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Complete Some Basic Concepts of Chemistry Class 11 notes, summary, MCQs, important questions, keywords, and exam tips for quick revision and exam success.

Introduction of the Chapter

Some Basic Concepts of Chemistry Class 11 is the foundation chapter of chemistry. It introduces the basic language and calculations used throughout chemistry. This chapter explains the importance of measurement, units, scientific notation, mole concept, atomic mass, molecular mass, percentage composition, empirical and molecular formula, and stoichiometry.

Understanding Some Basic Concepts of Chemistry is essential because all numerical chemistry problems depend on these ideas. From laboratory measurements to chemical reactions, these concepts help students understand matter quantitatively.

This chapter builds problem-solving skills and strengthens logical thinking, which is useful for board exams and competitive exams.

Short Notes (Quick Revision Points)

Chemistry deals with the composition, structure, and properties of matter.
Matter exists in three physical states: solid, liquid, and gas.
Physical properties can be observed without changing identity.
Chemical properties involve chemical change.
SI unit system is used for scientific measurements.
Significant figures indicate measurement accuracy.
Laws of chemical combination explain how elements combine.
Atomic mass is relative mass of an atom compared to 1/12 of carbon-12.
Mole concept connects microscopic particles with measurable quantities.
1 mole = particles.
Molar mass = mass of 1 mole of substance.
Empirical formula shows simplest ratio of atoms.
Molecular formula shows actual number of atoms.
Stoichiometry deals with quantitative relationships in reactions.
Limiting reagent determines product amount.

Detailed Summary (1200 Words)

Some Basic Concepts of Chemistry Class 11 provides the basic framework for understanding chemical science. Chemistry deals with matter, its properties, composition, and the changes it undergoes.

Nature of Matter

Matter is anything that occupies space and has mass. It exists in three physical states: solids, liquids, and gases. Solids have fixed shape and volume, liquids have fixed volume but variable shape, and gases have neither fixed shape nor volume.

Matter has physical properties such as color, density, and melting point. Chemical properties describe how substances react to form new substances.

Measurement and Units

Measurement is fundamental in chemistry. Scientists use the International System of Units (SI). Common units include:

Length: meter (m)
Mass: kilogram (kg)
Time: second (s)
Temperature: kelvin (K)
Amount of substance: mole (mol)

Accurate measurement is important. Scientific notation helps express very large or small numbers.

Uncertainty and Significant Figures

No measurement is perfectly exact. Significant figures indicate precision. Rules include:

All non-zero digits are significant.
Zeros between digits are significant.
Leading zeros are not significant.
Trailing zeros in decimals are significant.

Laws of Chemical Combination

Several laws explain how elements combine:

Law of Conservation of Mass (Antoine Lavoisier)
Mass is neither created nor destroyed in a chemical reaction.

Law of Definite Proportions (Joseph Proust)
A compound always contains elements in fixed proportions.

Law of Multiple Proportions (John Dalton)
Elements combine in simple whole-number ratios.

Atomic and Molecular Mass

Atomic mass is the average mass of atoms of an element compared with carbon-12.

Molecular mass is the sum of atomic masses of atoms in a molecule.

Formula mass is used for ionic compounds.

Mole Concept

The mole concept is central to Some Basic Concepts of Chemistry. It provides a link between microscopic particles and measurable quantities.

1 mole of any substance contains Avogadro’s number:

particles

Molar mass is the mass of one mole of a substance.

Examples:

1 mole of carbon = 12 g
1 mole of water = 18 g

Percentage Composition

It tells the mass percentage of each element in a compound.


\% \text{ element} = \frac{\text{mass of element}}{\text{molar mass}} \times 100

Empirical Formula

The empirical formula shows the simplest ratio of atoms in a compound.

Steps:

Convert percentage to grams.
Divide by atomic mass.
Find simplest ratio.

Molecular Formula

The molecular formula shows the actual number of atoms in a molecule.


