• The sequence of amino acids in a polypeptide chain.

• Determines the unique characteristics of the protein.

2. Secondary Structure

• Folding of the polypeptide chain into patterns like:

  • Alpha-helix (α-helix)

  • Beta-pleated sheet (β-sheet)

• Stabilized by hydrogen bonds.

3. Tertiary Structure

• Three-dimensional shape of the protein.

• Determined by interactions among R-groups (side chains) of amino acids.

4. Quaternary Structure

• Present in proteins with more than one polypeptide chain.

• Example: Hemoglobin (made of 4 polypeptide chains).

Composition of Proteins

Proteins are made of carbon (C), hydrogen (H), oxygen (O), and nitrogen (N).
Some proteins also contain sulfur (S), phosphorus (P), or iron (Fe).

Types of Proteins

1. Simple Proteins

   • Made entirely of amino acids.

   • Example: Albumin, Globulin

2. Conjugated Proteins

   • Contain a non-protein group (prosthetic group) in addition to amino acids.

   • Example: Glycoprotein (protein + carbohydrate), Hemoglobin (protein + heme)

Functions of Proteins

 

Function	Example/Explanation
Structural	Collagen in connective tissues, keratin in hair
Enzymatic	Amylase, lipase (catalyze reactions)
Hormonal	Insulin, glucagon (regulate metabolism)
Transport	Hemoglobin (carries oxygen)
Immune	Antibodies (fight infection)
Storage	Ferritin (stores iron)
Energy	Proteins can be broken down for energy if carbs/fats are low

Classification Based on Shape

• Fibrous Proteins: Long, strand-like, structural proteins.

  • Example: Keratin, Collagen

• Globular Proteins: Spherical, functional proteins.

  • Example: Enzymes, Hemoglobin, Antibodies

AMINO ACIDS

Amino acids are the building blocks of proteins. Each amino acid has the general formula:

H₂N-CH(R)-COOH

Where R = side chain unique to each amino acid

Examples of Amino Acids

• Essential amino acids (cannot be made by the body):

  • Lysine, Methionine, Leucine, Valine, etc.

• Non-essential amino acids (can be made by the body):

  • Glycine, Alanine, Serine, etc.

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1. Essential Amino Acids

These cannot be made by the body and must be obtained from the diet.

Amino Acid	3-Letter Code	Source / Notes
Lysine	Lys	Meat, eggs, beans
Methionine	Met	Fish, eggs, nuts
Leucine	Leu	Meat, soy, dairy
Isoleucine	Ile	Meat, eggs, cheese
Valine	Val	Meat, dairy, legumes
Threonine	Thr	Eggs, meat, milk
Phenylalanine	Phe	Dairy, soy, meat
Tryptophan	Trp	Milk, chocolate, turkey
Histidine	His	Meat, poultry, fish

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2. Non-Essential Amino Acids

These can be synthesized by the body.

Amino Acid	3-Letter Code	Source / Notes
Glycine	Gly	Gelatin, collagen
Alanine	Ala	Meat, poultry
Serine	Ser	Soy, dairy, eggs
Cysteine	Cys	Poultry, eggs, garlic
Tyrosine	Tyr	Meat, eggs, dairy
Aspartic acid	Asp	Meat, eggs
Glutamic acid	Glu	Meat, cheese, soy
Proline	Pro	Collagen-rich foods
Asparagine	Asn	Dairy, asparagus
Glutamine	Gln	Meat, eggs

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3. Special Amino Acids

• Selenocysteine (Sec, U) – sometimes called the 21st amino acid; contains selenium.

• Pyrrolysine (Pyl, O) – found in some microorganisms.

4. Key Points

• Amino acids are linked together by peptide bonds to form proteins.

• Each amino acid has:

  • Amino group (–NH₂)

  • Carboxyl group (–COOH)

  • Hydrogen atom (H)

  • Side chain (R group)

• Essential amino acids are important in diet; non-essential can be made in the body.

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Tests for Proteins

1. Biuret Test

   • Add Biuret reagent (NaOH + CuSO₄)

   • Purple color indicates presence of protein

2. Xanthoproteic Test

   • Add concentrated nitric acid

   • Yellow color indicates aromatic amino acids (like tyrosine, tryptophan)

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Summary

• Proteins are polymers of amino acids linked by peptide bonds.

• They have primary, secondary, tertiary, and quaternary structures.

• Proteins serve structural, functional, enzymatic, and regulatory roles.

• Proper protein intake is essential for growth, repair, immunity, and metabolism.

• Keratin = fibrous, structural protein

• Hemoglobin = globular, functional, conjugated protein

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✅ Tips for Students:

• Always link structure → function when answering protein questions.

• Remember fibrous = structural, globular = functional.

• Include examples and chemical/structural details for high marks.

Objective Questions

1. Proteins are polymers of:
A. Sugars
B. Amino acids
C. Fatty acids
D. Nucleotides

2. The bond linking amino acids in a protein is called:
A. Glycosidic bond
B. Peptide bond
C. Hydrogen bond
D. Ionic bond

3. Which of the following is an essential amino acid?
A. Glycine
B. Alanine
C. Lysine
D. Serine

4. Which protein is primarily structural in hair and nails?
A. Hemoglobin
B. Albumin
C. Keratin
D. Globulin

5. A protein that contains a non-protein prosthetic group is called:
A. Simple protein
B. Conjugated protein
C. Fibrous protein
D. Globular protein

6. Which of the following is non-essential?
A. Methionine
B. Leucine
C. Glycine
D. Isoleucine

7. The secondary structure of proteins can be:
A. Alpha-helix or beta-pleated sheet
B. Single chain of nucleotides
C. Fatty acid tail
D. Glycosidic linkage

8. Which test is used to detect proteins?
A. Benedict’s test
B. Biuret test
C. Iodine test
D. Molisch test

9. Hemoglobin is an example of:
A. Fibrous protein
B. Conjugated protein
C. Simple protein
D. Enzyme

10. Which element is found in all amino acids but not in carbohydrates or fats?
A. Carbon
B. Nitrogen
C. Hydrogen
D. Oxygen

Theory Questions

Short Answer

1. Define proteins.

2. Define amino acids.

3. Distinguish between essential and non-essential amino acids.

4. Name four functions of proteins in living organisms.

5. Define peptide bond and explain how it is formed.

Structured Questions

6. Draw and label the general structure of an amino acid.

7. Explain the four levels of protein structure (primary, secondary, tertiary, quaternary).

8. List three examples of simple proteins and three examples of conjugated proteins.

9. State the Biuret test procedure for detecting proteins.

10. Explain the difference between fibrous and globular proteins.

Higher-Level / Application Questions

11. Explain why amino acids are called the building blocks of proteins.

12. Hemoglobin contains iron as part of its structure. What type of protein is hemoglobin and why?

13. A student carries out a Biuret test and gets a violet color. What does this indicate?

14. Explain the importance of essential amino acids in the human diet.

15. Compare the structural features of keratin and hemoglobin in terms of protein classification.

STOICHIOMETRY

 Stoichiometry is the branch of chemistry that deals with the quantitative relationship between reactants and products in a chemical reaction.

