NCERT Science Notes - Class 9
Chapter 3 - Atoms and Molecules

Welcome to AJs Chalo Seekhen. This webpage is dedicated to Class 9 | Science | Chapter 3 - Atoms and Molecules. The chapter delves about the basic building blocks of matter: atoms and molecules. It covers the concept of atoms as the smallest unit of an element and molecules as combinations of atoms. Key topics include laws of chemical combination, Dalton’s atomic theory, the atomic and molecular masses, and the mole concept. The chapter explains how to write chemical formulas and the role of atoms and molecules in various chemical reactions. Understanding these fundamental concepts helps students grasp the basics of chemistry and the composition of substances in their surroundings. 🧪📚

NCERT Science Notes - Class 8 Chapter 9 - Friction notes ajs, cbse notes class 10 ajslearning, cbse notes ajs, ajs notes class 10, ajslearning, ajs chalo seekhen

NOTES

NCERT Science Notes - Class 9
Chapter 3 - Atoms and Molecules

3.0 - Introduction

Ancient Philosophical Theories on Matter:

Indian Philosophers:
- Around 500 BC, Indian philosopher Maharishi Kanad proposed that matter (padarth) could be divided repeatedly until reaching indivisible particles, which he called Parmanu (atoms).
- Pakudha Katyayama, another Indian philosopher, suggested that these indivisible particles usually exist in a combined form, creating various forms of matter.
Greek Philosophers:
- Around the same time, Greek philosophers Democritus and Leucippus also theorized about the divisibility of matter.
- Democritus introduced the term atomos, meaning indivisible, for the smallest particles that could no longer be divided.

Nature of Early Theories:

These ideas about matter and atoms were based on philosophical reasoning rather than scientific experiments.
There was no experimental validation for these theories until the 18th century, when advances in scientific methods allowed for testing and verification.

Recognition of Elements and Compounds:

By the late 18th century, scientists began to distinguish between elements (substances that cannot be broken down) and compounds (substances formed from two or more elements).
This distinction led to an interest in understanding how and why elements combine to form compounds.

Antoine Lavoisier’s Contributions:

Antoine L. Lavoisier, a French chemist, laid the foundation for modern chemical science.
He established the Laws of Chemical Combination, which describe how elements combine and interact in chemical reactions, setting the stage for the development of atomic theory.

3.1 - Laws of Chemical Combination

The laws of chemical combination describe how substances combine and react to form new substances. These laws were established by Antoine Lavoisier and Joseph L. Proust through extensive experimentation.

3.1.1 - Law of Conservation of Mass

Definition:
- The Law of Conservation of Mass states that mass can neither be created nor destroyed in a chemical reaction. This means that the total mass of the reactants before a reaction is equal to the total mass of the products after the reaction.
Explanation:
- During a chemical reaction, atoms are rearranged to form new substances, but the total number of atoms remains the same. As a result, the total mass does not change, as atoms are neither lost nor gained.
Example Activity (3.1) to Demonstrate the Law:
- Objective: To observe if there is a change in mass during a chemical reaction.
- Materials Needed:
  - Two sets of chemicals: (i) copper sulphate and sodium carbonate, (ii) barium chloride and sodium sulphate, or (iii) lead nitrate and sodium chloride.
- Procedure:
  1. Prepare a 5% solution of one set of chemicals, X and Y, in separate containers, each containing 10 mL of water.
  2. Place the solution of Y in a conical flask and suspend the solution of X in an ignition tube inside the flask, ensuring they do not mix initially.
  3. Seal the flask with a cork to prevent any escape of substances.
  4. Weigh the flask with its contents.
  5. Tilt the flask to allow the solutions to mix and observe the reaction.
  6. Weigh the flask again after the reaction.
- Observations and Questions:
  - What happens in the reaction flask? You may observe changes like color formation or precipitation, indicating a chemical reaction.
  - Does the mass of the flask and its contents change? According to the law, the mass should remain the same.
  - Why should we put a cork on the flask? The cork ensures that no gas or liquid escapes, maintaining the total mass.

This activity demonstrates that even though substances change form during a reaction, the total mass remains constant, verifying the Law of Conservation of Mass.