\text{Molecular Formula} = (\text{Empirical Formula})_n

Stoichiometry

Stoichiometry deals with quantitative relationships in chemical reactions.

Example:


2H_2 + O_2 → 2H_2O

This shows 2 moles of hydrogen react with 1 mole of oxygen.

Limiting Reagent

The limiting reagent is the reactant that is completely consumed first and limits product formation.

Importance of Stoichiometry

Helps calculate product yield
Used in chemical industries
Essential in laboratory experiments

Thus, Some Basic Concepts of Chemistry Class 11 builds numerical and conceptual understanding necessary for future chemistry learning.

Flowchart / Mind Map (Text-Based)

Chemistry
→ Matter
→ Properties (Physical / Chemical)
→ Measurement
→ SI Units
→ Significant Figures
→ Laws of Chemical Combination
→ Atomic & Molecular Mass
→ Mole Concept
→ Avogadro Number
→ Molar Mass
→ Percentage Composition
→ Empirical Formula
→ Molecular Formula
→ Stoichiometry
→ Limiting Reagent

Important Keywords with Meanings

Matter – Anything with mass and volume
Physical Property – Property observed without chemical change
Chemical Property – Property describing chemical reaction
Atomic Mass – Relative mass of an atom
Molecular Mass – Sum of atomic masses
Mole – Unit representing particles
Molar Mass – Mass of one mole of substance
Stoichiometry – Quantitative study of reactions
Empirical Formula – Simplest atom ratio
Molecular Formula – Actual atom count
Limiting Reagent – Reactant that limits product formation
Significant Figures – Digits showing measurement precision

Important Questions & Answers

Short Answer Questions

1. Define matter.
Anything that has mass and occupies space.

2. What is a mole?
A mole is the amount containing particles.

3. Define molar mass.
Mass of one mole of a substance.

4. State law of conservation of mass.
Mass remains constant during chemical reactions.

5. What is atomic mass?
Relative mass of an atom compared to carbon-12.

6. Define empirical formula.
Simplest ratio of atoms in a compound.

7. What is stoichiometry?
Study of quantitative relationships in reactions.

8. What is limiting reagent?
Reactant consumed first.

9. Define significant figures.
Digits indicating measurement accuracy.

10. Write Avogadro number.

Long Answer Questions

1. Explain the mole concept.
The mole concept links microscopic particles with measurable quantities. One mole contains particles. It allows calculation of masses, molecules, and reaction quantities.

2. Explain laws of chemical combination.
Lavoisier proposed conservation of mass; Proust stated definite proportions; Dalton proposed multiple proportions.

3. What are significant figures? Explain rules.
They show precision. Non-zero digits are significant; leading zeros are not; trailing decimals are significant.

4. Explain empirical and molecular formulas.
Empirical formula shows simplest ratio; molecular formula shows actual number of atoms.

5. Define stoichiometry with example.
It calculates quantities in reactions using balanced equations.

6. Explain percentage composition.
It gives percentage mass of elements in compounds.

7. Define atomic mass and molecular mass.
Atomic mass is relative mass of an atom; molecular mass is sum of atomic masses.

