Atoms and Molecules — Class 9 Science Notes (NCERT Chapter 3)

Atoms and molecules are the fundamental building blocks of matter, and understanding them is essential for learning chemistry. In this chapter, we explore the basic concepts of how matter is made up of tiny particles, the laws of chemical combination, and how atoms join together to form molecules. These Atoms and Molecules Class 9 Notes will help you simplify important NCERT concepts, prepare for exams, and build a strong foundation for higher studies in science.

Whether you are revising before a test or studying for the first time, this guide explains key terms, laws, and examples in an easy-to-understand manner with clear definitions, solved questions, and diagrams.

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Table of Contents

• Laws of Chemical Combination
• Dalton’s Atomic Theory
• Atoms and Molecules
• Atomic Mass and Molecular Mass
• Mole Concept
• Writing Chemical Formulae
• NCERT Exercise Solutions (Important Questions)
  • Short Answer Questions
  • Long Answer Questions
  • Numericals
• FAQs on Atoms and Molecules (Class 9)
• Summary (Quick Revision Notes)

Introduction to Atoms and Molecules

Definition of Atoms and Molecules

• Atoms: Atoms are the smallest unit of an element that retains its chemical properties. They are the basic building blocks of matter, consisting of a nucleus (containing protons and neutrons) surrounded by electrons. For example, a single hydrogen atom has one proton and one electron. Atoms are so tiny that they cannot be seen with the naked eye and are measured in nanometers (1 nm = 10⁻⁹ meters). (Relevant to atoms definition class 9.)
• Molecules: A molecule is a group of two or more atoms chemically bonded together, forming a single unit with distinct properties. Molecules can be made of the same type of atoms (e.g., O₂, oxygen gas) or different types of atoms (e.g., H₂O, water). The bonds holding atoms together in a molecule are strong chemical bonds, like covalent or ionic bonds. (Relevant to molecules definition.)

Importance in Understanding Matter

• Fundamental Units of Matter: Atoms and molecules are the foundation of all matter in the universe, whether solid, liquid, or gas. Understanding their structure and behavior helps explain how matter is composed and how substances interact. (Relevant to matter composition.)
• Chemical Reactions: Atoms and molecules are involved in chemical reactions, where they rearrange to form new substances. For instance, when hydrogen and oxygen atoms combine, they form water molecules, a process critical to life and industry.
• Properties of Substances: The arrangement and type of atoms in a molecule determine a substance’s properties, such as its color, state (solid, liquid, gas), or reactivity. For example, the molecular structure of graphite and diamond (both made of carbon atoms) results in vastly different properties.
• Scientific Advancements: Knowledge of atoms and molecules drives advancements in fields like chemistry, physics, biology, and materials science. It enables the development of new materials, medicines, and technologies.

Daily Life Connection

• Oxygen We Breathe: The air we breathe contains oxygen molecules (O₂), where two oxygen atoms are bonded together. These molecules are essential for respiration, providing energy to our cells. Without oxygen molecules, life as we know it would not exist.
• Water We Drink: Water (H₂O) is a molecule made of two hydrogen atoms and one oxygen atom bonded together. Its unique molecular structure makes it vital for drinking, cooking, cleaning, and supporting all forms of life.
• Everyday Materials: From the food we eat (e.g., glucose molecules, C₆H₁₂O₆) to the clothes we wear (e.g., cotton, made of cellulose molecules), atoms and molecules are part of everything around us. Even plastics, metals, and medicines are composed of specific arrangements of atoms and molecules.

By understanding atoms and molecules, we gain insight into the composition of the world around us, from the air we breathe to the water we drink, making it a fundamental concept for students and scientists alike.

Laws of Chemical Combination

Law of Conservation of Mass

Statement: The Law of Conservation of Mass, proposed by Antoine Lavoisier, states that the total mass of the reactants in a chemical reaction is equal to the total mass of the products. Matter is neither created nor destroyed during a chemical reaction.