3.1.2 - Law of Constant Proportions

(Questionnaire Format)

What is the Law of Constant Proportions?
- This law, also known as the Law of Definite Proportions, states that in any chemical substance, elements are always present in definite and constant proportions by mass, regardless of the source or method of preparation.
Who formulated the Law of Constant Proportions, and what observations led to this law?
- Joseph L. Proust formulated this law after observing that compounds are composed of elements in fixed mass ratios.
- For example:
  - In water (H₂O), the mass ratio of hydrogen to oxygen is always 1:8. If 9 grams of water is decomposed, it will yield 1 gram of hydrogen and 8 grams of oxygen.
  - In ammonia (NH₃), nitrogen and hydrogen are present in a constant mass ratio of 14:3, no matter the source.
How did John Dalton contribute to the explanation of this law?
- John Dalton provided a theory to explain why elements combine in fixed proportions. His Atomic Theory suggested that matter consists of indivisible particles called atoms, which combine in definite ratios to form compounds.

Dalton’s Atomic Theory

Who was John Dalton, and how did he develop his atomic theory?
- John Dalton (1766–1844) was a British chemist born into a poor family. He began teaching at twelve and later became a school principal. In 1808, he presented his atomic theory, which fundamentally changed the study of matter.
- Dalton's theory built on the philosophical ideas of ancient Greeks and incorporated the laws of chemical combination.
What are the postulates of Dalton’s Atomic Theory?
- All matter is composed of very tiny particles called atoms, which participate in chemical reactions.
- Atoms are indivisible particles; they cannot be created or destroyed during chemical reactions.
- Atoms of a given element are identical in terms of mass and chemical properties.
- Atoms of different elements have distinct masses and chemical properties.
- Atoms combine in small whole-number ratios to form compounds, explaining the fixed ratios observed in chemical combinations.
- In a given compound, the relative number and kinds of atoms remain constant, accounting for the consistency in chemical composition.
How does Dalton’s Atomic Theory support the laws of chemical combination?
- Law of Conservation of Mass: Since atoms are indivisible and cannot be created or destroyed, the total mass remains constant in a chemical reaction.
- Law of Constant Proportions: The fixed, whole-number ratios in which atoms combine explain why compounds always have elements in consistent mass proportions.

Dalton’s atomic theory laid the foundation for modern chemistry by explaining these fundamental laws. Even though later discoveries revealed that atoms are composed of smaller particles, his theory was crucial for advancing our understanding of matter.

3.2 - What is an Atom?

(Questionnaire Format)

What is an atom and why is it important?
- An atom is the basic building block of all matter, similar to how a small grain of sand can be the building block of an ant-hill or bricks form a wall.
- Atoms are incredibly small, yet they form the foundation of everything in the universe. Despite their tiny size, atoms are essential as they make up all substances around us.
How small are atoms?
- Atoms are smaller than anything we can directly observe or easily compare. To visualize:
  - Stacking millions of atoms would create a layer barely as thick as a single sheet of paper.
- Atomic radius is measured in nanometers (nm), where:
  - $1 \text{ nm} = \frac{1}{10^9} \text{ meters}$ 1 nm=1091 meters
  - $1 \text{ meter} = 10^9 \text{ nanometers}$ 1 meter=109 nanometers
What are the relative sizes of different particles and objects?
- Here are some examples of sizes measured in meters:
  - 10⁻¹⁰ m – Size of a hydrogen atom
  - 10⁻⁹ m – Size of a water molecule (H₂O)
  - 10⁻⁸ m – Size of a hemoglobin molecule
  - 10⁻⁴ m – Size of a grain of sand
  - 10⁻³ m – Size of an ant
  - 10⁻¹ m – Size of an apple
Why should we care about atoms despite their small size?
- Although atoms are incredibly small, they constitute everything we interact with in our daily lives. From the air we breathe to the food we eat, all are made up of atoms.
- Advances in technology now allow us to observe and even produce magnified images of atomic structures, which helps us understand material properties and chemical reactions at a fundamental level.

Understanding atoms allows us to grasp how matter is structured and behaves, making it an essential concept in both science and everyday life.