That is, the relationship between the number of moles of reactants and products in a chemical reaction 

In simple terms, stoichiometry helps us to calculate:

• How much reactant is needed

• How much product will be formed during a chemical reaction

Stoichiometry is based on the law of conservation of mass, which states that matter can neither be created nor destroyed in a chemical reaction.

Importance of Stoichiometry

Stoichiometry is used to:

• Calculate masses of reactants and products

• Determine the amount of substances in reactions

• Predict product yield

• Find limiting and excess reactants

• Design industrial chemical processes

Basic Terms in Stoichiometry

1. Mole

A mole is the amount of substance that contains
6.02 × 10²³ particles (Avogadro’s number).

The mole is also the unit of measurement in chemistry.

2. Molar Mass

The molar mass is the mass of one mole of a substance in grams (g/mol).

Example:
Molar mass of H₂O = 2(1) + 16 = 18 g/mol

3. Chemical Equation

A chemical equation shows the relationship between reactants and products.

Example:

         2H₂         +        O₂ →       2H₂O

         2mols               1mol         2mol

or        2g                 (2x16)g      (2x18)g

or         2g                    32g            22.4dm3

This means:
2 moles of hydrogen react with 1 mole of oxygen to form 2 moles of water.

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Types of Stoichiometric Calculations

1. Mole-to-Mole Calculations

This involves using the ratio of moles in a balanced equation. That is, mole-mole relationship 

Example 1:

How many moles of oxygen are needed to react with 4 moles of hydrogen?

2H₂ + O₂ → 2H₂O

From equation:

2 moles H₂ react with 1 mole O₂
So, 4 moles H₂ will need:

1/2 x 4 = 2 moles of O₂

2. Mass-to-Mole Calculations

Example 2:

What is the number of moles in 44 g of CO₂?

Molar mass of CO₂ = 12 + 2(16) = 44 g/mol

Moles = Mass       =    

           Molar mass   

  44 = 1mole
  44 

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3. Mass-to-Mass Calculations

Example 3:

What mass of CO₂ is produced when 10 g of CaCO₃ decomposes?

Equation:

CaCO₃ →CaO + CO₂

Step 1: Molar masses
CaCO₃ = 100 g/mol
CO₂ = 44 g/mol

From equation:
100 g CaCO₃ → 44 g CO₂

So,
10 g CaCO₃ →?

         10    = 4.4 g of CO₂
        100

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4. Volume-to-Volume (Gaseous Reactions)

At the same temperature and pressure, equal volumes of gases contain equal number of molecules.

Example 4:

What volume of oxygen is needed to react with 40 cm³ of hydrogen?

2H₂ + O₂ →2H₂O

2 volumes H₂ react with 1 volume O₂
So,
40 cm³ H₂ will need:

1/2x 40 = 20 cm³ of  O₂

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5. Limiting Reactant Calculations

The limiting reactant is the reactant that is completely used up first and stops the reaction.

Example 5:

If 2 g of hydrogen reacts with 16 g of oxygen, which is limiting?

2H₂ + O₂ →2H₂O

Moles:
H₂ = 2 ÷ 2 = 1 mole
O₂ = 16 ÷ 32 = 0.5 mole

Required ratio:
2H₂ : 1O₂
Actual ratio:
1H₂ : 0.5 O₂ → correct ratio

So, no reactant is in excess — both are completely used up.

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Percentage Yield

Not all reactions give maximum product.

Formula:

Percentage Yield =   Actual Yield        x 100
                              Theoretical Yield

Example 6:

If theoretical yield = 10 g and actual yield = 8 g

             8     = 100 = 80%
            10 

Summary

Stoichiometry helps chemists:

• Predict quantities in reactions

• Save materials

• Improve industrial efficiency

• Avoid wastage

OBJECTIVE QUESTIONS (WAEC/NECO)

1. Stoichiometry deals with the
A. speed of reactions
B. colour of substances
C. quantitative relationship between reactants and products
D. energy changes in reactions

2. The number of particles in one mole of a substance is
A. 3.01 × 10²³
B. 6.02 × 10²³
C. 1.00 × 10²³
D. 12.00 × 10²³

3. The molar mass of CO₂ is
A. 12 g/mol
B. 16 g/mol
C. 28 g/mol
D. 44 g/mol

4. How many moles are present in 18 g of water?
A. 0.5
B. 1
C. 2
D. 18

5. In the equation
2H₂ + O₂ → 2H₂O
the mole ratio of H₂ to O₂ is
A. 1:1
B. 1:2
C. 2:1
D. 2:2

6. What mass of NaCl contains 1 mole of NaCl?
A. 23 g
B. 35.5 g
C. 58.5 g
D. 46 g

7. At the same temperature and pressure, equal volumes of gases contain
A. equal masses
B. equal densities
C. equal number of molecules
D. equal pressures

8. Which of the following is the limiting reactant?
A. The reactant in excess
B. The reactant completely used up
C. The product formed
D. The catalyst

9. The formula for calculating percentage yield is
A. Actual × Theoretical
B. Actual ÷ Theoretical × 100
C. Theoretical ÷ Actual × 100
D. Actual − Theoretical

10. How many moles are in 44 g of CO₂?
A. 0.5
B. 1
C. 2
D. 44

11. What volume of oxygen is required to react with 40 cm³ of hydrogen?
(2H₂ + O₂ → 2H₂O)
A. 10 cm³
B. 20 cm³
C. 30 cm³
D. 40 cm³

12. Which law is the basis of stoichiometry?
A. Law of definite proportion
B. Law of multiple proportions
C. Law of conservation of mass
D. Law of gaseous volumes

13. The molar mass of CaCO₃ is
A. 40
B. 56
C. 84
D. 100

14. How many moles are present in 32 g of O₂?
A. 0.5
B. 1
C. 2
D. 16

15. Which of the following is NOT used in stoichiometric calculations?
A. Balanced equation
B. Molar mass
C. Temperature only
D. Mole ratio

THEORY QUESTIONS (WAEC/NECO)

Short Answer Questions

1. Define stoichiometry.
2. What is a mole?
3. State Avogadro’s number.
4. Define molar mass.
5. What is a limiting reactant?

