Chapter 1: Some Basic Concepts of Chemistry – Class 11 Notes for NEET/JEE

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1. Importance of Chemistry

🔹 What is Chemistry?

Chemistry is the branch of science that studies:

• Matter (anything that has mass and takes up space),
• Its composition (what it’s made of),
• Its properties (how it behaves), and
• The changes it undergoes (chemical reactions).

🔍 Why is Chemistry Important?

1. Everyday Life

Chemistry is everywhere — from the air we breathe to the food we eat. It explains:

• Why food spoils
• How soap cleans
• How medicines cure diseases
• Why milk turns sour

2. In Medicine

Chemistry helps in:

• Developing drugs and vaccines
• Understanding how medicines work in the body
• Making life-saving antibiotics (like penicillin)

Example: Paracetamol is made using organic chemistry reactions.

3. In Agriculture

• Chemistry creates fertilizers, pesticides, and insecticides that improve crop yield.
• Helps solve food shortage problems.

4. In Environment

• Helps understand pollution, climate change, ozone depletion, and greenhouse effect.
• Chemistry provides eco-friendly alternatives like biodegradable plastics.

5. In Industry

Used in:

• Manufacturing of textiles, cement, paint, plastic, glass
• Fuel production (petrol, diesel, LPG)
• Cosmetics and detergents

6. In Technology

• Battery and mobile technology use chemistry principles.
• Semiconductor chemistry powers your phone, computer, and internet.

🎯 Exam Relevance

• This topic helps you understand how Chemistry connects to biology, physics, and environment, which is important for NEET Biology applications, especially in:
  • Human body functions
  • Drug action mechanisms
  • Cellular reactions (enzymes, ATP production)

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2. 🧪 Nature of Matter

🔹 What is Matter?

Matter is anything that has mass and occupies space.

Examples:

• Air, water, soil, plants, animals, food, clothes — all are matter.

NEET me directly ya indirectly yeh concept base banata hai states of matter, physical vs chemical properties, aur composition of substances ke liye.

🔹 Physical States of Matter

Matter generally exists in three main physical states:

State	Properties
Solid	Fixed shape & volume, particles tightly packed, not compressible
Liquid	Fixed volume but no fixed shape, takes shape of container
Gas	Neither fixed shape nor volume, compressible, particles far apart

Exam Tip: Gaseous state is very important in physical chemistry. You’ll study gas laws in detail in Chapter: “States of Matter”.

🔹 Classification Based on Composition

1. Pure Substances

• Same type of particles throughout.
• Fixed composition and properties.
• Can be elements or compounds.

a. Elements

• Made of only one type of atom.
• Cannot be broken down into simpler substances.
• E.g., Hydrogen (H), Oxygen (O), Iron (Fe)

b. Compounds

• Made of two or more elements chemically combined.
• Fixed ratio.
• Can be broken into simpler substances using chemical methods.
• E.g., Water (H₂O), Carbon dioxide (CO₂)

Exam Tip: Understanding compounds helps in inorganic and organic chemistry. You’ll study ionic and covalent bonding later.

2. Mixtures

• Two or more substances physically mixed.
• Variable composition.
• Properties depend on individual components.
• Can be homogeneous or heterogeneous.

a. Homogeneous Mixture (Solution)

• Uniform composition throughout.
• Example: Salt in water, air.

b. Heterogeneous Mixture

• Non-uniform composition.
• Example: Oil in water, sand in iron.

Exam Relevance: You may get application-based questions like:
Which of the following is a homogeneous mixture?
Answer: Air

🔹 Difference Between Mixture and Compound

Property	Mixture	Compound
Composition	Variable	Fixed
Separation method	Physical methods	Only chemical methods
Properties	Same as components	Different from individual elements
Energy change	No significant energy change	Heat is often absorbed or released

🔹 Types of Properties of Matter

1. Physical Properties

• Can be measured without changing chemical identity.
• Examples: Color, boiling point, melting point, density.

2. Chemical Properties

• Observed during chemical reaction.
• Change the composition of substance.
• Examples: Reactivity with acid, flammability, rusting.

Exam Link: In chapters like Thermodynamics & Redox Reactions, physical vs chemical properties are important for concept clarity.