8. Explain limiting reagent with example.
It is the reactant consumed first and determines product amount.

9. Describe SI units used in chemistry.
Meter, kilogram, second, kelvin, mole.

10. Why is the mole concept important?
It helps in calculations, industrial production, and reaction prediction.

MCQs with Answers

Matter occupies:
A) time
B) space ✔
C) energy
D) light
SI unit of amount of substance:
A) gram
B) mole ✔
C) liter
D) kg
Avogadro number is:
A)
B) ✔
C)
D)
Law of conservation of mass given by:
A) Dalton
B) Proust
C) Lavoisier ✔
D) Rutherford
Empirical formula represents:
A) actual atoms
B) simplest ratio ✔
C) molecular mass
D) valency
Molecular mass of H₂O:
A) 16
B) 17
C) 18 ✔
D) 19
Limiting reagent determines:
A) speed
B) product formed ✔
C) color
D) temperature
Atomic mass unit is based on:
A) oxygen
B) hydrogen
C) carbon-12 ✔
D) nitrogen
1 mole of carbon weighs:
A) 6 g
B) 12 g ✔
C) 24 g
D) 1 g
Significant figures indicate:
A) size
B) accuracy ✔
C) mass
D) color
Formula mass applies to: ionic compounds ✔
Mole concept connects microscopic and macroscopic ✔
Stoichiometry uses balanced equations ✔
Law of definite proportions given by Proust ✔
Leading zeros are not significant ✔
Molar mass unit is g/mol ✔
Molecules in 1 mole = ✔
Mass remains constant in reactions ✔
Percent composition uses mass ✔
Limiting reagent is fully consumed ✔

Exam Tips & Value-Based Questions

Exam Tips

Always balance equations before calculations.
Learn atomic masses of common elements.
Practice mole conversions daily.
Use dimensional analysis to avoid mistakes.
Keep track of units carefully.

Value-Based Questions

1. Why is accurate measurement important in chemistry?
It ensures reliable results and safe experiments.

2. How does conservation of mass promote environmental responsibility?
It encourages waste control and recycling.

3. Why should chemicals be measured precisely in medicine?
To ensure correct dosage and patient safety.

4. How does stoichiometry help in reducing industrial waste?
It optimizes reactant use and reduces pollution.

5. Why is scientific notation useful in research?
It simplifies handling large and small numbers.

Conclusion 💫

Some Basic Concepts of Chemistry is the cornerstone of chemical science and forms the basis for all future chemistry learning. This chapter introduces students to the essential principles required to understand the composition and behavior of matter. From measurement and units to the mole concept and stoichiometry, each topic plays a crucial role in building analytical and numerical skills.

Understanding Some Basic Concepts of Chemistry helps students interpret chemical reactions quantitatively and accurately. The SI system of units ensures uniformity in scientific measurements across the world. Learning significant figures and scientific notation improves precision and accuracy in calculations, which is vital for laboratory work and examinations.

The laws of chemical combination provide the theoretical foundation for chemical reactions. These laws explain how substances combine and why reactions follow fixed ratios. Atomic mass and molecular mass help determine the mass relationships between elements and compounds, enabling students to perform accurate calculations.

The mole concept is one of the most important topics in Some Basic Concepts of Chemistry Class 11. It connects the microscopic world of atoms and molecules with measurable laboratory quantities. By understanding moles, molar mass, and Avogadro’s number, students can calculate the number of particles, masses, and reaction yields efficiently.

Stoichiometry further extends this understanding by enabling the calculation of reactants and products in chemical reactions. It plays a significant role in chemical industries, pharmaceuticals, environmental science, and laboratory research. The concept of limiting reagent ensures efficient use of resources and prevents wastage.

This chapter also strengthens logical reasoning and problem-solving abilities. Mastering numerical problems from Some Basic Concepts of Chemistry improves accuracy and confidence in examinations and competitive tests.

For students preparing for board exams, NEET, JEE, and other competitive exams, this chapter is extremely important. Questions based on mole concept, empirical formula, and stoichiometry frequently appear in examinations.

By thoroughly understanding Some Basic Concepts of Chemistry , students develop a strong foundation for advanced topics such as thermodynamics, equilibrium, and organic chemistry.

Regular practice, conceptual clarity, and careful unit handling are the keys to mastering this chapter. With strong fundamentals from Some Basic Concepts of Chemistry, students can confidently approach chemistry as a logical and interesting subject rather than a difficult one.

In summary, this chapter is not only essential for academic success but also for developing scientific thinking and real-world problem-solving skills.

Long Question Answer

1. Explain the Laws of Chemical Combination.

The laws of chemical combination explain how elements combine to form compounds.

Law of Conservation of Mass (Antoine Lavoisier):
Mass can neither be created nor destroyed during a chemical reaction. The total mass of reactants equals the total mass of products.

