Basic Concepts of Chemistry

CBSE English Medium - Willer Academy

Complete chapter notes in 3 structured lectures with expandable sections and interactive elements

Lecture 1: Matter & Measurement

Lecture 2: Atomic Structure

Lecture 3: Mole Concept

Table of Contents

• Lecture 1: Matter & Measurement
• Classification of Matter
• Properties of Matter
• Measurement & Units
• Lecture 2: Atomic Structure
• Laws of Chemical Combination
• Dalton's Atomic Theory
• Atomic & Molecular Mass
• Lecture 3: Mole Concept
• Mole & Avogadro's Number
• Stoichiometry
• Concentration Terms

1

Lecture 1: Matter & Measurement

Classification of Matter

Matter is anything that occupies space and has mass. It can be classified based on physical state or chemical composition.

Pure Substances: Elements and compounds with fixed composition and distinct properties.

Mixtures: Combination of two or more substances mixed physically.

Type	Characteristics	Examples
Element	Pure substance, one type of atom, cannot be broken down	Iron (Fe), Oxygen (O₂)
Compound	Pure substance, two or more elements chemically combined	Water (H₂O), Salt (NaCl)
Homogeneous Mixture	Uniform composition, single phase	Salt water, Air, Brass
Heterogeneous Mixture	Non-uniform composition, multiple phases	Sand and water, Salad, Granite

Exercise: Classify the following

1. Sugar dissolved in water
2. Iron filings and sulfur powder
3. Carbon dioxide
4. Milk

Properties of Matter

Physical Properties: Observed without changing composition

• Color, odor, density, melting point, boiling point, solubility
• Extensive (depends on quantity: mass, volume) vs Intensive (independent of quantity: density, melting point)

Chemical Properties: Observed during chemical reactions

• Flammability, reactivity, acidity/basicity, toxicity

States of Matter:

• Solid: Fixed shape and volume, particles closely packed
• Liquid: Fixed volume, takes container shape, particles can flow
• Gas: No fixed shape or volume, particles far apart
• Plasma: Ionized gas (not covered in basic chemistry)

Density = Mass / Volume

Measurement & Units

The International System of Units (SI) is used for scientific measurements:

Physical Quantity	SI Unit	Symbol
Mass	Kilogram	kg
Length	Meter	m
Time	Second	s
Temperature	Kelvin	K
Amount of substance	Mole	mol

Significant Figures: Meaningful digits in a measured quantity that indicate precision.

Rules:

• All non-zero digits are significant (e.g., 123 has 3 SF)
• Zeros between non-zero digits are significant (e.g., 1002 has 4 SF)
• Leading zeros are not significant (e.g., 0.0023 has 2 SF)
• Trailing zeros after decimal are significant (e.g., 2.300 has 4 SF)
• Exact numbers have infinite significant figures (e.g., counting numbers)

Exercise: Determine significant figures

1. 0.00450
2. 1200
3. 2.00 × 10³
4. 100.00

2

Lecture 2: Atomic Structure

Laws of Chemical Combination

Four fundamental laws govern chemical reactions:

1. Law of Conservation of Mass (Lavoisier):

"Mass is neither created nor destroyed in a chemical reaction."

Mass of reactants = Mass of products

2. Law of Definite Proportions (Proust):

"A chemical compound always contains the same elements in the same proportion by mass."

Example: Pure water always has H:O mass ratio of 1:8

3. Law of Multiple Proportions (Dalton):

"When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in a ratio of small whole numbers."

Example: Carbon forms CO and CO₂ with oxygen. Ratio of oxygen combining with fixed carbon is 1:2

4. Gay Lussac's Law of Gaseous Volumes:

"When gases combine, they do so in volumes that bear a simple ratio to one another and to the product (if gaseous), under the same conditions of temperature and pressure."