📌 Summary Chart: Nature of Matter

Category	Sub-category
Physical State	Solid, Liquid, Gas
Composition	Pure Substances (Elements, Compounds)
	Mixtures (Homogeneous, Heterogeneous)
Properties	Physical and Chemical

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3. Properties of Matter and Their Measurement

🔹 1. Physical Properties

🔍 Definition:

Physical properties are the characteristics of matter that can be observed or measured without changing its chemical identity.

🧾 Examples:

• Color
• Odor
• Melting point
• Boiling point
• Density
• Solubility
• Hardness
• State (solid/liquid/gas)

Exam Relevance: Questions may ask to identify physical vs chemical property.

🔹 2. Chemical Properties

🔍 Definition:

Chemical properties describe the ability of a substance to undergo a chemical change and form new substances.

🧾 Examples:

• Rusting of iron
• Burning of fuel
• Reactivity with acid or base
• Flammability
• Decomposition

Exam Example Question:
Which of the following is a chemical property?
A. Boiling point
B. Solubility
C. Reactivity with HCl
✅ Answer: C. Reactivity with HCl

🔹 3. Units of Measurement (SI Units)

🔍 What is a Unit?

A unit is a standard for measuring a physical quantity.

📐 Base SI Units:

Physical Quantity	SI Unit	Symbol
Length	metre	m
Mass	kilogram	kg
Time	second	s
Temperature	kelvin	K
Amount of substance	mole	mol
Electric current	ampere	A
Luminous intensity	candela	cd

Exam Tip: Units of length (m), mass (kg), time (s), and amount of substance (mol) are frequently used in NEET Chemistry & Physics numericals.

🔹 4. Uncertainty in Measurement

🔍 What is Uncertainty?

Measurement is never perfectly accurate. Uncertainty refers to the doubt or error associated with a measured value.

🔍 Types of Errors:

• Systematic Error: Repeated error due to instrument fault or wrong method.
• Random Error: Error due to sudden changes in environment (like wind, temperature).
• Human Error: Mistakes made by observer (like wrong reading).

🔍 Accuracy vs Precision:

• Accuracy = How close measurement is to true value.
• Precision = How close repeated measurements are to each other.

Exam Tip: Concept clarity in error types helps in solving numericals in later chapters like Thermodynamics, Equilibrium, etc.

🔹 5. Significant Figures

🔍 What are Significant Figures?

Digits in a number that carry meaningful information about its accuracy.

✅ Rules to Identify:

1. All non-zero digits are significant.
   → 123 = 3 significant figures
2. Zeros between non-zero digits are significant.
   → 1003 = 4 significant figures
3. Leading zeros are not significant.
   → 0.0032 = 2 significant figures
4. Trailing zeros after decimal are significant.
   → 2.500 = 4 significant figures
5. Exact numbers (like 100 students) have infinite significant figures.

🔍 Rounding Off:

If digit to be dropped is:

• <5 → previous digit stays same
• ≥5 → increase previous digit by 1

Exam Tip: Significant figures help in correctly reporting answers in numerical problems.

🔹 6. Dimensional Analysis

🔍 What is Dimensional Analysis?

A method to convert units from one system to another using conversion factors.

🧾 Example:

Convert 1 meter = ? inches
We know: 1 inch = 2.54 cm
So,
1 meter = 100 cm = 100 ÷ 2.54 = 39.37 inches

🔁 Formula Conversion Using Dimensional Analysis:

If
Speed = Distance / Time​
Then dimension of speed = [L] / [T] = [LT⁻¹]

Exam Tip: Dimensional analysis helps verify the correctness of equations and convert units (e.g., g to kg, L to mL, etc.)

📌 Quick Summary Table

Topic	Key Points
Physical Properties	Observed without chemical change (e.g., color, boiling point)
Chemical Properties	Observed during a chemical change (e.g., rusting, reactivity)
SI Units	Standard international units (m, kg, mol, K, etc.)
Uncertainty in Measurement	Every measurement has error; understand accuracy vs precision
Significant Figures	Digits that show precision in measurement
Dimensional Analysis	Used to convert units or verify equations

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4⚗️ Laws of Chemical Combination – NEET Detailed Explanation.

1. Law of Conservation of Mass

What does it mean?
In a chemical reaction, the total mass of the substances before the reaction (called reactants) is exactly equal to the total mass of substances after the reaction (called products). No mass is lost or gained.