3.2.1 - What are the Modern-Day Symbols of Atoms of Different Elements?
(Questionnaire Format)

How were the symbols for elements originally developed?
- John Dalton was the first scientist to use symbols to represent elements, associating each symbol with a specific quantity, meaning one atom of the element.
- Later, Jöns Jakob Berzelius proposed that element symbols should be derived from one or two letters of the element’s name, which is a convention still used today.
What are some historical origins of element names?
- Element names often originated from the place of discovery or from colors associated with the element:
  - Copper was named after Cyprus, where it was first found.
  - Gold was named based on the English word meaning "yellow."
Who is responsible for approving element names and symbols today?
- The International Union of Pure and Applied Chemistry (IUPAC) is the scientific organization responsible for approving the names, symbols, and units of elements globally.
What are the general rules for element symbols?
- Most element symbols are derived from the first one or two letters of the element’s English name.
- The first letter is always capitalized, while the second letter, if present, is lowercase.
  - Examples:
    - Hydrogen: H
    - Aluminium: Al (not AL)
    - Cobalt: Co (not CO)
- Some symbols use the first letter and another letter appearing later in the name.
  - Examples:
    - Chlorine: Cl
    - Zinc: Zn
Why do some elements have symbols derived from Latin, German, or Greek names?
- Symbols for certain elements reflect their names in other languages, mainly due to historical or traditional reasons:
  - Iron – Symbol: Fe, derived from the Latin name ferrum
  - Sodium – Symbol: Na, from the Latin name natrium
  - Potassium – Symbol: K, from the Latin name kalium
What is unique about each element's symbol?
- Each element has a unique chemical symbol, which helps scientists and chemists identify elements accurately and consistently, regardless of language or regional differences.

This system of symbols allows for a universal language in chemistry, enabling clear and precise communication about elements and their properties across the scientific community.

This table presents symbols for some common elements, organized as follows:

Element	Symbol	Element	Symbol	Element	Symbol
Aluminium	Al	Copper	Cu	Nitrogen	N
Argon	Ar	Fluorine	F	Oxygen	O
Barium	Ba	Gold	Au	Potassium	K
Boron	B	Hydrogen	H	Silicon	Si
Bromine	Br	Iodine	I	Silver	Ag
Calcium	Ca	Iron	Fe	Sodium	Na
Carbon	C	Lead	Pb	Sulphur	S
Chlorine	Cl	Magnesium	Mg	Uranium	U
Cobalt	Co	Neon	Ne	Zinc	Zn

This table illustrates the systematic approach of assigning symbols to elements, usually based on one or two letters derived from their English or Latin names. It serves as a quick reference for identifying elements by their standardized symbols.

3.2.2 - Atomic Mass

(Questionnaire Format)

What concept did Dalton’s atomic theory introduce regarding elements?
- Dalton’s theory introduced the idea that each element has a characteristic atomic mass, which helped explain why elements combine in fixed ratios, as seen in the Law of Constant Proportions.
Why was it challenging to determine the atomic mass of an atom?
- Measuring the mass of an individual atom directly was difficult due to the atom's extremely small size. Instead, scientists determined relative atomic masses using the masses of compounds and chemical reactions.
How was relative atomic mass initially determined?
- For example, in the compound carbon monoxide (CO):
  - It was found that 3 g of carbon combines with 4 g of oxygen.
  - Thus, carbon combines with $\frac{4}{3}$ times its mass of oxygen.
- If we initially assume the atomic mass of carbon is 1.0 u (unified mass unit), then oxygen's atomic mass would be about 1.33 u to maintain this ratio. However, this approach was later refined to ensure atomic masses were closer to whole numbers.
What was the initial standard used for atomic mass units?
- Scientists originally used 1/16th of the mass of naturally occurring oxygen as the atomic mass unit, which:
  - Was convenient because oxygen forms compounds with many elements.
  - Allowed for whole numbers as atomic masses for most elements.
What is the current standard for atomic mass, and why was it chosen?
- In 1961, the carbon-12 isotope was selected as the standard for atomic mass measurement:
  - One atomic mass unit (u) is defined as one-twelfth (1/12) the mass of a carbon-12 atom.
- This provided a universally accepted reference for atomic masses, allowing all elements' atomic masses to be expressed relative to carbon-12.
What analogy helps explain the concept of relative atomic mass?
- Imagine a fruit seller who uses a watermelon as a unit of measure, assigning it a mass of 12 units. He divides it into 12 equal pieces and then measures other fruits’ masses relative to one piece. In this analogy:
  - Each piece of watermelon represents an atomic mass unit (fmu).
  - Similarly, scientists use carbon-12 as a standard "unit" to measure the relative masses of atoms.