6. Calculate the number of moles in 22 g of CO₂.

7. What mass of CO₂ is produced when 50 g of CaCO₃ decomposes? according to the equation
CaCO₃ → CaO + CO₂

8. How many moles of oxygen are needed to react completely with 6 moles of hydrogen? given the equation below
2H₂ + O₂ → 2H₂O

9. Explain stoichiometry and state three of its applications.

10. Describe how to calculate the mass of a product formed from a given mass of reactant using a balanced chemical equation.

11. In a reaction, 10 g of calcium carbonate was heated and produced 3.5 g of carbon dioxide.
(a) Calculate the theoretical yield
(b) Calculate the percentage yield
CaCO₃ → CaO + CO₂

12. Explain the term “limiting reactant” and show with an example.

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REDOX REACTIONS (Oxidation–Reduction Reactions) at a glance

 

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Meaning of Redox Reaction

A redox reaction is a chemical reaction in which oxidation and reduction occur simultaneously.

The word redox comes from:

• RED → Reduction

• OX → Oxidation

In every redox reaction:

• One substance is oxidized

• Another substance is reduced

They always happen together — you cannot have oxidation without reduction.

Definitions of Oxidation

Oxidation can be defined in several ways depending on the concept used:

1. Oxidation in terms of Oxygen

Oxidation is the addition of oxygen to a substance.

Example:

2Mg + O₂ →2MgO

Magnesium is oxidized because it gains oxygen.

2. Oxidation in terms of Hydrogen

Oxidation is the removal of hydrogen from a substance.

Example:

H₂S → S + H₂

Hydrogen sulphide is oxidized because hydrogen is removed.

3. Oxidation in terms of Electrons

Oxidation is the loss of electrons by a substance.

Example:

Zn →Zn²⁺ + 2e^-

Zinc is oxidized because it loses electrons.

4. Oxidation in terms of Oxidation Number

Oxidation is an increase in oxidation number.

Example:
Fe²⁺ → Fe³⁺
Iron is oxidized because its oxidation number increases from +2 to +3.

5. Oxidation in terms of Electrochemistry

Oxidation occurs at the anode in an electrochemical cell.

Definitions of Reduction

Reduction is the opposite of oxidation and can also be defined in different ways:

1. Reduction in terms of Oxygen

Reduction is the removal of oxygen from a substance.

Example:

CuO + H₂ →Cu + H₂O

Copper (II) oxide is reduced because oxygen is removed.

2. Reduction in terms of Hydrogen

Reduction is the addition of hydrogen to a substance.

Example:

N₂ + 3H₂ → 2NH₃

Nitrogen is reduced because it gains hydrogen.

3. Reduction in term of Electrons

Reduction is the gain of electrons by a substance.

Example:

Cu²⁺ + 2e^- →Cu

Copper ions are reduced because they gain electrons.

4. Reduction in terms of Oxidation Number

Reduction is a decrease in oxidation number.

Example:
Cl₂⁰ → Cl⁻^-1
 0           -1

Chlorine is reduced because its oxidation number decreases.

5. Reduction in terms of Electrochemistry

Reduction occurs at the cathode in an electrochemical cell.

Oxidizing and Reducing Agents

Oxidizing Agent

An oxidizing agent is a substance that:

• Causes oxidation of another substance

• Itself gets reduced

Example:
KMnO₄, O₂, Cl₂

Reducing Agent

A reducing agent is a substance that:

• Causes reduction of another substance

• Itself gets oxidized

Example:
Zn, H₂, CO

Example of a Redox Reaction

Zn + CuSO₄ → ZnSO₄ + Cu

• Zinc is oxidized (loses electrons)

• Copper (II) ions are reduced (gain electrons)

• Zinc is the reducing agent

• Copper (II) sulphate is the oxidizing agent

Importance and Examples of Redox Reactions

Redox reactions are very important in daily life and industry, such as:

• Respiration in living cells

• Rusting of iron

• Burning of fuels

• Electrolysis

• Production of metals

• Batteries and cells

• Photosynthesis

Summary Table

Oxidation	Reduction
Gain of oxygen	Loss of oxygen
Loss of hydrogen	Gain of hydrogen
Loss of electrons	Gain of electrons
Increase in oxidation number	Decrease in oxidation number
Occurs at anode	Occurs at cathode

Test for oxidizing agents

To common test or reactions that are used to test for an oxidizing agent involves the action on iron (II) chloride and hydrogen sulphide.

a)       Reaction with FeCl₂

          When an oxidizing agent is added to green iron (II chloride; the iron (II) ions become oxidized to yellow or brown Fe³⁺.

          Fe²⁺ →     Fe³⁺ + e^–

          green         yellow/brown

b)      Reaction with hydrogen sulphide

          When hydrogen sulphide is bubbled through a solution of an oxidizing agent, the sulphide ions S^2– becomes oxidized to elemental sulphur; and this is seen or observed as yellow deposits sulphur,                    i.e. S^2– → S(s) + 2e^–.

Test for reducing agents

Two commonest reagents that are used to test for a reducing agent are

1        Acidified potassium tetraoxomanganate(VI) (KMnO₄) and acidified potassium heptaoxodichromate(I) (K₂Cr₂O₇).

a)   Action of potassium hyptaoxodichromate (VI) (K₂Cr₂O₇)

  When acidified potassium heptaoxodichromate (VI) (K₂Cr₂O₇) is added to a sample of a reducing agent, its colour changes from orange to green, due to the reduction of the dichromate (VI) ion  (Cr⁶⁺) (orange) to chromium (III) (Cr³⁺) ion green

     Cr⁶⁺  +  3e^– → Cr³⁺
  Orange                green

 

b) Test using acidified potassium tetraoxomangane(VI) (KMnO₄)

  When acidified potassium tetraoxomanganate (VII) to a sample of reducing agent, the purple colour changes to colourless: due to the reduction of the manganate ion from (+7) which is purple to (+2) which is colourless and a more stable oxidation state.

          MnO₄^- + 8H⁺ + 5e^– →Mn²⁺ + 4H₂O
          purple                       colourless         

    

     Mn⁷⁺ + 5e^– → Mn²
  purple              colourless

                   This reaction is reversible as the purple colour is restored when an oxidizing agent is reintroduced into the mixture.

