NCERT Notes Class 11 Chemistry (Part-I) Chapter 1: Some Basic Concepts of Chemistry (Free PDF)

Chemistry is an important branch of natural science that deals with the study of matter, energy, and the interactions between them. In this unit on some basic concepts of chemistry, we discuss the importance and development of chemistry, its branches, the nature and scope of the subject, the states and classification of matter, as well as key concepts such as mass, volume, density, atomic and molecular masses, mole concept, and stoichiometry.

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Introduction

Science aims to systematically understand and describe nature. Chemistry is the branch of science dealing with the composition, structure, properties, and reactions of material substances. Everyday changes such as curd formation, vinegar formation, and rusting of iron are chemical in nature. Chemistry evolved through practices like alchemy, iatrochemistry, and later modern chemistry.

Development of Chemistry

In this section, we have presented a sneak peek at the historical development of chemistry.

Alchemical Traditions in India

The alchemical traditions in India are briefly mentioned below.

• Ancient India had well-developed chemical knowledge, including:
  • Metallurgy, medicine, dyeing, glass and ceramic production, cosmetics, etc.
  • Found in texts like Rasayan Shastra, Rasvidya, and Rastantra.
• Excavations at Mohenjo-Daro and Harappa show the use of:
  • Baked bricks, glazed pottery, and gypsum cement.
  • Faience: a glass-like material used for ornaments.

Contributions to Chemistry

Some of the notable Indian teachers who contributed to the field of Chemistry in India are briefly mentioned below.

• Charaka Samhita: Describes sulphuric acid, nitric acid, and metal oxides.
• Sushruta Samhita: Highlights the use of alkalies in treatment.
• Kautilya’s Arthashastra: Discusses salt production, dyeing, and fermentation.
• Bhasma: Finely powdered metal used in Ayurvedic treatments (now considered to contain nanoparticles).
• Acharya Kanda (600 BCE): Proposed atomic theory using Paramāṇu — indivisible, eternal, in motion.
• Nagarjuna: Described mercury compounds and metal extraction.
• Chakrapani: Discovered mercury sulphide (HgS), and made early soap using alkalies and mustard oil.

Importance of Chemistry

• Chemistry is central to all sciences and is used in:
  • Weather prediction
  • Functioning of computers
  • Brain processes
• Chemistry helps in the production of:
  • Fertilisers, dyes, soaps, drugs, detergents, alloys, metals, and synthetic materials.
  • Healthcare: Drugs like cisplatin (for cancer) and AZT (for AIDS) were developed using chemical processes.
  • Environment: Safer CFC substitutes have been created using chemical knowledge.
• Chemistry contributes to the economy and employment generation through industries.

Nature of Matter

Matter is anything that has mass and occupies space. It exists in 3 forms mainly: solid, liquid, and gas. Below, we have tabulated the general differences between these matters.

State	Volume	Shape	Particle Arrangement
Solid	Definite	Definite	Closely packed
Liquid	Definite	Not Definite	Loosely packed
Gas	Not Definite	Not Definite	Widely spaced

Interconversion: Matter can be converted from one state to another by changing temperature or pressure

Classification of Matter

Below, we have provided a tabular description of the classification of matter as a mixture.

Type	Details	Examples
Mixture	Contains 2+ substances mixed physically	Air, tea, salt water
Homegeneous	Uniform composition	Sugar solution
Hetreogeneous	Non-Uniform composition	Salt and sugar, oil and water

Below, we have provided a tabular description of the classification of matter as pure substance.

Type	Details	Examples
Pure Substance	Same type of particles with fixed composition	Water, copper, glucose
Elements	Made of identical atoms	O₂, Na, H₂
Compounds	Atoms of different elements in a fixed ratio; chemically combined	H₂O, CO₂, NH₃

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Properties of Matter & their Measurement

Every substance has unique characteristics or properties. These properties can be classified into two categories — physical properties, such as colour, odour, melting point, boiling point, density, etc., and chemical properties, like composition, combustibility, reactivity with acids and bases, etc. Physical properties can be measured or observed without changing the identity or the composition of the substance. The measurement or observation of chemical properties requires a chemical change to occur. The examples of chemical properties are characteristic reactions of different substances; these include acidity or basicity, combustibility, etc.

Measurement of Physical Properties

Scientific investigation requires quantitative measurements. Quantities are expressed as a number followed by its unit. Example:

• Length = 6 m (where 6 is the number, and m is the unit of metre)

Earlier, two systems of measurement were used:

• English system
• Metric system

The metric system was preferred because it’s based on the decimal system. In 1960, an internationally accepted standard system of units was established, called the SI system (International System of Units).