Using the carbon-12 isotope as the standard provides a consistent and widely accepted way to measure atomic mass, which is fundamental in chemistry for understanding the behavior and properties of elements.

This table provides the atomic masses of a few common elements, measured in unified atomic mass units (u), where 1 u is equal to one-twelfth the mass of a carbon-12 atom:

Element	Atomic Mass (u)
Hydrogen	1
Carbon	12
Nitrogen	14
Oxygen	16
Sodium	23
Magnesium	24
Sulphur	32
Chlorine	35.5
Calcium	40

The atomic mass of each element represents the average mass of its atoms relative to carbon-12. This standardized comparison allows for consistency in measuring and calculating chemical properties and reactions across different elements.

3.2.3 - How Do Atoms Exist?

Can atoms of elements exist independently?
- Most atoms cannot exist independently because they are generally too reactive or unstable on their own.
How do atoms form stable structures?
- Atoms form molecules and ions to achieve stability. These molecules or ions then combine in large quantities to create visible matter.
What are molecules?
- A molecule is a stable group of two or more atoms held together by chemical bonds. Molecules can be composed of atoms of the same element (e.g., $O_2$ O2) or different elements (e.g., $H_2O$ H2O).
What are ions?
- An ion is an atom or group of atoms that has gained or lost one or more electrons, resulting in a net electric charge. Ions can be positively charged (cations) or negatively charged (anions).
How do molecules and ions form matter?
- Molecules and ions aggregate, or come together in large numbers, to form various types of matter that we can see, feel, or touch. This aggregation leads to the formation of solids, liquids, and gases.

In essence, atoms combine to form stable compounds, which aggregate to create the matter that constitutes the world around us. This process enables atoms to achieve stability and form the substances that are essential to life and the environment.

3.3 - What is a Molecule?

What is a molecule?
- A molecule is a group of two or more atoms that are chemically bonded together by attractive forces. These bonds hold the atoms tightly together, forming a stable unit.
How is a molecule defined in terms of its properties?
- A molecule is the smallest particle of an element or a compound that can exist independently and retains all the properties of that substance. This means that a molecule exhibits the chemical characteristics of the substance it represents.
Can molecules consist of atoms from the same element?
- Yes, molecules can be formed by atoms of the same element. For example:
  - An oxygen molecule ( $O_2$ ) consists of two oxygen atoms bonded together.
  - A nitrogen molecule ( $N_2$ ) consists of two nitrogen atoms bonded together.
Can molecules consist of atoms from different elements?
- Yes, molecules can also be composed of different elements. These are called compounds. For example:
  - Water ( $H_2O$ ) is a molecule made of two hydrogen atoms and one oxygen atom.
  - Carbon dioxide ( $CO_2$ ) is a molecule made of one carbon atom and two oxygen atoms.

In summary, molecules are the fundamental units that can independently exist and exhibit the properties of elements and compounds. They can be simple, consisting of atoms of the same element, or complex, comprising different elements chemically bonded together.