                  Mn²⁺ + 5e →Mn⁷⁺   
               colourless        purple

OBJECTIVE QUESTIONS 

1. A redox reaction is one in which

A. acids react with bases
B. electrons are transferred
C. salts are formed
D. heat is evolved

2. Oxidation is the process involving
A. gain of electrons
B. loss of electrons
C. gain of neutrons
D. loss of protons

3. Reduction is best defined as
A. loss of hydrogen
B. gain of oxygen
C. gain of electrons
D. loss of oxygen and hydrogen

4. In the reaction
Zn + Cu²⁺ → Zn²⁺ + Cu
the species oxidized is
A. Cu²⁺
B. Zn²⁺
C. Zn
D. Cu

5. Which of the following is a reducing agent?
A. A substance that gains electrons
B. A substance that is reduced
C. A substance that loses electrons
D. A substance that gains oxygen

6. Which of the following statements is correct?
A. Oxidation and reduction occur separately
B. Reduction involves loss of electrons
C. Oxidation involves gain of electrons
D. Oxidation and reduction occur simultaneously

7. In terms of oxygen, oxidation is defined as
A. removal of oxygen
B. addition of oxygen
C. removal of hydrogen
D. addition of hydrogen

8. In the reaction
2Mg + O₂ → 2MgO
Magnesium is
A. reduced
B. oxidized
C. neutralized
D. displaced

9. Which of the following is an oxidizing agent?
A. A substance that is oxidized
B. A substance that loses electrons
C. A substance that gains electrons
D. A substance that donates protons

10. The oxidation number of chlorine in Cl₂ is
A. +1
B. −1
C. 0
D. +2

11. Reduction involves
A. increase in oxidation number
B. decrease in oxidation number
C. no change in oxidation number
D. increase in atomic mass

12. Which reaction is a redox reaction?
A. NaOH + HCl → NaCl + H₂O
B. AgNO₃ + NaCl → AgCl + NaNO₃
C. Zn + H₂SO₄ → ZnSO₄ + H₂
D. CaCO₃ → CaO + CO₂

13. The electrode at which oxidation occurs is called the
A. cathode
B. anode
C. electrolyte
D. salt bridge

14. Which of the following processes is an example of oxidation?
A. Rusting of iron
B. Freezing of water
C. Melting of ice
D. Dissolving sugar

15. A substance that causes reduction and is itself oxidized is called
A. oxidizing agent
B. electrolyte
C. reducing agent
D. catalyst

16. In the reaction
Fe²⁺ → Fe³⁺ + e⁻
iron is
A. reduced
B. oxidized
C. neutralized
D. displaced

17. Which of the following is NOT a redox reaction?
A. Combustion
B. Respiration
C. Electrolysis
D. Neutralization

18. In a redox reaction, electrons are transferred from
A. oxidizing agent to reducing agent
B. reducing agent to oxidizing agent
C. acid to base
D. salt to water

19. Which of the following changes represents reduction?
A. Cu → Cu²⁺
B. Fe²⁺ → Fe³⁺
C. Cl₂ → 2Cl⁻
D. Na → Na⁺

20. The main feature of a redox reaction is
A. formation of precipitate
B. evolution of gas
C. transfer of electrons
D. formation of salt

THEORY QUESTIONS 

Short Answer Questions

1. Define oxidation and reduction using the electron concept.

2. What is a redox reaction?

3. State two ways by which oxidation can be defined.

4. State two ways by which reduction can be defined.

5. What is meant by an oxidizing agent?

6. What is meant by a reducing agent?

7. Why must oxidation and reduction always occur together?

8. Identify the oxidizing and reducing agents in the reaction:
Zn + CuSO₄ → ZnSO₄ + Cu

Structured / Calculation-Based Questions

9. For the reaction:
2Fe³⁺ + Sn²⁺ → 2Fe²⁺ + Sn⁴⁺
(a) State the species oxidized
(b) State the species reduced
(c) Identify the oxidizing agent
(d) Identify the reducing agent

10. In the reaction:
Mg + 2HCl → MgCl₂ + H₂
(a) Which substance is oxidized?
(b) Which substance is reduced?
(c) State the role of magnesium.

Essay Questions

11. Define oxidation and reduction. Explain each using:
(a) oxygen
(b) hydrogen
(c) electrons
(d) oxidation numbers.

12. With examples, explain what is meant by oxidizing and reducing agents.

13. Describe an experiment to show that oxidation and reduction occur simultaneously.

14. Explain the importance of redox reactions in everyday life and industry.

15. Using suitable examples, distinguish between oxidation and reduction.

Practical-Oriented Questions

16. Explain how rusting of iron is a redox reaction.

17. Describe the redox processes that occur during electrolysis of molten sodium chloride.

18. Explain how respiration in living cells is a redox process.

19. Balance the following redox equation using the ion-electron method:

MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺ (in acidic medium)

20. State three industrial applications of redox reactions.

ALDEHYDES (R–CHO)

Definition

Aldehydes are organic compounds that contain the formyl functional group (–CHO), in which a carbon atom is double-bonded to oxygen and single-bonded to hydrogen.

General formula:
R–CHO

where R is an alkyl or aryl group.

Nomenclature (Naming of Aldehydes)

IUPAC Naming Rules

1. Identify the longest carbon chain containing the –CHO group.

2. The carbon of the –CHO group is always numbered carbon 1.

3. Replace the –e ending of the parent alkane with –al.

Examples:

Molecular Formula	IUPAC Name	Common Name
HCHO	Methanal	Formaldehyde
CH₃CHO	Ethanal	Acetaldehyde
C₂H₅CHO	Propanal	Propionaldehyde
C₆H₅CHO	Benzaldehyde	Benzaldehyde

Occurrence of Aldehydes

I. Found in natural substances such as vanilla (vanillin) and cinnamon (cinnamaldehyde).

• Present in essential oils, perfumes, and flavorings.

• Produced during oxidation of alcohols.

Preparation of Aldehydes

1. Oxidation of Primary Alcohols

Aldehydes are prepared by controlled oxidation of primary alcohols.

R–CH₂OH + [O] → R–CHO + H₂O

Oxidizing agents:

• Acidified potassium dichromate (VI) (K₂Cr₂O₇/H₂SO₄)

• Acidified potassium permanganate (KMnO₄)

NOTE:  Excess oxidation converts aldehydes to carboxylic acids.

2. Dehydrogenation of Alcohols

Passing alcohol vapour over heated copper (300°C).

R–CH₂OH ---{Cu}--> R–CHO + H₂

3. Hydrolysis of Geminal Dihalides

R–CHX₂ + ---{H₂O}--->R–CHO

Physical Properties of Aldehydes

1. Lower aldehydes are colorless liquids with sharp smells.

2. They are polar compounds.

3. Lower members are soluble in water due to hydrogen bonding.

4. Boiling points are higher than alkanes but lower than alcohols.

5. Formaldehyde exists as a gas at room temperature.

Chemical Properties of Aldehydes

1. Oxidation (Reducing Nature)

Aldehydes are easily oxidized to carboxylic acids.