The International System of Units (SI)

Based on 7 fundamental physical quantities, each with its own unit.

Physical Quantity	Symbol	SI Unit	Expressed as
Length	l	metre	m
Mass	m	kilogram	kg
Time	t	second	s
Electric Current	I	ampere	A
Thermodynamic Temperature	T	kelvin	K
Amount of Substance	n	mole	mol
Luminous Intensity	lv	candela	cd

In India, the National Physical Laboratory (NPL), New Delhi, maintains the national standards of measurements. Units are redefined when newer, more accurate methods are developed.

Mass and Weight

Mass and weight are discussed briefly below.

Mass:

• Mass is defined as the amount of matter present in a substance.
• It is an intrinsic property of matter.
• The mass of a substance remains constant and does not vary with location or gravitational pull.
• It is measured using an analytical balance in laboratories.
• The SI unit of mass is the kilogram (kg).
• In laboratories, smaller units like grams (g) are commonly used.

Weight:

• Weight is the force with which an object is attracted towards the Earth.

• It depends on mass and gravitational acceleration (g).
  • Weight = Mass × g
• Unlike mass, weight is not constant. It varies from place to place depending on the gravity.
  • Example: A body weighs less on the moon than on Earth.
• The SI unit of weight is also newton (N) (being a force), but it is commonly expressed in kg-force in everyday life.

Volume

The volume is briefly described below.

• Volume refers to the space occupied by matter.
• It is an extensive property and varies with the amount of substance.
• SI unit of volume: cubic metre (m³)
• Commonly used units in the laboratory:
  • 1 litre (L) = 1000 millilitres (mL)
  • 1 L = 1000 cm³ = 1 dm³
• Instruments used to measure volume:
  • Graduated cylinders
  • Pippttes
  • Burettes
  • Volumetric flasks

Density

The density is briefly described below.

• Density is defined as the mass per unit volume of a substance.
• It is represented by the formula: Density = Mass / Volume
• SI unit: kg/m³
• Common unit in laboratories: g/cm³ or g/mL
• It helps in identifying substances and determining purity.

Temperature

• Temperature is a measure of the degree of hotness or coldness of a body.
• It determines the direction of heat flow.
• Heat flows from a body at a higher temperature to a lower temperature.
• Common temperature scales:
  1. Celsius scale (°C)
  2. Fahrenheit scale (°F)
  3. Kelvin scale (K) — the SI unit of temperature

Conversion formulas:

• K = °C + 273.15
• °F = (°C × 9/5) + 32
• The Kelvin scale begins at absolute zero (0 K), the lowest possible temperature where the motion of particles ceases.
• It does not have negative values.

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Uncertainty in Measurement

In scientific experiments, uncertainty may arise due to limitations of instruments and human errors.

Scientific Notation

It is used to express very large or very small numbers in compact form. General form:
N × 10ⁿ, where 1 ≤ N < 10 and n is an integer.

Examples:

• 0.00016 = 1.6 × 10⁻⁴
• 232.508 = 2.32508 × 10²
• 6,02,200,000,000,000,000,000,000 = 6.022 × 10²³ (Avogadro’s number)
• Useful for representing:
  • Atomic and molecular sizes
  • Number of particles in a mole

Significant Figures

The number of meaningful digits in a measured quantity is called significant figures. These include all digits that are known reliably, plus one uncertain digit.

Rules for determining significant figures:

1. All non-zero digits are significant.
   • Example: 237 → 3 significant figures
2. Zeros between non-zero digits are significant.
   • Example: 1005 → 4 significant figures
3. Leading zeros are not significant.
   • Example: 0.0025 → 2 significant figures
4. Trailing zeros in a number with a decimal point are significant.
   • Example: 3.00 → 3 significant figures
5. Exact numbers (e.g., 2 pens, 12 eggs) have infinite significant figures.

Rounding off significant figures:

• If the digit to be dropped is less than 5, the preceding digit is not changed.
• If the digit to be dropped is more than 5, the preceding digit is increased by one.
• If the digit to be dropped is exactly 5, round to make the preceding digit even.

Rules for operations:

• Addition/Subtraction:
  • Final answer should have the same number of decimal places as the number with the fewest decimal places.
• Multiplication/Division:
  • Final result should have the same number of significant figures as the quantity with the least number of significant figures.

Dimensional Analysis

It is also called the factor label method or unit factor method. Used to convert units from one system to another using conversion factors.