Example: 2H₂(g) + O₂(g) → 2H₂O(g) shows volume ratio 2:1:2

Dalton's Atomic Theory

John Dalton proposed the atomic theory in 1808:

• Matter consists of indivisible atoms
• All atoms of a given element are identical in mass and properties
• Atoms of different elements have different masses and properties
• Compounds form when atoms combine in fixed ratios
• Atoms are neither created nor destroyed in chemical reactions

Modifications to Dalton's Theory:

• Atoms are divisible (discovery of subatomic particles: electrons, protons, neutrons)
• Isotopes: Atoms of same element with different mass numbers
• Atoms can be transformed in nuclear reactions
• All atoms of an element are not identical (isotopes)

Exercise: Limitations of Dalton's Theory

Explain how the discovery of isotopes challenges Dalton's theory that all atoms of an element are identical.

Atomic & Molecular Mass

Atomic Mass: Mass of an atom relative to carbon-12 isotope

Atomic mass = (Mass of one atom of the element) / (1/12th mass of one C-12 atom)

Molecular Mass: Sum of atomic masses of all atoms in a molecule

Example: Molecular mass of H₂O = (2 × 1.008) + (1 × 16.00) = 18.016 u

Formula Mass: For ionic compounds that don't exist as discrete molecules

Example: Formula mass of NaCl = Atomic mass of Na + Atomic mass of Cl = 23.0 + 35.5 = 58.5 u

Element	Symbol	Atomic Mass (u)
Hydrogen	H	1.008
Carbon	C	12.01
Oxygen	O	16.00
Sodium	Na	23.0
Chlorine	Cl	35.5

3

Lecture 3: Mole Concept

Mole & Avogadro's Number

The mole is the SI unit for amount of substance that contains as many elementary entities as there are atoms in exactly 12g of carbon-12.

Avogadro's Number (Nₐ):

1 mole = 6.022 × 10²³ particles (atoms, molecules, ions)

Molar Mass: Mass of one mole of a substance in grams (numerically equal to atomic/molecular mass in u)

Example: Molar mass of H₂O = 18.016 g/mol

Concept	Formula	Example
Number of moles (n)	n = Mass / Molar mass	Moles in 36g H₂O: 36/18 = 2 moles
Number of particles	N = n × Nₐ	Molecules in 2 moles H₂O: 2 × 6.022×10²³
Mass from moles	Mass = n × Molar mass	Mass of 0.5 mole O₂: 0.5 × 32 = 16g

Exercise: Mole Calculations

1. How many moles are in 25g of calcium carbonate (CaCO₃)?
2. Calculate the number of oxygen atoms in 1g of oxygen gas (O₂).

Stoichiometry

Calculation of reactants and products in chemical reactions based on balanced equations.

Steps for Stoichiometric Calculations:

1. Write balanced chemical equation
2. Convert given quantities to moles
3. Use mole ratios from balanced equation
4. Convert moles to required units

Example: How many grams of O₂ are needed to burn 36g of carbon?

C + O₂ → CO₂

Mole ratio: 1 mol C : 1 mol O₂

Moles of C = 36g / 12g/mol = 3 moles

Moles of O₂ needed = 3 moles

Mass of O₂ = 3 × 32 = 96g

Limiting Reagent: The reactant that is completely consumed and determines the amount of product formed.

Excess Reagent: The reactant present in greater quantity than needed.

Concentration Terms

Ways to express concentration of solutions:

Term	Formula	Unit
Mass Percentage	(Mass of solute / Mass of solution) × 100	%
Mole Fraction (x)	n₁ / (n₁ + n₂ + ...)	Dimensionless
Molarity (M)	Moles of solute / Liters of solution	mol/L (M)
Molality (m)	Moles of solute / kg of solvent	mol/kg

Important Relationship:

Molarity (M) and density (d):

M = (Mass percentage × d × 10) / Molar mass of solute

Exercise: Concentration Calculations

1. Calculate the molarity of a solution containing 5g of NaOH in 250mL solution.
2. What is the molality of a solution with 20g glucose in 500g water?