3.3.1 - MOLECULES OF ELEMENTS

What constitutes the molecules of an element?
- The molecules of an element are composed of the same type of atoms. For example, an oxygen molecule ( $O_2$ O2) consists of two oxygen atoms.
Can molecules of some elements consist of single atoms?
- Yes, molecules of certain elements, particularly noble gases like argon (Ar) and helium (He), consist of single atoms. These are known as monoatomic molecules because they contain only one atom.
How do molecules form in nonmetals?
- Most nonmetals form molecules containing two or more atoms of the same element. For example:
  - Diatomic molecules: These are molecules with two atoms. Oxygen ( $O_2$ O2) and nitrogen ( $N_2$ N2) are examples of diatomic molecules, where two atoms of the same element bond together.
  - Triatomic molecules: When three atoms of oxygen bond together, they form ozone ( $O_3$ O3).
What is atomicity, and how does it vary among elements?
- Atomicity is the number of atoms present in a molecule of an element.
- Examples of atomicity for nonmetals:
  - Argon (Ar): 1 (monoatomic)
  - Oxygen ( $O_2$ ): 2 (diatomic)
  - Ozone ( $O_3$ ): 3 (triatomic)
- For metals and some elements like carbon, the atomicity is not simple because they consist of a large and indefinite number of atoms bonded in complex structures.
How do metals differ in structure compared to nonmetals?
- Metals and some other elements, such as carbon, do not form discrete molecules with a specific atomicity. Instead, they form metallic lattices or network structures, where atoms are bonded together in large, indefinite numbers.

Understanding the concept of atomicity and molecular structures helps to distinguish between different types of elemental molecules, from simple monoatomic gases to complex metallic lattices.

3.3.2 - Molecules of Compounds

What is a molecule of a compound?
- A molecule of a compound consists of atoms of different elements joined together in fixed proportions by mass. These proportions remain constant in each molecule of a specific compound.
How do elements combine in compounds?
- Elements combine in specific ratios to form compounds. For example:
  - Water (H₂O): Hydrogen and oxygen combine in a ratio of 1:8 by mass.
  - Ammonia (NH₃): Nitrogen and hydrogen combine in a ratio of 14:3 by mass.
  - Carbon dioxide (CO₂): Carbon and oxygen combine in a ratio of 3:8 by mass.
How can we find the ratio by the number of atoms in a compound?
- To determine the ratio by the number of atoms, use the atomic masses of the elements from Table 3.2 and the mass ratios from Table 3.4.
- Example Calculation for Water (H₂O) Atomic mass of hydrogen = 1 u, atomic mass of oxygen = 16 u.
  - For a mass ratio of 1:8 (H), calculate the number of atoms:
    - Mass of two hydrogen atoms (2 × 1 u) = 2 u.
    - Mass of one oxygen atom = 16 u.
  - Simplified ratio by number of atoms is H = 2 : 1
Activity 3.2: Calculate the ratio by number of atoms for each compound using their atomic masses and given mass ratios:
- Water (H₂O): H = 2:1
- Ammonia (NH₃): N = 1:3
- Carbon dioxide (CO₂): C = 1:2

Understanding the ratio by number of atoms helps illustrate how elements combine to form compounds, maintaining consistent proportions that define the compound's chemical structure and properties.

3.3.3 - WHAT IS AN ION?

What are ions?
- Ions are charged species that can consist of a single atom or a group of atoms with a net electrical charge. Ions form when atoms gain or lose electrons, resulting in a positive or negative charge.
What are the types of ions?
- Cations: Positively charged ions formed when atoms lose electrons.
  - Example: Sodium ion ( $\text{Na}^+$ )
- Anions: Negatively charged ions formed when atoms gain electrons.
  - Example: Chloride ion ( $\text{Cl}^-$ )
What are ionic compounds?
- Ionic compounds are formed from the electrostatic attraction between cations and anions. They typically consist of metals (forming cations) and nonmetals (forming anions).
- Examples of ionic compounds and their constituting elements with mass ratios:
  - Calcium Oxide (CaO): Calcium and oxygen with a mass ratio of 5:2.
  - Magnesium Sulphide (MgS): Magnesium and sulphur with a mass ratio of 3:4.
  - Sodium Chloride (NaCl): Sodium and chlorine with a mass ratio of 23:35.5.
What are polyatomic ions?
- Polyatomic ions are groups of atoms bonded together that carry an overall charge. These ions act as a single charged entity in chemical reactions.
- Examples of polyatomic ions will be explored further in the next chapter.

Understanding ions and their role in forming ionic compounds is essential, as these compounds exhibit unique properties and are a major class of substances in chemistry.