Examples:

1. Convert 3 inches to cm:
   1 inch = 2.54 cm
   → 3 × 2.54 = 7.62 cm
   
2. Convert 2 L to m³:
   1 L = 1 dm³ = 10⁻³ m³
   → 2 L = 2 × 10⁻³ m³
   
3. Convert 2 days into seconds:
   1 day = 24 × 60 × 60 = 86400 seconds
   → 2 days = 172800 seconds

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Laws of Chemical Combinations

These laws were developed after experimental observations regarding how elements combine to form compounds.

Law of Conservation of Mass

It was proposed by Antoine Lavoisier in 1789. It states that:

“Mass can neither be created nor destroyed during a chemical reaction.”

Example: When 10 g of calcium reacts with 17 g of chlorine, 27 g of calcium chloride is formed.

Law of Definite Proportions

It was proposed by Joseph Proust. It states that:

“A given compound always contains the same elements in the same proportion by mass.”

Example: Water from any source contains 88.89% oxygen and 11.11% hydrogen by mass.

Law of Multiple Proportions

It was proposed by John Dalton in 1803. It states that:

“If two elements combine to form two or more compounds, the masses of one element that combine with a fixed mass of the other are in the ratio of small whole numbers.”

Example: CO and CO₂:

• 12 g of carbon combines with 16 g (in CO) and 32 g (in CO₂) of oxygen
• Ratio = 16:32 = 1:2

Gay Lussac’s Law of Gaseous Volumes

It was proposed by Gay Lussac in 1808. It states that:

“When gases combine or are produced in a chemical reaction, they do so in simple whole number ratios by volume, provided all gases are measured at the same temperature and pressure.”

Example: H₂ + O₂ → H₂O
2 volumes of hydrogen + 1 volume of oxygen → 2 volumes of water vapor

Avogadro’s Law

It was proposed by Amedeo Avogadro in 1811. It states that:

“Equal volumes of all gases, at the same temperature and pressure, contain equal number of molecules.”

Avogadro’s number (Nₐ): 6.022 × 10²³ particles/mol

Dalton’s Atomic Theory

Dalton’s Atomic Theory was proposed by John Dalton in 1808.

Postulates of Dalton’s Atomic Theory:

1. Matter is made up of tiny indivisible particles called atoms.
2. Atoms of the same element are identical in all respects.
3. Atoms of different elements have different properties and masses.
4. Compounds are formed when atoms combine in simple whole-number ratios.
5. Atoms can neither be created nor destroyed in chemical reactions.

Atomic and Molecular Masses

In this section, we have provided a summary of the concept of atomic and molecular masses.

Atomic Mass

Atoms are extremely small, so their masses are expressed relative to a standard. The standard used is the carbon-12 isotope. 1 atomic mass unit (1 amu or 1 u) = 1/12 the mass of one atom of carbon-12.

• In grams:
  1 u = 1.66056 × 10⁻²⁴ g
• Example:
  • Mass of one hydrogen atom = 1.6736 × 10⁻²⁴ g
    → in atomic mass units = 1.0078 u

Average Atomic Mass

In our universe, many elements exist as isotopes. The atomic mass of such elements is the weighted average of the masses of the isotopes based on their natural abundance.

Formula:
Average atomic mass = Σ (isotopic mass × % abundance) / 100

Example – Chlorine:

• ³⁵Cl (75.77%, mass = 34.9689 u)
• ³⁷Cl (24.23%, mass = 36.9659 u)

Average atomic mass =
(75.77 × 34.9689 + 24.23 × 36.9659)/100 = 35.5 u

Molecular Mass

The molecular mass is the sum of the atomic masses of all atoms present in a molecule.

Formula:
Molecular mass = Σ (atomic masses of constituent atoms)

Example – CH₄ (Methane):

• C = 12.01 u, H = 1.008 u
• Molecular mass = 12.01 + 4 × 1.008 = 16.043 u

Formula Mass

For ionic compounds (which do not exist as discrete molecules), we use the term formula mass instead of molecular mass. It is calculated the same way: by adding the atomic masses of the ions in the empirical formula.

Example – NaCl:

• Na = 22.99 u, Cl = 35.45 u
• Formula mass = 22.99 + 35.45 = 58.44 u

Mole Concept and Molar Masses

The mole concept and molar masses are discussed below.