Why?
Atoms are just rearranged in chemical reactions; they are not created or destroyed.

Example:
If you burn 5 grams of charcoal in air, the ash and gases formed will together weigh 5 grams. Nothing disappears or is created from nothing.

Exam tip:
This law helps us balance chemical equations. The total mass on both sides of the equation must be equal.

2. Law of Definite Proportions (Constant Composition)

What does it mean?
A pure chemical compound always contains the same elements in exactly the same fixed ratio by mass, no matter where it comes from or how it was made.

Why?
Because the atoms combine in fixed ratios to form a compound.

Example:
Water is always made up of hydrogen and oxygen in a 1:8 mass ratio. So, 1 gram of hydrogen always combines with 8 grams of oxygen to make water.

Exam tip:
This law confirms that compounds are pure substances with a fixed formula.

3. Law of Multiple Proportions

What does it mean?
If two elements can combine to form more than one compound, then the masses of one element that combine with a fixed mass of the other element are in the ratio of small whole numbers.

Why?
Because atoms combine in different ratios to make different compounds.

Example:
Carbon and oxygen form two compounds:

• Carbon monoxide (CO) — 12 g carbon combines with 16 g oxygen
• Carbon dioxide (CO₂) — 12 g carbon combines with 32 g oxygen

The ratio of oxygen masses = 16:32 = 1:2 (small whole numbers)

Exam tip:
This law supports the atomic theory and helps understand molecular formulas.

4. Gay-Lussac’s Law of Gaseous Volumes

What does it mean?
When gases react together at the same temperature and pressure, the volumes of the reacting gases and the volumes of the products (if gaseous) are in simple whole number ratios.

Why?
Because gas particles are counted by volume at the same conditions.

Example:
2 volumes of hydrogen gas react with 1 volume of oxygen gas to produce 2 volumes of water vapor:
2 H₂ + 1 O₂ → 2 H₂O (volumes ratio 2:1:2)

Exam tip:
This law is useful for solving problems related to gases and their reactions.

5. Avogadro’s Law

What does it mean?
Equal volumes of all gases, at the same temperature and pressure, contain the same number of molecules.

Why?
Because volume depends on the number of gas particles when temperature and pressure are fixed.

Example:
22.4 liters of hydrogen gas and 22.4 liters of oxygen gas at standard temperature and pressure (STP) contain the same number of molecules.

Exam tip:
This law helps in calculating molar volume and understanding mole concept in gases.

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5. Dalton’s Atomic Theory

What is Dalton’s Atomic Theory?

Dalton’s Atomic Theory was made by John Dalton in 1803. It tells us about the smallest parts of matter, called atoms, and how they behave in chemical reactions.

Main points of Dalton’s Atomic Theory:

1. Atoms are very small particles that cannot be divided.

• Everything is made of tiny particles called atoms.
• Atoms cannot be broken into smaller parts (Dalton thought so).

2. All atoms of the same element are exactly the same.

• Atoms of one element are identical in size, mass, and properties.
• Atoms of different elements are different from each other.

3. Atoms of different elements combine in simple ratios to make compounds.

• Atoms join in whole number ratios like 1:1, 2:1, etc.
• Example: Water (H₂O) is made of 2 hydrogen atoms and 1 oxygen atom.

4. Atoms cannot be created or destroyed in chemical reactions.

• Atoms just rearrange to form new substances.
• This supports the law of conservation of mass.

5. Atoms join, separate, or rearrange during chemical reactions.

• In reactions, atoms change partners but do not break or disappear.

What Dalton’s theory does not explain (limitations):

• Atoms are not actually indivisible—they have smaller parts (protons, neutrons, electrons).
• Some atoms of the same element can have different masses (called isotopes), but Dalton didn’t know this.
• Not all atoms of an element are exactly identical.

Why is Dalton’s Theory important for NEET?

• It helps us understand that matter is made of atoms.
• It explains how atoms combine in fixed ratios to make compounds.
• It supports balancing chemical equations and conservation of mass.