3.4 - Writing Chemical Formulae

What is a chemical formula?
- A chemical formula is a symbolic representation of the composition of a compound, showing the elements involved and their ratios.
What is valency, and why is it important for writing chemical formulae?
- Valency refers to the combining power or capacity of an element, indicating how many atoms of other elements an atom can bond with.
- Valency is essential for determining how atoms will combine to form compounds, ensuring that the resulting compound is stable.
What are the rules for writing chemical formulae?
- Balancing Charges: The charges or valencies of ions in a compound must balance, resulting in a neutral compound.
- Order of Elements: In compounds consisting of a metal and a non-metal, the metal's name or symbol is written first. For example:
  - Calcium oxide: $\text{CaO}$
  - Sodium chloride: $\text{NaCl}$
  - Iron sulphide: $\text{FeS}$
- Polyatomic Ions: If a compound contains polyatomic ions, and more than one ion is needed to balance the compound, the polyatomic ion is enclosed in parentheses with a subscript indicating the number of ions.
  - Example: $\text{Mg(OH)}_2$ for magnesium hydroxide.
  - No parentheses are needed if only one polyatomic ion is present, e.g., $\text{NaOH}$ for sodium hydroxide.
Activity: Using valencies from Table 3.6, practice writing chemical formulae:
- Determine the valency of each element or ion.
- Combine elements or ions so that the total positive and negative charges balance.
For example:
- Octopus Analogy: If O (octopus) has a valency of 8, and H (human) has a valency of 2, then O can combine with four H atoms. The formula would be $\text{OH}_4$ , indicating the number of atoms of each element.
Valencies of Common Ions (Table 3.6):
- Elements have different valencies which affect their bonding:
  - Sodium (Na): +1
  - Magnesium (Mg): +2
  - Aluminum (Al): +3
- Non-metallic elements:
  - Chloride (Cl): -1
  - Oxide (O): -2
- Polyatomic ions:
  - Ammonium ( $\text{NH}_4^+$ ): +1
  - Sulphate ( $\text{SO}_4^{2-}$ ): -2
  - Phosphate ( $\text{PO}_4^{3-}$ ): -3

These rules provide a systematic approach to writing correct chemical formulae for various compounds, ensuring that the elements and ions are combined in a way that reflects their chemical properties and stability.

Names and symbol of some ions

Valency	Name of ion	Symbol	Non-metallic element	Symbol	Polyatomic ions	Symbol
1	Sodium	Na⁺	Hydrogen	H⁺	Ammonium	NH₄⁺
	Potassium	K⁺	Hydride	H^-	Hydroxide	OH^-
	Silver	Ag⁺	Chloride	Cl^-	Nitrate	NO₃^-
	Copper (I)	Cu⁺	Bromide	Br^-	Hydrogen carbonate	HCO₃^-
			Iodide	I^-
2	Magnesium	Mg²⁺	Oxide	O^2-	Carbonate	CO₃^2-
	Calcium	Ca²⁺	Sulphide	S^2-	Sulphite	SO₃^2-
	Zinc	Zn²⁺			Sulphate	SO₄^2-
	Iron (II)	Fe²⁺
	Copper (II)	Cu²⁺
3	Aluminium	Al³⁺	Nitride	N^3-	Phosphate	PO₄^3-
	Iron (III)	Fe³⁺

3.4 - Writing Chemical Formulae

What is a chemical formula?
- A chemical formula is a symbolic representation of the composition of a compound, showing the elements involved and their ratios.
What is valency, and why is it important for writing chemical formulae?
- Valency refers to the combining power or capacity of an element, indicating how many atoms of other elements an atom can bond with.
- Valency is essential for determining how atoms will combine to form compounds, ensuring that the resulting compound is stable.
What are the rules for writing chemical formulae?
- Balancing Charges: The charges or valencies of ions in a compound must balance, resulting in a neutral compound.
- Order of Elements: In compounds consisting of a metal and a non-metal, the metal's name or symbol is written first. For example:
  - Calcium oxide: $\text{CaO}$
  - Sodium chloride: $\text{NaCl}$
  - Iron sulphide: $\text{FeS}$
- Polyatomic Ions: If a compound contains polyatomic ions, and more than one ion is needed to balance the compound, the polyatomic ion is enclosed in parentheses with a subscript indicating the number of ions.
  - Example: $\text{Mg(OH)}_2$ for magnesium hydroxide.
  - No parentheses are needed if only one polyatomic ion is present, e.g., $\text{NaOH}$ for sodium hydroxide.
Activity: Using valencies from Table 3.6, practice writing chemical formulae:
- Determine the valency of each element or ion.
- Combine elements or ions so that the total positive and negative charges balance.
For example:
- Octopus Analogy: If O (octopus) has a valency of 8, and H (human) has a valency of 2, then O can combine with four H atoms. The formula would be $\text{OH}_4$ , indicating the number of atoms of each element.
Valencies of Common Ions (Table 3.6):
- Elements have different valencies which affect their bonding:
  - Sodium (Na): +1
  - Magnesium (Mg): +2
  - Aluminum (Al): +3
- Non-metallic elements:
  - Chloride (Cl): -1
  - Oxide (O): -2
- Polyatomic ions:
  - Ammonium ( $\text{NH}_4^+$ ): +1
  - Sulphate ( $\text{SO}_4^{2-}$ ): -2
  - Phosphate ( $\text{PO}_4^{3-}$ ): -3