Mole Concept

1 mole is defined as the amount of a substance that contains as many entities (atoms, molecules, ions, etc.) as there are in 12 g of carbon-12. 1 mole = 6.022 × 10²³ particles (Avogadro’s number, Nₐ). So:

• 1 mole of atoms = 6.022 × 10²³ atoms
• 1 mole of molecules = 6.022 × 10²³ molecules
• 1 mole of electrons = 6.022 × 10²³ electrons

Molar Mass

Molar mass = Mass of one mole of a substance. It is numerically equal to the atomic/molecular/formula mass, but expressed in grams per mole (g/mol)

Examples:

• H₂O: Molecular mass = 18.02 u → Molar mass = 18.02 g/mol
• NaCl: Formula mass = 58.44 u → Molar mass = 58.44 g/mol

Percentage Composition

It refers to the mass per cent of each element present in a compound.

Formula:
Mass % of element = (Mass of element in 1 mol of compound / Molar mass of compound) × 100

Example – H₂O

• Molar mass = 2 × 1.008 + 16.00 = 18.016 g/mol
• Mass % of H = (2.016 / 18.016) × 100 = 11.19%
• Mass % of O = (16.00 / 18.016) × 100 = 88.81%

Example – C₂H₆ (Ethane)

• C = 12.01, H = 1.008
• Molar mass = 2 × 12.01 + 6 × 1.008 = 30.07 g/mol
• Mass % of C = (24.02 / 30.07) × 100 = 79.89%
• Mass % of H = (6.048 / 30.07) × 100 = 20.11%

Empirical Formula and Molecular Formula

In this section, we have discussed the empirical formula and molecular formula.

Empirical Formula

The simplest whole-number ratio of atoms of each element in a compound.

Molecular Formula

It represents the actual number of atoms of each element present in one molecule. It is always a whole-number multiple of the empirical formula.

Formula:
Molecular formula = (Empirical formula) × n

Where,
n = Molar mass / Empirical formula mass

Steps to Determine Empirical Formula:

1. Assume 100 g of the compound (mass % = grams).
2. Convert masses to moles using atomic masses.
3. Divide each mole value by the smallest mole number.
4. Multiply to get a whole-number ratio (if needed).
5. Write the empirical formula using these ratios.

Steps to Determine Molecular Formula:

Use the formula:
Molecular formula = Empirical formula × (Molar mass / Empirical formula mass)

Stoichiometry and Stoichiometric Calculations

Stoichiometry deals with quantitative relationships between reactants and products in a balanced chemical reaction.

Example Reaction:

CH₄ + 2O₂ → CO₂ + 2H₂O

This means:

• 1 mole of CH₄ reacts with 2 moles of O₂
• Produces 1 mole of CO₂ and 2 moles of H₂O

So, if we have:

• 16 g CH₄ reacts with 64 g O₂ → gives 44 g CO₂ + 36 g H₂O

Limiting Reagent

In a chemical reaction, the limiting reagent is the reactant that is completely consumed first, limiting the amount of product formed. The other reactants are in excess.

Steps to Identify the Limiting Reagent:

1. Calculate the moles of all reactants.
2. Use mole ratios from the balanced equation to determine which reactant would produce the least amount of product.
3. The reactant that forms the least amount of product is the limiting reagent.

Class 11 Chemistry (Part-I) Chapter 1: Some Basic Concepts of Chemistry- Important Formulas 

Here are the important formulas and equations from NCERT Class 11 Chemistry Chapter Some Basic Concepts of Chemistry.

• Number of atoms in 1 mole (Avogadro’s number)
  N = 6.022 × 10²³
• Molar mass
  Molar mass = Mass of 1 mole of a substance (in grams)
• Mole formula
  Number of moles = Given mass / Molar mass
• Number of particles (atoms, molecules, ions, etc.)
  Number of particles = Moles × Avogadro’s number
• Mass from moles
  Mass = Number of moles × Molar mass
• Volume of gas at STP
  1 mole of gas at STP = 22.4 L
  Volume = Moles × 22.4 L (at STP)
• Percentage composition
  % of element = (Mass of element in 1 mole / Molar mass of compound) × 100
• Empirical formula calculation
  • Convert % composition to grams
  • Divide by atomic masses to get moles
  • Divide all mole values by the smallest mole to get the simplest ratio
• Molecular formula
  Molecular formula = (Empirical formula) × (Molar mass / Empirical formula mass)
• Limiting reagent
  The reagent that produces the lesser amount of product is the limiting reagent.
• Concentration in molarity
  Molarity (M) = Moles of solute / Volume of solution in litres
• Law of Conservation of Mass
  Total mass of reactants = Total mass of products
• Law of Definite Proportions
  A given compound always contains the same elements in the same proportion by mass.

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