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6. Atomic and Molecular Masses

1. Atomic Mass

• Definition:
  Atomic mass is the mass of one atom of an element, usually measured in atomic mass units (amu).
• Atomic mass unit (amu):
  1 amu is defined as 1/12th of the mass of one carbon-12 atom.
• Example:
  Hydrogen atom has atomic mass approximately 1 amu. Oxygen atom has atomic mass approximately 16 amu.
• Note:
  The atomic mass you see on the periodic table is often an average of all isotopes of that element (see next point).

2. Average Atomic Mass

• Definition:
  Most elements have more than one isotope (atoms with the same number of protons but different number of neutrons). The average atomic mass is the weighted average of the masses of all isotopes of that element, based on their natural abundance.
• How to calculate:

Average atomic mass = ∑ (fractional abundance × mass of isotope)

• Example:
  Chlorine has two isotopes:
• Cl-35 (mass = 35 amu, abundance = 75%)
• Cl-37 (mass = 37 amu, abundance = 25%)

Average atomic mass = (0.75 × 35) + (0.25 × 37) = 26.25 + 9.25 = 35.5 amu

• This is why chlorine’s atomic mass on the periodic table is 35.5 amu, not a whole number.

3. Molecular Mass (Molecular Weight)

• Definition:
  Molecular mass is the sum of the atomic masses of all the atoms in a molecule. It is the total mass of one molecule of a compound, measured in amu.
• Example:
  Water (H₂O):
• Atomic mass of H = 1 amu
• Atomic mass of O = 16 amu
  Molecular mass = (2 × 1) + (1 × 16) = 18 amu
• Note:
  Molecular mass is used for molecules that exist as discrete units (like H₂O, CO₂, O₂).

4. Formula Mass (Formula Weight)

• Definition:
  Formula mass is the sum of atomic masses of all atoms in the formula unit of a compound. It is used for ionic compounds that do not exist as discrete molecules but as a repeating pattern of ions (like NaCl).
• Example:
  Sodium chloride (NaCl):
• Atomic mass of Na = 23 amu
• Atomic mass of Cl = 35.5 amu (average atomic mass)
  Formula mass = 23 + 35.5 = 58.5 amu
• Difference from molecular mass:
  Formula mass is used for ionic compounds (like salts), molecular mass for molecular compounds.

Summary Table:

Term	Meaning	Used for	Unit
Atomic Mass	Mass of one atom of an element	Elements	amu
Average Atomic Mass	Weighted average mass of isotopes	Elements	amu
Molecular Mass	Sum of atomic masses in a molecule	Molecular compounds	amu
Formula Mass	Sum of atomic masses in a formula unit	Ionic compounds	amu

Exam Tips:

• Always remember amu = atomic mass unit = 1/12th of Carbon-12 mass.
• Average atomic mass explains why periodic table atomic masses are decimals.
• Molecular mass = sum of atoms in molecule; formula mass = sum in formula unit.
• Know the difference between molecular and formula mass clearly.
• Practice calculating these with given isotopes and compounds.

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7. Mole Concept and Molar Masses

🔹 1. What is a Mole?

• Definition:
  A mole is a unit used to count very small particles like atoms, molecules, ions, or electrons.
  It’s just like saying 1 dozen = 12 things, but here:
  👉 1 mole = 6.022 × 10²³ particles
• This number (6.022 × 10²³) is called Avogadro’s Number or Avogadro’s Constant.

📌 Why use mole?

Atoms and molecules are extremely small. It’s hard to count them directly. So, we use the mole as a counting unit to simplify chemical calculations.

🔹 2. Avogadro’s Constant (6.022 × 10²³)

• Avogadro’s number is the number of particles in 1 mole of any substance.
  🔹 It could be atoms (in elements)
  🔹 Molecules (in compounds)
  🔹 Ions (in ionic compounds)

📌 Examples:

• 1 mole of carbon atoms = 6.022 × 10²³ atoms of carbon
• 1 mole of water (H₂O) = 6.022 × 10²³ molecules of water
• 1 mole of Na⁺ ions = 6.022 × 10²³ sodium ions

🔹 3. Molar Mass

• Definition:
  Molar mass is the mass of 1 mole of a substance (element or compound).
  It is numerically equal to the atomic/molecular/formula mass, but the unit is grams per mole (g/mol).