Example (Magnesium + Oxygen → Magnesium Oxide): When magnesium (Mg) reacts with oxygen (O₂) to form magnesium oxide (MgO), the mass remains conserved. For instance, if \(24 \, \text{g}\) of magnesium reacts with \(16 \, \text{g}\) of oxygen, the resulting magnesium oxide weighs \(40 \, \text{g} \, (24 + 16 = 40)\). The chemical equation is:

\(2Mg + O_{2} \;\;\longrightarrow\;\; 2MgO\)

Simple Experiment (Lavoisier’s Experiment): Antoine Lavoisier demonstrated this law by heating mercury in a sealed container with air. The mercury reacted with oxygen to form mercury oxide. He measured the mass of the container, mercury, and air before and after heating. The total mass remained unchanged, proving that mass is conserved in chemical reactions. A similar classroom experiment can involve burning a small piece of magnesium ribbon in a crucible and measuring the mass before and after to confirm the law. (Relevant to law of conservation of mass example.)

Law of Constant Proportion

Statement: The Law of Constant Proportion, proposed by Joseph Proust, states that a chemical compound always contains its constituent elements in a fixed ratio by mass, regardless of the source or method of preparation. (Relevant to law of constant proportion class 9.)

Example (Water = H₂O Always 1:8 Ratio by Mass): In water (\(H_{2}O\)), hydrogen and oxygen are always present in a mass ratio of \(1:8\). For example, in \(9 \, \text{g}\) of water, \(1 \, \text{g}\) is hydrogen and \(8 \, \text{g}\) is oxygen. This ratio holds true whether the water comes from a river, rain, or a lab synthesis. The chemical formula \(H_{2}O\) reflects this fixed proportion.

Real-Life Examples:

• Sugar: Table sugar (sucrose, \(C_{12}H_{22}O_{11}\)) always has carbon, hydrogen, and oxygen in a fixed mass ratio (approximately \(6:1:8\) by mass), whether extracted from sugarcane or sugar beets.
• Salt: Common salt (sodium chloride, \(NaCl\)) always contains sodium and chlorine in a fixed mass ratio of \(23:35.5\), whether obtained from seawater or a salt mine.

These laws form the foundation of chemistry, explaining how matter behaves during chemical reactions and ensuring consistency in the composition of compounds. (Relevant to laws of chemical combination class 9.)

Dalton’s Atomic Theory

Postulates

John Dalton proposed his atomic theory in 1808, laying the foundation for modern chemistry. The key postulates of Dalton’s Atomic Theory are: (Relevant to postulates of dalton atomic theory and dalton’s atomic theory class 9.)

• All matter is composed of extremely small, indivisible particles called atoms.
• Atoms of the same element are identical in size, mass, and chemical properties, but differ from atoms of other elements.
• Atoms cannot be created, destroyed, or divided into smaller particles during chemical reactions.
• Atoms combine in simple, whole-number ratios to form compounds (e.g., water is formed by hydrogen and oxygen atoms in a \(2:1\) ratio).
• In chemical reactions, atoms are rearranged, combined, or separated, but their total number remains conserved.

Limitations

While Dalton’s theory was groundbreaking, later discoveries revealed its limitations: (Relevant to limitations of dalton theory.)

• Inability to Explain Isotopes: Dalton assumed all atoms of an element have identical masses. However, isotopes (e.g., \(^{12}C\) and \(^{14}C\)) are atoms of the same element with different masses due to varying numbers of neutrons, contradicting Dalton’s postulate.
• Subatomic Particles: Modern discoveries showed that atoms are divisible into smaller particles like protons, neutrons, and electrons, challenging the idea of indivisible atoms.
• Failure to Explain Allotropes: Dalton’s theory could not account for allotropes (e.g., carbon in the form of diamond and graphite), where the same element exhibits different properties due to varying atomic arrangements.
• Chemical Bonding and Molecular Structure: The theory does not explain the nature of chemical bonds or why atoms combine in specific ways, which later models like the electron cloud model addressed.
• Nuclear Reactions: Dalton’s idea that atoms cannot be created or destroyed does not apply to nuclear reactions, where atoms can undergo fission or fusion, altering their identity.

Despite these limitations, Dalton’s Atomic Theory provided a fundamental framework for understanding matter and paved the way for advancements in atomic and molecular science.

Atoms and Molecules

Atoms

Definition of Atom: An atom is the smallest unit of an element that retains its chemical properties and participates in chemical reactions. It consists of a nucleus (containing protons and neutrons) surrounded by electrons. Atoms are the fundamental building blocks of all matter, too small to be seen without advanced microscopes, and are measured in nanometers (\(1 \, \text{nm} = 10^{-9} \, \text{m}\)). (Relevant to atom and molecule difference.)