Law of Definite Proportions (Joseph Proust):
A chemical compound always contains the same elements in the same fixed proportion by mass.

Law of Multiple Proportions (John 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 simple whole-number ratios.

These laws form the basis of modern chemistry.

2. Describe the SI System of Units and its Importance.

The International System of Units (SI) is the standard system used worldwide for scientific measurements.

Base Units:

Length → meter (m)
Mass → kilogram (kg)
Time → second (s)
Temperature → kelvin (K)
Amount of substance → mole (mol)

Importance:

Ensures uniformity worldwide
Avoids confusion in measurements
Essential for scientific communication
Used in laboratories and industries

3. What are Significant Figures? Explain the Rules.

Significant figures represent the precision of a measurement.

Rules:

All non-zero digits are significant.
Zeros between digits are significant.
Leading zeros are not significant.
Trailing zeros after decimal are significant.

Example: 0.00450 has 3 significant figures.

Significant figures help express measurement accuracy.

4. Explain the Mole Concept in Detail.

The mole concept connects microscopic particles to measurable quantities.

One mole of any substance contains:


6.022 \times 10^{23}

particles (Avogadro number).

Key Points:

1 mole of carbon = 12 g
1 mole of water = 18 g
1 mole contains equal number of particles

Importance:

Helps calculate masses and particles
Essential in chemical reactions
Used in stoichiometry

5. Define Atomic Mass and Molecular Mass.

Atomic Mass:
The relative mass of an atom compared to 1/12 of the mass of carbon-12.

Molecular Mass:
The sum of atomic masses of all atoms in a molecule.

Example:
Water (H₂O)
= 2(1) + 16 = 18 u

These values help in calculating molar mass and stoichiometric relationships.

6. Explain Percentage Composition and its Calculation.

Percentage composition tells the mass percentage of each element in a compound.


\% \text{Element} = \frac{\text{Mass of element}}{\text{Molar mass}} \times 100

Example: Water (H₂O)

Hydrogen =
Oxygen =

This helps determine formulas and purity.

7. What is an Empirical Formula? Explain the Steps to Find It.

The empirical formula shows the simplest whole-number ratio of atoms in a compound.

Steps:

Convert percentages into grams.
Divide by atomic masses.
Find simplest ratio.
Multiply to get whole numbers.

Example: CH₂O is empirical formula of glucose.

8. Explain Molecular Formula and How It is Determined.

The molecular formula shows the actual number of atoms in a molecule.


\text{Molecular Formula} = (\text{Empirical Formula})_n

where


n = \frac{\text{Molecular mass}}{\text{Empirical formula mass}}

Example: Empirical formula = CH₂O
Molecular mass = 180
Empirical mass = 30
n = 6

Molecular formula = C₆H₁₂O₆

9. Explain Stoichiometry with an Example.

Stoichiometry deals with quantitative relationships in chemical reactions.

Example reaction:


2H_2 + O_2 → 2H_2O

This means:

2 moles hydrogen react with 1 mole oxygen
2 moles water are formed

Uses:

Calculate product yield
Industrial production
Laboratory calculations

10. What is a Limiting Reagent? Explain with Example.

The limiting reagent is the reactant that is completely consumed first in a reaction and determines the amount of product formed.

Example:

If 2 moles of H₂ react with 1 mole O₂:


2H_2 + O_2 → 2H_2O

If only 1 mole H₂ is available, hydrogen becomes the limiting reagent and only 1 mole of water will form.

Importance:

Prevents wastage of chemicals
Used in industrial efficiency
Helps calculate product yield

Short Question Answer

1. What is a physical property?

A physical property is a characteristic that can be observed without changing the substance’s chemical identity (e.g., color, melting point).

2. What is a chemical property?

A chemical property describes the ability of a substance to undergo a chemical change (e.g., burning, rusting).

3. What is scientific notation?

Scientific notation is a way of expressing very large or very small numbers using powers of 10.