Willer Academy Chemistry Notes | CBSE English Medium

Complete Chapter Notes in 3 Lectures | © 2023 Willer Academy

Basic Concepts of Chemistry

रसायन विज्ञान की मूल अवधारणाएँ

CBSE English & Hindi Medium - Willer Academy

Complete chapter notes in 3 structured lectures with bilingual content (English and Hindi)

Show Hindi Content

Lecture 1: Matter & Measurement

Lecture 2: Atomic Structure

Lecture 3: Mole Concept

Table of Contents

• Lecture 1: Matter & Measurement
• Classification of Matter
• Properties of Matter
• Measurement & Units
• Lecture 2: Atomic Structure
• Laws of Chemical Combination
• Dalton's Atomic Theory
• Atomic & Molecular Mass
• Lecture 3: Mole Concept
• Mole & Avogadro's Number
• Stoichiometry
• Concentration Terms

1

Lecture 1: Matter & Measurement

Classification of Matter

Matter is anything that occupies space and has mass. It can be classified based on physical state or chemical composition.

Pure Substances: Elements and compounds with fixed composition and distinct properties.

Mixtures: Combination of two or more substances mixed physically.

Type	Characteristics	Examples
Element	Pure substance, one type of atom, cannot be broken down	Iron (Fe), Oxygen (O₂)
Compound	Pure substance, two or more elements chemically combined	Water (H₂O), Salt (NaCl)
Homogeneous Mixture	Uniform composition, single phase	Salt water, Air, Brass
Heterogeneous Mixture	Non-uniform composition, multiple phases	Sand and water, Salad, Granite

पदार्थ का वर्गीकरण

पदार्थ कुछ भी है जो स्थान घेरता है और जिसमें द्रव्यमान होता है। इसे भौतिक अवस्था या रासायनिक संरचना के आधार पर वर्गीकृत किया जा सकता है।

• शुद्ध पदार्थ: तत्व और यौगिक जिनकी निश्चित संरचना और विशिष्ट गुण होते हैं
• मिश्रण: दो या दो से अधिक पदार्थों का भौतिक रूप से मिश्रण

उदाहरण:

• तत्व: लोहा (Fe), ऑक्सीजन (O₂)
• यौगिक: पानी (H₂O), नमक (NaCl)
• समांगी मिश्रण: नमकीन पानी, वायु, पीतल
• विषमांगी मिश्रण: रेत और पानी, सलाद, ग्रेनाइट

Exercise: Classify the following

1. Sugar dissolved in water
2. Iron filings and sulfur powder
3. Carbon dioxide
4. Milk

Properties of Matter

Physical Properties: Observed without changing composition

• Color, odor, density, melting point, boiling point, solubility
• Extensive (depends on quantity: mass, volume) vs Intensive (independent of quantity: density, melting point)

Chemical Properties: Observed during chemical reactions

• Flammability, reactivity, acidity/basicity, toxicity

States of Matter:

• Solid: Fixed shape and volume, particles closely packed
• Liquid: Fixed volume, takes container shape, particles can flow
• Gas: No fixed shape or volume, particles far apart
• Plasma: Ionized gas (not covered in basic chemistry)

Density = Mass / Volume

पदार्थ के गुण

भौतिक गुण: संरचना बदले बिना देखे जा सकने वाले गुण

• रंग, गंध, घनत्व, गलनांक, क्वथनांक, घुलनशीलता
• विस्तारी गुण (मात्रा पर निर्भर: द्रव्यमान, आयतन) vs गहन गुण (मात्रा से स्वतंत्र: घनत्व, गलनांक)

रासायनिक गुण: रासायनिक अभिक्रियाओं के दौरान देखे जाते हैं

• ज्वलनशीलता, अभिक्रियाशीलता, अम्लीयता/क्षारीयता, विषाक्तता

पदार्थ की अवस्थाएँ:

• ठोस: निश्चित आकार और आयतन, कण सघन रूप से व्यवस्थित
• द्रव: निश्चित आयतन, पात्र का आकार लेता है, कण प्रवाहित हो सकते हैं
• गैस: निश्चित आकार या आयतन नहीं, कण दूर होते हैं

घनत्व = द्रव्यमान / आयतन

Measurement & Units

The International System of Units (SI) is used for scientific measurements:

Physical Quantity	SI Unit	Symbol
Mass	Kilogram	kg
Length	Meter	m
Time	Second	s
Temperature	Kelvin	K
Amount of substance	Mole	mol

Significant Figures: Meaningful digits in a measured quantity that indicate precision.