Symbol Representation: Atoms are represented by chemical symbols based on their element names, as per the periodic table. Examples include:

• \(H\): Hydrogen
• \(O\): Oxygen
• \(Na\): Sodium
• \(Cl\): Chlorine

These symbols are universal and used in chemical equations to denote specific elements.

Atomic Mass Unit (amu): The atomic mass unit (amu), also called a dalton, is a standard unit for measuring the mass of atoms. It is defined as one-twelfth the mass of a carbon-12 atom, approximately \(1.66 \times 10^{-27} \, \text{kg}\). For example, hydrogen has an atomic mass of about \(1 \, \text{amu}\), and oxygen has about \(16 \, \text{amu}\). This unit helps compare the relative masses of atoms. (Relevant to atomic mass unit class 9.)

Molecules

Definition of Molecule: A molecule is a group of two or more atoms chemically bonded together, forming a single unit with distinct properties. These bonds can be covalent or ionic, and molecules represent the smallest unit of a compound that retains its chemical properties. For example, a water molecule (\(H_{2}O\)) consists of two hydrogen atoms and one oxygen atom. (Relevant to atom and molecule difference.)

Types of Molecules:

• Monoatomic: Consists of a single atom, typically noble gases. Example: \(He\) (Helium).
• Diatomic: Consists of two atoms bonded together. Example: \(O_{2}\) (Oxygen).
• Polyatomic: Consists of three or more atoms bonded together. Example: \(CO_{2}\) (Carbon dioxide).

Examples:

• \(O_{2}\): A diatomic molecule of oxygen, where two oxygen atoms are bonded, essential for respiration.
• \(CO_{2}\): A polyatomic molecule with one carbon atom and two oxygen atoms, produced during combustion and respiration.
• \(H_{2}O\): A polyatomic molecule with two hydrogen atoms and one oxygen atom, forming water, vital for life.

Understanding atoms and molecules is key to grasping how matter is structured and behaves in chemical reactions, forming the basis of chemistry. (Relevant to molecule examples class 9.)

Atomic Mass and Molecular Mass

Definition of Atomic Mass

The atomic mass of an element is the average mass of its atoms, expressed in atomic mass units (amu), where \(1 \, \text{amu} \approx \tfrac{1}{12} \, \text{mass of a carbon-12 atom} \approx 1.66 \times 10^{-27} \, \text{kg}\). It accounts for the weighted average of all naturally occurring isotopes of the element. For example, the atomic mass of hydrogen is approximately \(1 \, \text{amu}\), and that of oxygen is about \(16 \, \text{amu}\). Understanding atomic mass is crucial for studying chemical reactions and compositions in chemistry, particularly for students exploring the topic in class 9. (Relevant to atomic mass class 9.)

Molecular Mass

The molecular mass of a compound is the sum of the atomic masses of all the atoms in its molecule or formula unit, also measured in amu. For molecular compounds, it is calculated by adding the atomic masses of the constituent atoms as per the molecular formula. For ionic compounds, it is often referred to as the formula mass, but the calculation method remains the same. This concept helps determine the mass of substances involved in chemical reactions. (Relevant to formula mass.)

Examples of Molecular Mass Calculation

Here are some practical examples to illustrate how molecular mass is calculated (relevant to molecular mass examples):

• Water (\(H_{2}O\)):

  Molecular formula: \(H_{2}O\)
  Atomic mass of hydrogen (\(H\)) = \(1 \, \text{amu}\)
  Atomic mass of oxygen (\(O\)) = \(16 \, \text{amu}\)
  Molecular mass = \((2 \times 1) + (1 \times 16) = 2 + 16 = 18 \, \text{amu}\)

• Carbon Dioxide (\(CO_{2}\)):

  Molecular formula: \(CO_{2}\)
  Atomic mass of carbon (\(C\)) = \(12 \, \text{amu}\)
  Atomic mass of oxygen (\(O\)) = \(16 \, \text{amu}\)
  Molecular mass = \((1 \times 12) + (2 \times 16) = 12 + 32 = 44 \, \text{amu}\)

• Sodium Chloride (\(NaCl\)):

  Formula unit: \(NaCl\)
  Atomic mass of sodium (\(Na\)) = \(23 \, \text{amu}\)
  Atomic mass of chlorine (\(Cl\)) = \(35.5 \, \text{amu}\)
  Formula mass = \((1 \times 23) + (1 \times 35.5) = 23 + 35.5 = 58.5 \, \text{amu}\)

These calculations of atomic and molecular mass are essential for understanding the quantitative aspects of chemical reactions and the composition of substances in everyday life, such as water, carbon dioxide, and table salt.