📌 Examples:

Substance	Atomic/Molecular Mass (amu)	Molar Mass (g/mol)
H	1	1 g/mol
O	16	16 g/mol
H₂O	18	18 g/mol
NaCl	58.5	58.5 g/mol

💡 Remember:

• Atomic mass (amu) → Mass of 1 atom
• Molar mass (g/mol) → Mass of 1 mole of atoms

🔹 4. Relationship Between Mole, Mass, and Number of Particles

Number of moles (n) = Given mass of substance(g) / Molar mass (g/mol)

💬 Formula 2:

Number of particles (atoms/molecules/ions) = Moles × 6.022×10²³

💬 Formula 3 (for mass):

Mass of substance (g) = Number of moles × Molar mass

🔹 5. Exam Tips & Tricks

• ✅ Always use correct units (g for mass, mol for mole, g/mol for molar mass).
• ✅ Use Avogadro’s number only when asked about atoms/molecules/ions.
• ✅ Be careful: O₂ is a molecule, but O is an atom.
• ✅ Molar mass is a bridge between mass ↔ moles ↔ number of particles.

🔹 Summary Table

Quantity	Symbol	Formula
Number of moles	n	n = mass / molar mass
Number of particles	N	N = moles × 6.022 × 10²³
Mass of substance	m	m = n × molar mass

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🔹 What is Percentage Composition?

Percentage composition tells us how much percent by mass of each element is present in a compound.

It shows the relative mass contribution of each element to the total mass of the compound.

🔹 Why is it important?

• It helps us analyze a compound’s formula.
• Useful in finding empirical or molecular formula.
• Frequently used in stoichiometry and chemical calculations in NEET.

🔹 Formula for Percentage Composition:-

Percentage of an element = ( Total mass of that element in 1 mole of compound​ / Molar mass of compound) × 100

🔹 Step-by-step Method:-

1. Write molecular formula of compound.
2. Calculate molar mass of compound.
3. Find mass of each element in the formula.
4. Use the formula to calculate % of each element.
5. Final check: All percentages should add up to ~100%.

🔹 Exam Tips

• ✅ Always use atomic masses correctly (e.g. H = 1, O = 16, N = 14).
• ✅ Don’t forget the number of atoms of each element in the formula.
• ✅ Your final percentages must add up to ~100%.
• ✅ NEET often asks this in the form of MCQs or numerical problems.

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9. Empirical and Molecular Formula

🔹 1. What is Empirical Formula?

Empirical formula shows the simplest whole number ratio of atoms of each element in a compound.

👉 It does not show the actual number of atoms, just the ratio.

✅ Example:

• H₂O₂ (Hydrogen Peroxide)
  • Empirical Formula = HO
  • Molecular Formula = H₂O₂

🔹 2. What is Molecular Formula?

Molecular formula shows the actual number of atoms of each element in a molecule of the compound.

👉 It is usually a whole-number multiple of the empirical formula.

✅ Example:

• Glucose → Molecular Formula = C₆H₁₂O₆
  • Empirical Formula = CH₂O

🔹 3. Difference Between Empirical and Molecular Formula

Feature	Empirical Formula	Molecular Formula
Shows	Simplest ratio	Actual atoms per molecule
May be same as MF?	Yes	Always actual formula
Example (H₂O₂)	HO	H₂O₂
Use	Used for unknown compounds	Used when molar mass is known

🔹 4. Relation Between Empirical and Molecular Formula

Molecular Formula = n × (Empirical Formula)

Where:

• n = Molar Mass of Compound​ / Empirical Formula Mass​

👉 Value of n must be a whole number.

🔹 5. Steps to Find Empirical Formula from Percentage Composition

Suppose % composition is given. Follow these steps:

✅ Step 1:

Write % of each element (treat them as grams if not specified).

✅ Step 2:

Convert the mass of each element to moles:- Moles = Given mass / Molar mass​s​

✅ Step 3:

Divide all mole values by the smallest mole value.

✅ Step 4:

Round or multiply to get whole numbers (1, 2, 3,…).

✅ Step 5:

Write empirical formula using these whole numbers.

🔹 6. How to Find Molecular Formula (when molar mass is known)

After finding empirical formula:

1. Find empirical formula mass (sum of atomic masses in the EF).
2. Use: n = Molar Mass / Empirical Formula Mass
3. Multiply empirical formula by nnn to get the molecular formula.