R–CHO + [O] → R–COOH
 

a. Tollens’ Test (Silver Mirror Test)

• Aldehydes reduce ammoniacal silver nitrate to silver.

• Produces a silver mirror.

✔️ Distinguishes aldehydes from ketones.

b. Fehling’s Solution Test

• Aldehydes reduce Fehling’s solution to a brick-red precipitate (Cu₂O).

2. Reduction

Aldehydes are reduced to primary alcohols.

R–CHO + [H] → R–CH₂OH

Reducing agents:

• NaBH₄

• LiAlH₄

• Hydrogen/Nickel catalyst

3. Addition Reactions

Due to the polar C=O bond, aldehydes undergo nucleophilic addition.

Example:
R–CHO + HCN → R–CH(OH)CN

4. Reaction with Sodium Hydrogen Sulphite

Forms crystalline addition compounds (used for purification).

Uses of Aldehydes

1. Formaldehyde (Methanal)

• Used as formalin for preserving specimens.

• Manufacture of plastics, resins, and disinfectants.

• Used in textile and leather industries.

2. Ethanal (Acetaldehyde)

• Manufacture of acetic acid.

• Used in perfumes and dyes.

3. Benzaldehyde

• Used in flavorings and perfumes (almond smell).

Differences Between Aldehydes and Ketones

Aldehydes	Ketones
Have –CHO group	Have –CO– group
Easily oxidized	Not easily oxidized
Give Tollens’ test	Do not
Carbonyl carbon attached to H	No H attached

Important tip

• Aldehydes are reducing agents.

• Prepared from primary alcohols only.

• Distinguished using Tollens’ and Fehling’s tests.

• General formula: R–CHO

 

OBJECTIVE QUESTIONS

1. The functional group present in aldehydes is
   A. –COOH
   B. –CO–
   C. –CHO
   D. –OH

2. The general formula of aliphatic aldehydes is
   A. R–CO–R
   B. R–COOH
   C. R–CHO
   D. R–OH

3. The IUPAC name of CH₃CHO is
   A. Methanal
   B. Ethanal
   C. Propanal
   D. Acetone

4. Which of the following alcohols on oxidation produces an aldehyde?
   A. Tertiary alcohol
   B. Secondary alcohol
   C. Primary alcohol
   D. Phenol

5. Which reagent gives a silver mirror with aldehydes?
   A. Fehling’s solution
   B. Benedict’s solution
   C. Tollens’ reagent
   D. Acidified KMnO₄

6. Aldehydes can be oxidized to form
   A. Ketones
   B. Alcohols
   C. Carboxylic acids
   D. Esters

7. Which aldehyde exists as a gas at room temperature?
   A. Ethanal
   B. Methanal
   C. Propanal
   D. Benzaldehyde

8. Aldehydes act as reducing agents because they
   A. accept hydrogen easily
   B. are easily oxidized
   C. are acidic
   D. have low boiling points

9. Which of the following does NOT react with aldehydes?
   A. Tollens’ reagent
   B. Fehling’s solution
   C. Sodium hydrogen sulphite
   D. Schiff’s reagent

10. Reduction of an aldehyde produces a
    A. secondary alcohol
    B. tertiary alcohol
    C. primary alcohol
    D. carboxylic acid

11. Which of the following aldehydes has an aromatic ring?
    A. Methanal
    B. Ethanal
    C. Propanal
    D. Benzaldehyde

12. The oxidation of aldehydes to acids involves
    A. loss of hydrogen
    B. gain of hydrogen
    C. loss of oxygen
    D. no chemical change

13. Aldehydes undergo nucleophilic addition mainly due to the
    A. non-polar nature of C=O
    B. polarity of the carbonyl group
    C. presence of –OH group
    D. acidic nature of hydrogen

14. Which of the following can be used to purify aldehydes?
    A. Dilute acids
    B. Sodium hydroxide
    C. Sodium hydrogen sulphite
    D. Calcium chloride

15. The common name of methanal is
    A. Acetaldehyde
    B. Formaldehyde
    C. Propionaldehyde
    D. Benzaldehyde

KETONES at a glance

 

Definition

Ketones are organic compounds that contain the carbonyl group (C=O) bonded to two alkyl or aryl groups.
The carbonyl carbon is not attached to a hydrogen atom.

General Formula

R–CO–R′

Where R and R′ are alkyl or aryl groups.

📌 Important point: Ketones must have at least three carbon atoms.

Functional Group

• Carbonyl group (C=O)

• Located within the carbon chain, not at the end.

Nomenclature (IUPAC Naming)

Ketones are named by:

• Replacing the ending –e in the corresponding alkane name with –one

• Indicating the position of the carbonyl group if necessary

Examples:

• Propanone (acetone) → CH₃COCH₃

• Butanone → CH₃COCH₂CH₃

• Pentan-2-one

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Preparation of Ketones

1. Oxidation of Secondary Alcohols
   Secondary alcohol + [O] → Ketone

Example:

        Propan-2-ol → Propanone

2. Dry Distillation of Calcium Salts of Carboxylic Acids

3. Hydration of Alkynes (acid-catalysed)

Physical Properties

• Colourless liquids (lower members)

• Pleasant or sharp smell

• Boiling points higher than alkanes but lower than alcohols

• Soluble in water (lower members like propanone)

• Polar due to the C=O group

Chemical Properties

1. Resistance to Oxidation

• Ketones are not easily oxidized

• Do not react with:

  • Tollens’ reagent

  • Fehling’s solution

📌 Important distinction from aldehydes

2. Reduction

• Reduced to secondary alcohols

3. Addition Reactions

• React with hydrogen cyanide (HCN)

• React with sodium hydrogen sulphite

Laboratory Tests for Ketones

1. 2,4-Dinitrophenylhydrazine (2,4-DNPH) Test

• Yellow or orange precipitate

• Confirms presence of carbonyl group

2. Iodoform Test

• Positive for methyl ketones

• Yellow precipitate of iodoform (CHI₃)

Example:

• Propanone gives a positive iodoform test

Uses of Ketones

• Solvents (e.g. acetone)

• Nail polish remover

• Paint and varnish industries

• Pharmaceutical production

• Plastic and fibre manufacture

Differences Between Aldehydes and Ketones

Aldehydes	Ketones
–CHO group	–CO– group
Easily oxidized	Not easily oxidized
Positive Tollens’ test	Negative Tollens’ test
At end of chain	In the middle of chain

Examples of Ketones

• Propanone (acetone)