Mole Concept

Definition of Mole

A mole is a fundamental unit in chemistry that represents a specific quantity of particles, such as atoms, molecules, or ions. One mole contains exactly \( 6.022 \times 10^{23} \) particles, a value known as Avogadro’s number (denoted as \( N_A \)). This constant allows chemists to count particles by weighing substances, making it essential for understanding chemical quantities, especially for students exploring the mole concept in class 9. (Relevant to avogadro number class 9.)

Relation Between Mole, Mass, and Number of Particles

The mole concept links the number of particles, mass, and molar mass of a substance through two key formulas:

• Moles from mass: \[ n = \frac{\text{Given mass}}{\text{Molar mass}} \] where \( n \) = number of moles, given mass in grams, and molar mass in g/mol.
• Number of particles: \[ \text{Number of particles} = n \times N_A \] where \( N_A = 6.022 \times 10^{23} \). (Relevant to mole concept class 9.)

Numericals on Mole Concept

Below are example calculations to demonstrate the application of the mole concept: (Relevant to numericals on mole concept.)

Example 1: Calculate the number of moles in 9 g of water.

• Given: Mass of water (H\(_2\)O) = 9 g.
• Molar mass of H\(_2\)O: \[ (2 \times 1) + (1 \times 16) = 18 \, \text{g/mol} \]
• Formula: \[ n = \frac{9}{18} \]
• Calculation: \[ n = 0.5 \, \text{moles} \]
• Answer: There are 0.5 moles in 9 g of water.

Example 2: Find the number of molecules in 22 g of CO\(_2\).

• Given: Mass of CO\(_2\) = 22 g.
• Molar mass of CO\(_2\): \[ (1 \times 12) + (2 \times 16) = 44 \, \text{g/mol} \]
• Step 1: Calculate moles: \[ n = \frac{22}{44} = 0.5 \, \text{moles} \]
• Step 2: Calculate molecules: \[ \text{Number of molecules} = 0.5 \times 6.022 \\ \times 10^{23} = 3.011 \times 10^{23} \]
• Answer: There are \( 3.011 \times 10^{23} \) molecules in 22 g of CO\(_2\).

Example 3: Mass of 2 moles of NaCl.

• Given: \( n = 2 \, \text{moles} \).
• Molar mass of NaCl: \[ (1 \times 23) + (1 \times 35.5) = 58.5 \, \text{g/mol} \]
• Formula: \[ \text{Mass} = n \times \text{Molar mass} \]
• Calculation: \[ \text{Mass} = 2 \times 58.5 = 117 \, \text{g} \]
• Answer: The mass of 2 moles of NaCl is 117 g.

Writing Chemical Formulae

Valency

Definition of Valency: Valency is the combining capacity of an element, determined by the number of electrons an atom can gain, lose, or share to achieve a stable electron configuration. It indicates how many bonds an atom can form with other atoms. Valency is a key concept for students learning to write chemical formulae in class 9, as it helps predict how elements combine to form compounds. (Relevant to valency class 9.)

Examples of Valency

• Hydrogen (H): Valency = 1 (can form one bond by sharing or losing one electron).
• Oxygen (O): Valency = 2 (can form two bonds by gaining or sharing two electrons).
• Sodium (Na): Valency = 1 (loses one electron to form a positive ion).
• Chlorine (Cl): Valency = 1 (gains or shares one electron to form a negative ion).

Rules for Writing Formulae

To write the chemical formula of a compound, follow these steps, often using the cross-method to ensure the correct ratio of atoms based on their valencies. This method is widely taught when learning about the formula of compounds in class 9. (Relevant to writing chemical formula class 9 and formula of compounds class 9.)

Cross-Method Explanation

1. Write the symbols of the elements or ions involved, with the positive ion (usually a metal or hydrogen) on the left and the negative ion (usually a non-metal) on the right.
2. Write the valency of each element or ion above its symbol.
3. Cross-multiply the valencies: the valency of the first element becomes the subscript of the second, and vice versa.
4. Simplify the ratio of subscripts to the smallest whole numbers, if possible.
5. Omit the subscript if it is 1.
6. For polyatomic ions, enclose them in parentheses if their subscript is greater than 1.