Rules:

• All non-zero digits are significant (e.g., 123 has 3 SF)
• Zeros between non-zero digits are significant (e.g., 1002 has 4 SF)
• Leading zeros are not significant (e.g., 0.0023 has 2 SF)
• Trailing zeros after decimal are significant (e.g., 2.300 has 4 SF)
• Exact numbers have infinite significant figures (e.g., counting numbers)

मापन एवं मात्रक

वैज्ञानिक मापन के लिए अंतर्राष्ट्रीय मात्रक प्रणाली (SI) का उपयोग किया जाता है:

• द्रव्यमान: किलोग्राम (kg)
• लंबाई: मीटर (m)
• समय: सेकंड (s)
• तापमान: केल्विन (K)
• पदार्थ की मात्रा: मोल (mol)

सार्थक अंक: मापी गई मात्रा में सार्थक अंक जो परिशुद्धता को दर्शाते हैं।

नियम:

• सभी शून्येतर अंक सार्थक होते हैं (जैसे 123 में 3 सार्थक अंक)
• शून्येतर अंकों के बीच के शून्य सार्थक होते हैं (जैसे 1002 में 4 सार्थक अंक)
• अग्र शून्य सार्थक नहीं होते (जैसे 0.0023 में 2 सार्थक अंक)
• दशमलव के बाद के अनुगामी शून्य सार्थक होते हैं (जैसे 2.300 में 4 सार्थक अंक)
• निश्चित संख्याओं में अनंत सार्थक अंक होते हैं (जैसे गिनती की संख्याएँ)

Exercise: Determine significant figures

1. 0.00450
2. 1200
3. 2.00 × 10³
4. 100.00

2

Lecture 2: Atomic Structure

Laws of Chemical Combination

Four fundamental laws govern chemical reactions:

1. Law of Conservation of Mass (Lavoisier):

"Mass is neither created nor destroyed in a chemical reaction."

Mass of reactants = Mass of products

2. Law of Definite Proportions (Proust):

"A chemical compound always contains the same elements in the same proportion by mass."

Example: Pure water always has H:O mass ratio of 1:8

3. Law of Multiple Proportions (Dalton):

"When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in a ratio of small whole numbers."

Example: Carbon forms CO and CO₂ with oxygen. Ratio of oxygen combining with fixed carbon is 1:2

रासायनिक संयोग के नियम

1. द्रव्यमान संरक्षण का नियम (लावॉज़ियर):

"रासायनिक अभिक्रिया में द्रव्यमान न तो उत्पन्न होता है और न ही नष्ट होता है।"

अभिकारकों का द्रव्यमान = उत्पादों का द्रव्यमान

2. निश्चित अनुपात का नियम (प्राउस्ट):

"एक रासायनिक यौगिक में तत्व सदैव द्रव्यमान के अनुसार समान अनुपात में उपस्थित होते हैं।"

उदाहरण: शुद्ध जल में हाइड्रोजन : ऑक्सीजन का द्रव्यमान अनुपात सदैव 1:8 होता है

3. गुणित अनुपात का नियम (डाल्टन):

"जब दो तत्व एक से अधिक यौगिक बनाते हैं, तो एक तत्व का द्रव्यमान जो दूसरे तत्व के निश्चित द्रव्यमान से संयोग करता है, छोटे पूर्ण संख्याओं के अनुपात में होता है।"

उदाहरण: कार्बन ऑक्सीजन के साथ CO और CO₂ बनाता है। निश्चित कार्बन के साथ संयुक्त ऑक्सीजन का अनुपात 1:2 है