3.4.1 - Formulae of Simple Compounds

What are binary compounds?
- Binary compounds consist of only two different elements. These are the simplest types of compounds, where elements combine in definite ratios based on their valencies.
How are chemical formulae written for binary compounds?
- The chemical formula of a binary compound is written by listing the symbols of the constituent elements, with the cation (positively charged ion) first, followed by the anion (negatively charged ion).
- Use the criss-cross method to balance the valencies. The valency of each element is criss-crossed to become the subscript of the other element.
Examples of writing chemical formulae:
- Hydrogen Chloride (HCl):
  - Symbols: $H$ and $Cl$
  - Valencies: $H = 1$ , $Cl = 1$
  - Since the valencies are equal, the formula is $HCl$ .
- Hydrogen Sulphide ( $H2S):$
  - Symbols: $H$ and $S$
  - Valencies: $H = 1$ , $S = 2$
  - Criss-cross the valencies: $H_2S$ .
  - This shows two hydrogen atoms combine with one sulfur atom.
- Carbon Tetrachloride ( $CCl4):$
  - Symbols: $C$ and $Cl$
  - Valencies: $C = 4$ , $Cl = 1$
  - Criss-cross the valencies: $C C l_{4}$
  - Four chlorine atoms bond with one carbon atom.
- Magnesium Chloride ( $MgCl_2$ ):
  - Symbols: $Mg^{2+}$ and $Cl^-$
  - Charges: $Mg = +2$ , $Cl = -1$
  - Criss-cross the charges: $MgCl_2$ .
  - Two chloride ions are needed to balance one magnesium ion.
Why are charges not indicated in the final formula?
- In the chemical formula, only the ratios of the ions are shown. The charges are balanced by the criss-cross method, resulting in a neutral compound, so there is no need to indicate the charges.

This method ensures that chemical formulae are correctly written to reflect the combining capacities of elements, resulting in neutral and stable compounds.

More Examples of Writing Chemical Formulae

Formula for Aluminium Oxide ( $\text{Al}_2\text{O}_3$ ):
- Symbols: $\text{Al}$ and $\text{O}$
- Charges: Aluminium ( $\text{Al}^{3+}$ ) and Oxygen ( $\text{O}^{2-}$ )
- Criss-Cross Method:
  - The 3 from $\text{Al}^{3+}$ becomes the subscript for oxygen, and the 2 from $\text{O}^{2-}$ becomes the subscript for aluminium.
  - Formula: $\text{Al}_2\text{O}_3$ , indicating two aluminium atoms bond with three oxygen atoms.
Formula for Calcium Oxide ( $CaO):$
- Symbols: $\text{Ca}$ and $\text{O}$
- Charges: Calcium ( $\text{Ca}^{2+}$ ) and Oxygen ( $\text{O}^{2-}$ )
- Criss-Cross Method:
  - The charges are equal, so they can be simplified. Although $\text{Ca}_2\text{O}_2$ might be obtained by criss-crossing, it reduces to $\text{CaO}$ .
  - Formula: $\text{CaO}$ , meaning one calcium atom bonds with one oxygen atom.

This approach ensures the correct ratio of ions in the compound, balancing charges to form a neutral substance. Reducing subscripts when the valencies are equal is standard practice to simplify the formula to its most basic form.