Examples of Writing Formulae

1. Sodium Chloride (NaCl)

• Elements: Sodium (Na, valency = 1), Chlorine (Cl, valency = 1).
• Cross-method: \[ \text{Na}^1 \text{Cl}^1 \;\;\rightarrow\;\; \text{Na}_1 \text{Cl}_1 \;\;\rightarrow\;\; \text{NaCl} \] (Subscript 1 omitted)

2. Water (H₂O)

• Elements: Hydrogen (H, valency = 1), Oxygen (O, valency = 2).
• Cross-method: \[ \text{H}^1 \text{O}^2 \;\;\rightarrow\;\; \text{H}_2 \text{O}_1 \;\;\rightarrow\;\; \text{H}_2\text{O} \] (Subscript 1 omitted)

3. Calcium Chloride (CaCl₂)

• Elements: Calcium (Ca, valency = 2), Chlorine (Cl, valency = 1).
• Cross-method: \[ \text{Ca}^2 \text{Cl}^1 \;\;\rightarrow\;\; \text{Ca}_1 \text{Cl}_2 \;\;\rightarrow\;\; \text{CaCl}_2 \] (Subscript 1 omitted)

4. Aluminium Oxide (Al₂O₃)

• Elements: Aluminium (Al, valency = 3), Oxygen (O, valency = 2).
• Cross-method: \[ \text{Al}^3 \text{O}^2 \;\;\rightarrow\;\; \text{Al}_2 \text{O}_3 \] (Simplified ratio, no further reduction needed)

These rules and the cross-method provide a systematic way to write chemical formulae, ensuring accurate representation of compounds based on the valency of their constituent elements.

NCERT Exercise Solutions (Important Questions)

Short Answer Questions

Q1. Define atom and molecule.

• Atom: An atom is the smallest unit of an element that retains its chemical properties. It cannot be further divided into simpler substances by ordinary chemical means. For example, a hydrogen atom (H) represents the smallest particle of hydrogen.
• Molecule: A molecule is formed when two or more atoms combine chemically. It can be of the same element (O₂, H₂) or different elements (H₂O, CO₂). Molecules are the basic units of compounds that exhibit independent existence.

Q2. State Dalton’s atomic theory.

Dalton’s atomic theory (1808) proposed the following main points:

1. All matter is made up of indivisible particles called atoms.
2. Atoms of the same element are identical in mass, size, and properties.
3. Atoms of different elements have different masses and properties.
4. Atoms combine in simple whole-number ratios to form compounds.
5. Atoms cannot be created or destroyed in chemical reactions, only rearranged.

This theory laid the foundation for modern atomic structure.

Q3. Difference between law of conservation of mass and law of constant proportion.

Law	Statement	Example
Law of Conservation of Mass	Mass is neither created nor destroyed in a chemical reaction.	In the reaction: H₂ + O₂ → H₂O, the total mass of reactants equals the total mass of products.
Law of Constant Proportion	A chemical compound always contains elements in a fixed ratio by mass.	In water (H₂O), hydrogen and oxygen are always in the ratio 1:8 by mass, irrespective of source.

Long Answer Questions

Q4. Explain mole concept with examples.

The mole concept is a way to count particles in chemistry.

• One mole of a substance contains 6.022 × 10²³ particles (Avogadro’s number). These particles may be atoms, molecules, or ions.

Examples:

• 1 mole of oxygen atoms (O) = 6.022 × 10²³ atoms of oxygen.
• 1 mole of oxygen gas (O₂) = 6.022 × 10²³ O₂ molecules.
• 1 mole of water (H₂O) = 18 g = 6.022 × 10²³ molecules of water.

The mole concept helps in relating mass, number of particles, and volume of gases at standard conditions.

Q5. Derive chemical formula using valency.

The valency of an element determines how atoms combine to form compounds. The rules to derive chemical formula are:

1. Write symbols of the elements.
2. Write their valencies below the symbols.
3. Cross-multiply valencies to balance charges.
4. Simplify if possible.

Examples:

• Sodium (Na, valency 1) + Chlorine (Cl, valency 1) → Formula: NaCl
• Calcium (Ca, valency 2) + Oxygen (O, valency 2) → Formula: CaO
• Aluminium (Al, valency 3) + Oxygen (O, valency 2) → Cross multiplication → Al₂O₃

Thus, chemical formulae represent the ratio in which atoms combine.

Numericals

Q6. Solve mole concept problems from NCERT.