Dalton's Atomic Theory

John Dalton proposed the atomic theory in 1808:

• Matter consists of indivisible atoms
• All atoms of a given element are identical in mass and properties
• Atoms of different elements have different masses and properties
• Compounds form when atoms combine in fixed ratios
• Atoms are neither created nor destroyed in chemical reactions

Modifications to Dalton's Theory:

• Atoms are divisible (discovery of subatomic particles: electrons, protons, neutrons)
• Isotopes: Atoms of same element with different mass numbers
• Atoms can be transformed in nuclear reactions
• All atoms of an element are not identical (isotopes)

डाल्टन का परमाणु सिद्धांत

जॉन डाल्टन ने 1808 में परमाणु सिद्धांत प्रस्तावित किया:

• पदार्थ अविभाज्य परमाणुओं से मिलकर बना होता है
• किसी तत्व के सभी परमाणु द्रव्यमान और गुणों में समान होते हैं
• भिन्न तत्वों के परमाणुओं के भिन्न द्रव्यमान और गुण होते हैं
• यौगिक तब बनते हैं जब परमाणु निश्चित अनुपात में संयोग करते हैं
• रासायनिक अभिक्रियाओं में परमाणु न तो बनते हैं और न नष्ट होते हैं

डाल्टन के सिद्धांत में संशोधन:

• परमाणु विभाज्य हैं (उपपरमाणुक कणों की खोज: इलेक्ट्रॉन, प्रोटॉन, न्यूट्रॉन)
• समस्थानिक: एक ही तत्व के परमाणु जिनके द्रव्यमान संख्या भिन्न होती है
• परमाणु नाभिकीय अभिक्रियाओं में परिवर्तित हो सकते हैं
• एक तत्व के सभी परमाणु समान नहीं होते (समस्थानिक)

Exercise: Limitations of Dalton's Theory

Explain how the discovery of isotopes challenges Dalton's theory that all atoms of an element are identical.

Atomic & Molecular Mass

Atomic Mass: Mass of an atom relative to carbon-12 isotope

Atomic mass = (Mass of one atom of the element) / (1/12th mass of one C-12 atom)

Molecular Mass: Sum of atomic masses of all atoms in a molecule

Example: Molecular mass of H₂O = (2 × 1.008) + (1 × 16.00) = 18.016 u

Formula Mass: For ionic compounds that don't exist as discrete molecules

Example: Formula mass of NaCl = Atomic mass of Na + Atomic mass of Cl = 23.0 + 35.5 = 58.5 u

परमाणु एवं आणविक द्रव्यमान

परमाणु द्रव्यमान: कार्बन-12 समस्थानिक के सापेक्ष परमाणु का द्रव्यमान

परमाणु द्रव्यमान = (तत्व के एक परमाणु का द्रव्यमान) / (एक C-12 परमाणु के द्रव्यमान का 1/12)

आणविक द्रव्यमान: एक अणु में सभी परमाणुओं के परमाणु द्रव्यमानों का योग

उदाहरण: H₂O का आणविक द्रव्यमान = (2 × 1.008) + (1 × 16.00) = 18.016 u

सूत्र द्रव्यमान: आयनिक यौगिकों के लिए जो अलग-अलग अणुओं के रूप में नहीं होते

उदाहरण: NaCl का सूत्र द्रव्यमान = Na का परमाणु द्रव्यमान + Cl का परमाणु द्रव्यमान = 23.0 + 35.5 = 58.5 u

Element	Symbol	Atomic Mass (u)
Hydrogen	H	1.008
Carbon	C	12.01
Oxygen	O	16.00
Sodium	Na	23.0
Chlorine	Cl	35.5

3

Lecture 3: Mole Concept

Mole & Avogadro's Number

The mole is the SI unit for amount of substance that contains as many elementary entities as there are atoms in exactly 12g of carbon-12.