• Butanone

• Cyclohexanone

WAEC / NECO EXAM TIPS

• Ketones do not reduce Tollens’ or Fehling’s solution

• Oxidation of secondary alcohols gives ketones

• Methyl ketones give a positive iodoform test

OBJECTIVE QUESTIONS

1. Ketones are organic compounds that contain the functional group
   A. –OH
   B. –COOH
   C. –CHO
   D. –CO–

2. The general formula of ketones is
   A. R–CHO
   B. R–COOH
   C. R–CO–R′
   D. R–OH

3. Which of the following is the simplest ketone?
   A. Methanone
   B. Ethanal
   C. Propanone
   D. Ethanoic acid

4. Ketones differ from aldehydes because ketones
   A. are easily oxidized
   B. reduce Fehling’s solution
   C. have the carbonyl group at the end of the chain
   D. have the carbonyl group within the chain

5. Which reagent can be used to distinguish between an aldehyde and a ketone?
   A. Bromine water
   B. Tollens’ reagent
   C. Sodium hydroxide
   D. Dilute acid

6. Ketones do NOT react with
   A. Hydrogen cyanide
   B. Sodium hydrogen sulphite
   C. 2,4-dinitrophenylhydrazine
   D. Tollens’ reagent

7. The oxidation of secondary alcohols produces
   A. aldehydes
   B. ketones
   C. alkanes
   D. carboxylic acids

8. Which of the following ketones gives a positive iodoform test?
   A. Butanone
   B. Pentanone
   C. Cyclohexanone
   D. Benzophenone

9. The yellow precipitate formed in the iodoform test is
   A. CHCl₃
   B. CHBr₃
   C. CHI₃
   D. C₂H₅I

10. Ketones can be reduced to form
    A. primary alcohols
    B. secondary alcohols
    C. tertiary alcohols
    D. carboxylic acids

11. Which of the following reagents gives an orange precipitate with ketones?
    A. Fehling’s solution
    B. Tollens’ reagent
    C. 2,4-dinitrophenylhydrazine
    D. Benedict’s solution

12. The boiling points of ketones are generally
    A. lower than alkanes
    B. equal to alcohols
    C. higher than alkanes but lower than alcohols
    D. higher than carboxylic acids

13. Propanone is commonly known as
    A. formaldehyde
    B. acetaldehyde
    C. acetone
    D. acetic acid

14. Which of the following is NOT a use of ketones?
    A. Solvent
    B. Nail polish remover
    C. Paint manufacture
    D. Fuel for engines

15. Ketones are best described as
    A. non-polar compounds
    B. ionic compounds
    C. polar compounds
    D. basic compounds

THEORY QUESTIONS

   

NUCLEAR CHEMISTRY summary note for student

 

Introduction

Nuclear chemistry is the branch of chemistry that deals with changes in the nucleus of an atom. These changes involve the emission of particles or radiation and are known as nuclear reactions. Unlike ordinary chemical reactions, nuclear reactions involve the nucleus and may result in the formation of new elements.

THĺE NUCLEUS AND NUCLEAR STABILITY

The nucleus of an atom contains:

• Protons (positively charged)

• Neutrons (neutral)

The stability of a nucleus depends on the neutron–proton ratio. Unstable nuclei undergo radioactive decay to become more stable.

RADIOACTIVITY

Radioactivity is the spontaneous disintegration of unstable atomic nuclei with the emission of radiation.

Types of Radioactive Radiation

1.  Characteristics of Alpha (α) particles

i.  they are positively charged particles that resembles the helium nucleus

ii.  they are deflected towards the negative plate 

iii.  They have Low penetrating power (Can be stopped by paper or skin)

iv.  They have high ionizing power

 Example:

         ²³⁸₉₂U → ²³⁴₉₀Th + ⁴₂He

2. Beta (β) particles

i. Fast-moving electrons

ii. Negatively charged 

iii. Moderate penetrating power (Stopped by thin aluminum sheet)

iv. they are deflected towards the positive plate in an electrostatic field.

      Example:

     ¹⁴₆C → ¹⁴₇N + β⁻

3. Gamma (γ) rays 

i. High-energy electromagnetic radiation 

ii. They have No mass and no charge. 

iii. They have Very high penetrating power ( Stopped by thick lead or concrete)

iv. They are not affected by an electrostatic field

NUCLEAR EQUATIONS

In nuclear reactions:

• Mass number is conserved

• Atomic number is conserved

HALF-LIFE

Half-life is the time taken for half the number of radioactive atoms in a substance to decay.

Example:
If a substance has a half-life of 10 years:

• After 10 years → ½ remains

• After 20 years → ¼ remains

TYPES OF NUCLEAR REACTIONS

1. Nuclear Fission

Nuclear fission is the splitting of a heavy nucleus into two lighter nuclei with the release of a large amount of energy and radiation

Example:

 ²³⁵₉₂U + ¹₀n → ¹⁴¹₅₆Ba + ⁹²₃₆Kr + 3¹₀n + energy

Uses of fission:

i.  Nuclear power plants for power generation

ii. Atomic bombs

2. Nuclear Fusion

Nuclear fusion is the combination of two light nuclei to form a heavier nucleus with the release of energy and radiation.

Example:

²₁H + ³₁H → ⁴₂He + energy

Uses of fusion:

It is a source of energy in the sun and stars

It used for making Hydrogen bomb

DIFFERENCE BETWEEN FISSION AND FUSION

Fission	Fusion
Splitting of heavy nucleus	Combination of light nuclei
Produces radioactive waste	Produces little waste
Used in nuclear reactors	Occurs in the sun
Lower temperature required	Extremely high temperature required

ARTIFICIAL TRANSMUTATION

Artificial transmutation is the conversion of one element into another by bombarding the nucleus with particles. Thats is, causing radioactivity to occur artificially

Example:

  ¹⁴₇N + ⁴₂He → ¹⁷₈O + ¹₁H

USES OF RADIOISOTOPES

In Medicin Cancer treatment (radiotherapy)

i. Tracers in diagnosis

ii. Sterilization of medical equipment

iii. Treatment of cancerous cells

In Industry

i. Detecting cracks in metals

ii. Thickness control in manufacturing Packaging materials 

iii. 