🔹 8. Common NEET/JEE Tips:

• Always use accurate atomic masses.
• If mole ratios are like 1.5, 2.5, etc., multiply all values by 2 to make them whole.
• Empirical formula is crucial in organic chemistry and stoichiometry.
• NEET often gives % composition or mass of elements.

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10. Stoichiometry & Stoichiometric Calculations

🔹 1. What is Stoichiometry?

Stoichiometry is the study of quantitative relationships between reactants and products in a chemical reaction.

It is based on the law of conservation of mass.

✍️ Example:

In the reaction: 2H₂​ + O₂​ → 2H₂​O

This tells us:

• 2 moles of hydrogen react with 1 mole of oxygen
• to produce 2 moles of water

🔹 2. Balanced Chemical Equations

A balanced chemical equation shows the exact number of moles of reactants and products.

Balancing ensures that the number of atoms of each element is equal on both sides of the equation.

✍️ Example:

C₃​H₈ ​+ 5O₂ ​→ 3CO₂ ​+ 4H₂​O

Means:

• 1 mol propane + 5 mol oxygen → 3 mol carbon dioxide + 4 mol water

🔹 3. Mole-to-Mole Calculations

These involve converting from moles of one substance to moles of another using the mole ratio from the balanced equation.

✅ Formula:

Moles of unknown = Moles of known × ( Coefficient of unknown ​/ Coefficient of known)

🔹 4. Mole-to-Mass Calculations

First, convert known moles to mass using:- Mass = Moles × Molar Mass

✍️ Example:

How many grams of H₂O are formed from 2 mol H₂?

Equation: 2H₂​ + O₂​ → 2H₂​O

Moles of H₂O = 2
Molar mass of H₂O = 18 g/mol

Answer: 2×18=36 g

🔹 5. Mass-to-Mass Calculations

Steps:

1. Convert given mass to moles
2. Use mole ratio from balanced equation
3. Convert back to mass

✍️ Example:

How many grams of CO₂ are produced from 44 g of propane (C₃H₈)?

1. Molar mass of C₃H₈ = 44 g/mol
   → Moles = 44 / 44 = 1 mol
2. From equation:- C₃​H₈ ​+ 5O₂ ​→ 3CO₂ ​+ 4H₂​O

1 mol C₃H₈ → 3 mol CO₂

3. Moles of CO₂ = 3
   Molar mass CO₂ = 44 g/mol

→ Mass = 3 × 44 = 132 g

🔹 6. Limiting Reagent (or Limiting Reactant)

It is the reactant that gets used up first in a chemical reaction. It limits the amount of product formed.

Steps to Identify:

1. Convert mass to moles for all reactants
2. Divide by their stoichiometric coefficients
3. The smaller value indicates the limiting reagent

✍️ Example:

Given: 4 mol H₂ and 3 mol O₂
Reaction: 2H₂​ + O₂​ → 2H₂​O

Check:

• H₂ needed = 2 mol → makes 2 mol H₂O
• O₂ needed = 1 mol → makes 2 mol H₂O

From 4 mol H₂ → needs 2 mol O₂
We have only 3 mol O₂ → enough
So H₂ is limiting reagent

🔹 7. Theoretical Yield

The maximum amount of product expected from a given amount of reactants, assuming complete reaction.

Formula:

Theoretical yield = Moles of product × Molar mass of product

🔹 8. Percentage Yield

Tells us how efficient the reaction is.

Formula:

Percentage Yield = (Actual Yield / Theoretical Yield​) × 100

✍️ Example:

Theoretical yield = 100 g
Actual yield = 80 g

Percentage yield = (80/100​) × 100 = 80%

🔹 Summary Table

Concept	Formula / Key Idea
Balanced Equation	Equal atoms on both sides
Mole-to-mole	Use mole ratio from equation
Mole-to-mass	Multiply moles by molar mass
Mass-to-mass	Convert to moles → use ratio → convert to mass
Limiting Reagent	Reactant that finishes first
Theoretical Yield	Max possible product (in g)
% Yield	(Actual / Theoretical) × 100

🔹 NEET & JEE Tips:

• Always balance the equation first.
• Use correct molar masses and units.
• NEET asks direct mole calculations and limiting reagent MCQs.
• JEE can include more numerical stoichiometry and word problems.
• Practice mass-to-mass conversions and % yield questions.

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