Additional Examples of Writing Chemical Formulae with Polyatomic Ions

Formula for Sodium Carbonate ( $\text{Na}_2\text{CO}_3$ ):
- Symbols: $\text{Na}$ (Sodium) and $\text{CO}_3$ (Carbonate)
- Charges: Sodium ( $\text{Na}^+$ ) and Carbonate ( $\text{CO}_3^{2-}$ )
- Criss-Cross Method:
  - The charge of 2 from carbonate becomes the subscript for sodium, indicating two sodium ions are needed to balance one carbonate ion.
  - Formula: $\text{Na}_2\text{CO}_3$ .
Formula for Ammonium Sulphate ( $(\text{NH}_4)_2\text{SO}_4$ ):
- Symbols: $\text{NH}_4$ (Ammonium) and $\text{SO}_4$ (Sulphate)
- Charges: Ammonium ( $\text{NH}_4^+$ ) and Sulphate ( $\text{SO}_4^{2-}$ )
- Criss-Cross Method:
  - The charge of 2 from sulphate becomes the subscript for ammonium, indicating two ammonium ions are required to balance one sulphate ion.
  - Formula: $(\text{NH}_4)_2\text{SO}_4$ , with brackets around $\text{NH}_4$ because there are two ammonium ions.
Formula for Calcium Hydroxide ( $\text{Ca(OH)}_2$ ):
- Symbols: $\text{Ca}$ (Calcium) and $\text{OH}$ (Hydroxide)
- Charges: Calcium ( $\text{Ca}^{2+}$ ) and Hydroxide ( $\text{OH}^-$ )
- Criss-Cross Method:
  - The charge of 2 from calcium means two hydroxide ions are needed.
  - Formula: $\text{Ca(OH)}_2$ Brackets are used around $\text{OH}$ to indicate there are two hydroxide groups.

Note: Brackets are used around polyatomic ions when there is more than one of the ion in the formula. If there is only one polyatomic ion, brackets are not necessary. This ensures clarity in indicating the number of each group in the compound.

3.5 - Molecular Mass

3.5.1 - MOLECULAR MASS

What is molecular mass?
- Molecular mass is the sum of the atomic masses of all atoms in a molecule. It represents the relative mass of a molecule expressed in atomic mass units (u).
How is molecular mass calculated?
- To find the molecular mass, add the atomic masses of each atom in the molecule, considering the number of each type of atom.
Examples of calculating molecular mass:
- Water ( $H 2 O):$
  - Atomic mass of hydrogen = 1 u, oxygen = 16 u.
  - Water contains 2 hydrogen atoms and 1 oxygen atom.
  - Molecular mass of $\text{H}_2\text{O} = (2 \times 1) + (1 \times 16) = 18$ u.
- Nitric Acid ( $\text{HNO}_3$ ):
  - Atomic mass of hydrogen = 1 u, nitrogen = 14 u, oxygen = 16 u.
  - Molecular mass of $\text{HNO}_3 = 1 + 14 + (3 \times 16) = 63$ u.

3.5.2 - Formula Unit Mass

What is formula unit mass?
- Formula unit mass is similar to molecular mass but is used for ionic compounds instead of molecules. It is the sum of the atomic masses of all atoms in the formula unit of an ionic compound.
How is formula unit mass calculated?
- Formula unit mass is calculated in the same way as molecular mass: by summing the atomic masses of all atoms in the formula unit.

Understanding molecular mass and formula unit mass is crucial for working with substances in chemistry, as these masses are fundamental for determining quantities in chemical reactions.

NCERT Science Class 9 | Science | Chapter 3 - Atoms and Molecules

NCERT Science Class 9 | Science | Chapter 3 - Atoms and Molecules

OUR SERVICES

Doubt Solving 1-on-1

Dedicated team provides prompt assistance and individual guidance.

NCERT Visualized

Engaging visuals enhance understanding of complex concepts.

Career Counselling

Engaging visuals enhance understanding of complex concepts.

Section-wise Tests

Assess understanding and track progress through topic-specific tests

NOTES

3.0 - Introduction

3.1 - Laws of Chemical Combination