In Agriculture

i. Food preservation

ii. Pest control (radiations are used to destroy the reproductive cells of male insect)

In Archaeology

i. Carbon-14 dating to determine age of fossils

HAZARDS OF NUCLEAR RADIATION

• Causes cancer

• Damages living tissues

• Leads to genetic mutations

• Can cause radiation sickness

Safety Measures

• Use of lead materials for shielding

• Wearing protective clothing

• Proper disposal of radioactive waste

ADVANTAGES OF NUCLEAR ENERGY

• Produces large amount of energy

• Requires small amount of fuel

• No greenhouse gas emission during operation

DISADVANTAGES OF NUCLEAR ENERGY

• Radioactive waste disposal problem

• Risk of nuclear accidents

• High cost of setup and maintenance

SUMMARY (AT A GLANCE)

• Nuclear chemistry deals with changes in atomic nuclei

• Radioactivity involves alpha, beta, and gamma radiation

• Half-life measures rate of decay

• Nuclear reactions include fission and fusion

• Nuclear energy has many uses but also serious hazards

OBJECTIVE QUESTIONS

1. Nuclear chemistry mainly deals with changes in the
A. electron cloud
B. outer shell electrons
C. nucleus of an atom
D. valence electrons

2. Which of the following particles has the greatest penetrating power?
A. Alpha particles
B. Beta particles
C. Gamma rays
D. Protons

3. An alpha particle consists of
A. one proton
B. one electron
C. two protons and two neutrons
D. two electrons and two protons

4. Which radiation is deflected most by an electric field?
A. Alpha rays
B. Beta rays
C. Gamma rays
D. Neutron rays

5. Radioactivity is best described as
A. a chemical change
B. a physical change
C. spontaneous nuclear disintegration
D. a reversible reaction

6. The half-life of a radioactive substance is the time taken for
A. all atoms to decay
B. half of the atoms to decay
C. the activity to stop
D. the mass to double

7. If a radioactive substance has a half-life of 5 days, how long will it take for three-quarters of it to decay?
A. 5 days
B. 10 days
C. 15 days
D. 20 days

8. Which of the following is conserved in a nuclear reaction?
A. Number of electrons
B. Chemical properties
C. Atomic number and mass number
D. Physical state

9. The splitting of a heavy nucleus into lighter nuclei is known as
A. nuclear fusion
B. radioactive decay
C. nuclear fission
D. artificial transmutation

10. Nuclear fusion occurs mainly in
A. nuclear reactors
B. atomic bombs
C. the sun and stars
D. radioactive waste

11. One major difference between nuclear reactions and chemical reactions is that nuclear reactions
A. involve electrons
B. involve energy changes only
C. involve the nucleus
D. are reversible

12. Which of the following is used as a moderator in a nuclear reactor?
A. Graphite
B. Uranium
C. Plutonium
D. Lead

13. Carbon-14 is mainly used for
A. treating cancer
B. food preservation
C. determining the age of fossils
D. generating electricity

14. Which radiation is stopped by a sheet of paper?
A. Gamma rays
B. Beta particles
C. Alpha particles
D. Neutron rays

15. In nuclear fission, energy is released because
A. mass is conserved
B. mass is converted into energy
C. electrons are transferred
D. atoms are rearranged

16. Artificial transmutation involves
A. natural decay of elements
B. fusion of light nuclei
C. conversion of one element to another by bombardment
D. spontaneous disintegration

17. Which of the following is a hazard of nuclear radiation?
A. Increase in melting point
B. Formation of alloys
C. Genetic mutation
D. Improved conductivity

18. The SI unit of radioactivity is
A. joule
B. becquerel
C. watt
D. volt

19. Which statement about gamma rays is correct?
A. They are negatively charged
B. They have mass
C. They are electromagnetic waves
D. They are easily stopped by paper

20. Which of the following is an advantage of nuclear energy?
A. Produces smoke
B. Requires large fuel quantity
C. Produces large energy from small fuel
D. Produces no waste

---

FARADAY’S LAWS OF ELECTROLYSIS

 

Faraday’s First Law of Electrolysis: State that the mass (m) of an element discharged during electrolysis is directly proportional to the quantity of electricity (Q) passing through the electrolyte

Mathematically

M α Q

Q = It

M α It         removing the sign of proportionality we have 

M = ZIt   

Where Z is a constant known as the electrochemical equivalent of the substance.

M = Mass of substance in gram

Q = Quantity of electricity in coulombs

I = Current in ampere

t = Time in seconds

Verification of Faraday's First law of electrolysis

Faraday’s Second Law of Electrolysis: State that when the same quantity of electricity is passed through solutions of different electrolytes, the relative number of moles of the elements discharged at each electrode is inversely proportional to the charges on the ions of each of the element

According to Faraday the minimum quantity of electricity required to liberate one mole of a univalent ion during electrolysis is equal to 1 Faraday and 

1 Faraday = 96500 coulombs 

 

Mass              = Quantity of Electricity

Molar mass       Faraday

 

 

Verification Of Faraday’s Second Law 

Method

 

                               

  

1. Fill two beakers up to ⅔ of their volumes with 1M of copper (II) tetraoxosulphate (VI) solution and 1M solution of silver trioxonitrate (V) solution.

2. Weigh and place two clean plates of copper and silver electrodes in their respective solutions 

3. Connect a battery and complete the circuit as shown above, attach a variable resistor adjusted to maintain a steady current of 0.5A. Allow the current to   pass through the solution for 25 minutes.

4. The cathode is removed, washed with water dried and then reweighed to obtain the masses of copper and silver deposited

5. The ratio of the number of moles of copper and silver deposited is then calculated.

6. The process is repeated to obtain at least three more readings for accuracy.

 

Amounts n (Number of moles) = Mass of element deposited

                                                                  Its relative atomic mass.

Observation: On passing the same quantity of electricity through the solutions, the ratio of the number of moles of copper and silver deposited is 1:2.

This ratio is inversely proportional to the ratio of the charges on the ions, Cu²⁺  and  Ag⁺ , or the number of moles of electrons required to liberate 1 mole each of the ions.

Cu²⁺_(aq)  +  2e^- → Cu(s)             and         Ag⁺_(aq)  + e^-          →         Ag(s)

                    2 moles    63.5g                                   1 mole                     108g

Conclusion:  When the same quantity of electricity is passed through a solution of different electrolytes the relative number of moles of the elements deposited are inversely proportional to the charges on the ions of each of the elements respectively.

 

CALCULATION BASED ON THE FIRST LAW OF ELECTROLYSIS

1. In an electrolysis experiment, the ammeter records a steady current of 1A. The mass of copper deposited in 30mins is 0.66g. Calculate the error in the ammeter reading. [electrochemical equivalent of copper =0.00033gC-1]            

Solution

M = ZIt

M = 0.66g, Z = 0.00033gC-1, t = 30mins = 1800 seconds

I = M/Zt

I = 0.66/0.00033 x 1800

I = 0.66/0.594

I = 1.11A

The error in the ammeter reading is 1.11 – 1 = 0.11A

 

2. Calculated the time in minutes, required to plate a substance of total surface area 300cm², a layer of copper 0.6mm thick, if a constant current of 2A is maintained. Assuming the density of copper is 8.8g/cm3 and one coulomb liberates 0.00033g copper.