1. How many moles are present in 9 g of water?
   Molar mass of H₂O = 18 g/mol
   Moles = Mass ÷ Molar mass = 9 ÷ 18 = 0.5 mol
2. Calculate the number of atoms in 46 g of sodium (Na).
   Molar mass of Na = 23 g/mol
   Moles = 46 ÷ 23 = 2 mol
   Atoms = 2 × 6.022 × 10²³ = 1.204 × 10²⁴ atoms

Q7. Extra numericals for practice.

1. Find the mass of 3 moles of carbon dioxide (CO₂).
2. Calculate the number of molecules in 36 g of water.
3. How many moles are there in 11.2 L of oxygen gas at STP?
4. Find the number of atoms in 0.5 moles of magnesium.
5. Determine the molecular mass of H₂SO₄ and calculate the number of moles in 98 g of H₂SO₄.

This covers atoms and molecules important questions class 9 along with numericals class 9 science chapter 3 for exam preparation.

FAQs on Atoms and Molecules (Class 9)

This section covers class 9 atoms and molecules FAQs with a focus on laws of chemical combination questions and mole concept easy explanation for exam preparation.

Q1. What is Dalton’s atomic theory class 9?

Dalton’s atomic theory states that all matter is made up of tiny indivisible particles called atoms. Atoms of the same element are identical in mass and properties, while atoms of different elements differ in mass and properties. Atoms combine in simple whole-number ratios to form compounds and cannot be created or destroyed during a chemical reaction. This theory forms the basis of modern chemistry.

Q2. What is mole concept in simple words?

The mole concept is a way of counting particles in chemistry. In simple words, one mole of any substance contains 6.022 × 10²³ particles (atoms, molecules, or ions). For example, 1 mole of water = 18 g = 6.022 × 10²³ molecules of H₂O. It helps in easily calculating mass, volume, and number of particles in a substance. (Relevant for mole concept easy explanation).

Q3. What is the difference between atom and molecule?

• Atom: The smallest unit of an element that cannot be further divided by ordinary chemical means. Example: H, O, Na.
• Molecule: A group of two or more atoms chemically bonded together. It can be of the same element (O₂, N₂) or of different elements (H₂O, CO₂).

Q4. Give 2 examples of laws of chemical combination.

1. Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.
   Example: In H₂ + O₂ → H₂O, total mass of reactants = total mass of products.
2. Law of Constant Proportion: A compound always contains elements in a fixed ratio by mass.
   Example: In water (H₂O), hydrogen and oxygen are always present in the ratio 1:8 by mass.

(Relevant for laws of chemical combination questions in class 9).

Q5. Write chemical formula of calcium chloride.

Calcium (Ca) has valency 2 and chlorine (Cl) has valency 1. By cross multiplication, the formula becomes CaCl₂. This means one calcium atom combines with two chlorine atoms to form calcium chloride.

Summary (Quick Revision Notes)

Laws of Chemical Combination

• Law of Conservation of Mass: Mass can neither be created nor destroyed in a chemical reaction.
• Law of Constant Proportion: A compound always contains elements in a fixed ratio by mass, irrespective of its source.

Dalton’s Atomic Theory (5 Main Postulates)

1. All matter is made up of indivisible particles called atoms.
2. Atoms of the same element are identical in mass and properties.
3. Atoms of different elements have different masses and properties.
4. Atoms combine in simple whole-number ratios to form compounds.
5. Atoms cannot be created or destroyed in a chemical reaction, only rearranged.

Key Terms

• Atom: Smallest unit of an element that retains its properties.
• Molecule: Combination of two or more atoms chemically bonded (O₂, H₂O).
• Atomic Mass: Mass of a single atom of an element expressed in atomic mass units (amu).
• Molecular Mass: Sum of atomic masses of all atoms in a molecule.

Mole Concept and Avogadro’s Number

• Mole: The SI unit for amount of substance. One mole contains 6.022 × 10²³ particles (Avogadro’s number).
• 1 mole of atoms = 6.022 × 10²³ atoms.
• 1 mole of molecules = 6.022 × 10²³ molecules.
• 1 mole of gas at STP = 22.4 L.

Chemical Formula Writing Rules

1. Write the symbols of elements involved.
2. Note their valencies.
3. Cross-multiply the valencies to balance charges.
4. Simplify to get the simplest whole-number ratio.

Examples: Na (1) + Cl (1) → NaCl; Ca (2) + Cl (1) → CaCl₂; Al (3) + O (2) → Al₂O₃