Avogadro's Number (Nₐ):

1 mole = 6.022 × 10²³ particles (atoms, molecules, ions)

Molar Mass: Mass of one mole of a substance in grams (numerically equal to atomic/molecular mass in u)

Example: Molar mass of H₂O = 18.016 g/mol

मोल एवं आवोगाद्रो संख्या

मोल पदार्थ की मात्रा की SI इकाई है जिसमें उतने ही मूल इकाइयाँ होती हैं जितने परमाणु ठीक 12g कार्बन-12 में होते हैं।

आवोगाद्रो संख्या (Nₐ):

1 मोल = 6.022 × 10²³ कण (परमाणु, अणु, आयन)

मोलर द्रव्यमान: किसी पदार्थ के एक मोल का ग्राम में द्रव्यमान (संख्यात्मक रूप से u में परमाणु/आणविक द्रव्यमान के बराबर)

उदाहरण: H₂O का मोलर द्रव्यमान = 18.016 g/mol

Concept	Formula	Example
Number of moles (n)	n = Mass / Molar mass	Moles in 36g H₂O: 36/18 = 2 moles
Number of particles	N = n × Nₐ	Molecules in 2 moles H₂O: 2 × 6.022×10²³
Mass from moles	Mass = n × Molar mass	Mass of 0.5 mole O₂: 0.5 × 32 = 16g

Exercise: Mole Calculations

1. How many moles are in 25g of calcium carbonate (CaCO₃)?
2. Calculate the number of oxygen atoms in 1g of oxygen gas (O₂).

Stoichiometry

Calculation of reactants and products in chemical reactions based on balanced equations.

Steps for Stoichiometric Calculations:

1. Write balanced chemical equation
2. Convert given quantities to moles
3. Use mole ratios from balanced equation
4. Convert moles to required units

Example: How many grams of O₂ are needed to burn 36g of carbon?

C + O₂ → CO₂

Mole ratio: 1 mol C : 1 mol O₂

Moles of C = 36g / 12g/mol = 3 moles

Moles of O₂ needed = 3 moles

Mass of O₂ = 3 × 32 = 96g

रससमीकरणमिति (स्टॉइकियोमेट्री)

संतुलित समीकरणों के आधार पर रासायनिक अभिक्रियाओं में अभिकारकों और उत्पादों की गणना।

रससमीकरणमितीय गणना के चरण:

1. संतुलित रासायनिक समीकरण लिखें
2. दी गई मात्राओं को मोल में बदलें
3. संतुलित समीकरण से मोल अनुपात का उपयोग करें
4. मोल को आवश्यक इकाइयों में बदलें

उदाहरण: 36g कार्बन को जलाने के लिए कितने ग्राम O₂ की आवश्यकता होगी?

C + O₂ → CO₂

मोल अनुपात: 1 मोल C : 1 मोल O₂

C के मोल = 36g / 12g/mol = 3 मोल

आवश्यक O₂ के मोल = 3 मोल

O₂ का द्रव्यमान = 3 × 32 = 96g

Concentration Terms

Ways to express concentration of solutions:

Term	Formula	Unit
Mass Percentage	(Mass of solute / Mass of solution) × 100	%
Mole Fraction (x)	n₁ / (n₁ + n₂ + ...)	Dimensionless
Molarity (M)	Moles of solute / Liters of solution	mol/L (M)
Molality (m)	Moles of solute / kg of solvent	mol/kg

सांद्रता पद

विलयन की सांद्रता व्यक्त करने के तरीके:

• द्रव्यमान प्रतिशत: (विलेय का द्रव्यमान / विलयन का द्रव्यमान) × 100
• मोल प्रभाज (x): n₁ / (n₁ + n₂ + ...)
• मोलरता (M): विलेय के मोल / विलयन के लीटर
• मोललता (m): विलेय के मोल / विलायक के kg

Exercise: Concentration Calculations

1. Calculate the molarity of a solution containing 5g of NaOH in 250mL solution.
2. What is the molality of a solution with 20g glucose in 500g water?

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