Solution

Given that area = 300cm² , thickness = 0.6mm = 0.06cm

Mass = 0.00033g, density = 8.8g/cm3

Density = mass/volume

Mass = density x volume

Mass = 8.8 x 300 x 0.06

Mass = 158.4g

From M = Zit

t = m/ZI

t = 158.4/2 x 0.00033

t = 240000secs

t= 4000mins

 

 

CALCULATION BASED ON THE SECOND LAW OF ELECTROLYSIS

 

6. A current of 4.5A is passed through a solution of gold salt for 1 hour 45 minutes. Calculate

(i) The mass of gold deposited  

(ii) The number of moles of gold deposited

(iii) If the same current is used, find the time taken for 5.5g of gold to be deposited (Au = 197, 1 Faraday = 96500c)

Solution

(i) Au+   +     e-→ Au

    197g       1F       197g

Mass              = Quantity of Electricity
Molar mass       Faraday

M        Q
Mm     F

 

Mass = ?

Molar mass = 197g

Quantity of electricity = I x t =

I = 4.5A

t = 1 hour 45minutes = 105 minutes = 105 x 60 = 6300 seconds

Quantity of electricity = I x t = 4.5 x 6300 = 28,350C

Faraday = 96500F

 

Mass         =    28,350
197                  96500

Mass    =   0.29378
197

Mass of Gold deposited = 57.88g

 

(ii) Number of mole = Mass
                                    Molar mass

 

Number of mole = 57.88
                                197

The number of mole of Gold deposited = 0.30mol

 

(iii) Mass              =    Quantity of Electricity
      Molar mass              Faraday

M        Q
Mm     F

5.5    =  Quantity of Electricity

197                  96500 

0.02792 = Quantity of Electricity

                            96500

Quantity of Electricity = 2694.28C. This is the quantity of electricity (Q) required for 5.5g of Au to be deposited.

Q = It

t = Q/I

t = 2694
        4.5

t = 598.7

The time taken is 9.98 minutes

 

OBJECTIVE QUESTION

Choose the correct option from A – D.

 1. Faraday’s first law of electrolysis states that the mass of a substance deposited is proportional to the

A. time taken
B. voltage applied
C. quantity of electricity passed
D. temperature

2. The quantity of electricity passed during electrolysis is equal to

A. current × resistance
B. voltage × time
C. current × time
D. resistance × time

3. The SI unit of quantity of electricity is

A. ampere
B. volt
C. coulomb
D. ohm

4. One coulomb is equal to

A. 1 A × 1 s
B. 1 V × 1 s
C. 1 Ω × 1 s
D. 1 J × 1 s

5. According to Faraday’s first law, if the current is doubled, the mass deposited will

A. halve
B. remain constant
C. double
D. become zero

6. Faraday’s second law relates mass deposited to

A. current used
B. time of electrolysis
C. quantity of electricity
D. equivalent weight of the substance

7. The equivalent weight of an element is its

A. atomic mass
B. molecular mass
C. atomic number
D. atomic mass ÷ valency

8. According to Faraday’s second law, equal quantities of electricity will deposit masses proportional to their

A. atomic numbers
B. densities
C. melting points
D. equivalent weights

9. Which of the following will deposit the highest mass for the same quantity of electricity?

A. Na
B. Mg
C. Al
D. Ag

10. A current of 2 A flows for 10 seconds. The quantity of electricity passed is

A. 5 C
B. 10 C
C. 20 C
D. 40 C

11. The unit of current is

A. coulomb
B. ampere
C. volt
D. ohm

12. If the time of electrolysis is tripled, the mass deposited will

A. remain the same
B. halve
C. double
D. triple

13. The formula connecting mass, charge and electrochemical equivalent is

A. m = It
B. m = ZIt
C. Q = It
D. Z = mIt

14. In electrolysis, the electrochemical equivalent (Z) is the

A. mass of substance per second
B. mass per ampere
C. mass deposited per coulomb
D. atomic mass

15. Which of the following obeys Faraday’s laws?

A. Diffusion
B. Neutralization
C. Combustion
D. Electrolysis

16. When a metal ion gains electrons during electrolysis, the process is called

A. oxidation
B. ionization
C. reduction
D. dissociation

17. Which electrode gains mass during electrolysis?

A. anode
B. cathode
C. electrolyte
D. voltmeter

18. The amount of substance deposited depends on all except

A. current
B. time
C. equivalent weight
D. temperature

19. The greater the valency of a metal, the

A. greater the mass deposited
B. smaller the equivalent weight
C. higher the electrochemical equivalent
D. lower the mass deposited

20. Faraday’s laws are used to determine

A. atomic structure
B. chemical bonding
C. amount of substances produced during electrolysis
D. rate of reaction

THEORY QUESTION

1. At what time must a current of 5Amp pass through a solution of zinc sulphate to deposited 1g of zinc. Electrochemical equivalent (e.c.e) = 0.0003387

2. In an electrolysis experiment, a cathode of mass 5g is found to weigh 5.01g, after a current of 5A flows for 50 seconds. What is the electrochemical equivalent for the deposited substance?

 

3. The electrochemical equivalent of silver is 0.0012g/c. if 0.36g of silver is to be deposited by electrolysis on a surface by passing a steady current for 5.0 minutes. Calculate the value of the current.

4. Calculate the current that must be passed into a solution of aluminium salt for 1hr.30minutes in order to deposited 1.5g of Aluminium (Al = 27)

5. 0.222g of a divalent metal is deposited when a current of 0.45A is passed through a solution of its salt for 25 minutes. Calculate the relative atomic mass of the metal. (1 Faraday = 96500 coulombs)

 

6. A given quantity of electricity was passed through three electrolytic cells connected in series containing solutions of Silver trioxonitrate (V), Copper (II) tetraoxosulphate (VI) and Sodium Chloride respectively. If 10.5g of Copper are deposited in the second electrolytic cell. Calculate

(a) The mass of Silver deposited in the first cell.

(b) The Volume of Chloride liberated in the third cell at 180C and 760mmHg pressure. (Ag=108, Cu=63.5, 1Faraday=96500C, molar volume of gases at s.t.p =22.4dm³.)              

                

7. Calculated the time in minutes, required to plate a substance of total surface area 300cm2, a layer of copper 0.6mm thick, if a constant current of 2A is maintained. Assuming the density of copper is 8.8g/cm³ and one coulomb liberates 0.00033g copper.