4. What is the formula mass?

Formula mass is the sum of atomic masses of all atoms present in an ionic compound.

5. What is the mass of one mole of oxygen atoms?

The mass of one mole of oxygen atoms is 16 g.

6. Define Avogadro’s number.

Avogadro’s number is , the number of particles present in one mole of a substance.

7. What are significant figures in measurement?

Significant figures are the digits in a measurement that indicate its precision.

8. State the law of definite proportions.

A chemical compound always contains the same elements in a fixed ratio by mass. This law was proposed by Joseph Proust.

9. What is the gram atomic mass of an element?

The gram atomic mass is the atomic mass of an element expressed in grams.

10. What is a balanced chemical equation?

A balanced chemical equation has equal numbers of atoms of each element on both sides of the equation.

Assertion Reason 💫

Directions:
For each question, choose the correct option:

A. Both Assertion (A) and Reason (R) are true, and R is the correct explanation of A.
B. Both A and R are true, but R is not the correct explanation of A.
C. A is true, but R is false.
D. A is false, but R is true.

1.

Assertion (A): Matter has mass and occupies space.
Reason (R): Matter is made up of atoms and molecules.

Answer: B

2.

Assertion (A): The SI unit of temperature is kelvin.
Reason (R): Kelvin scale starts from absolute zero.

Answer: A

3.

Assertion (A): One mole of all gases occupies the same volume at STP.
Reason (R): Equal volumes of gases contain equal number of molecules at same temperature and pressure.

Answer: A
(This is based on Avogadro’s law.)

4.

Assertion (A): Significant figures represent precision in measurement.
Reason (R): All zeros in a number are significant.

Answer: C

5.

Assertion (A): Atomic mass is expressed in atomic mass unit (u).
Reason (R): One atomic mass unit is defined as 1/12 the mass of a carbon-12 atom.

Answer: A

6.

Assertion (A): Empirical formula represents the simplest ratio of atoms.
Reason (R): Molecular formula represents the actual number of atoms in a molecule.

Answer: B

7.

Assertion (A): Mass is conserved in a chemical reaction.
Reason (R): Atoms are neither created nor destroyed during a chemical reaction.

Answer: A
(This law was established by Antoine Lavoisier.)

8.

Assertion (A): Molar mass of a substance is numerically equal to its molecular mass.
Reason (R): Molecular mass is expressed in grams.

Answer: C
(Molecular mass is in atomic mass units.)

9.

Assertion (A): Limiting reagent controls the amount of product formed.
Reason (R): It is completely consumed during the reaction.

Answer: A

10.

Assertion (A): Stoichiometry is based on balanced chemical equations.
Reason (R): Balanced equations follow the law of conservation of mass.

Answer: A

Cased Based Questions 💫

Case 1: Measurement and Units

A student measures the length of a table as 2.50 m using a meter scale.

Questions:

How many significant figures are present?
What does the last digit represent?
Write the value in centimeters.

Answers:

3 significant figures
Measurement precision
250 cm

Case 2: Scientific Notation

The mass of an atom is g.

Questions:

Write the number in scientific notation.
Why is scientific notation useful?

Answers:

g
It simplifies writing very small or large numbers.

Case 3: Law of Conservation of Mass

When 10 g of calcium carbonate is heated, it forms 5.6 g calcium oxide and carbon dioxide gas.

Questions:

Calculate the mass of CO₂ formed.
Name the law verified.
Who proposed this law?

Answers:

4.4 g
Law of conservation of mass
Antoine Lavoisier

Case 4: Mole Concept

A sample contains molecules of oxygen gas.

Questions:

How many moles are present?
What is the mass of this oxygen sample?

Answers:

1 mole
32 g

Case 5: Atomic and Molecular Mass

Calculate the molecular mass of carbon dioxide (CO₂).

Questions:

Write atomic masses used.
Calculate molecular mass.
What is its molar mass?

Answers:

C = 12, O = 16
44 g/mol

Case 6: Percentage Composition

A compound contains 40% carbon, 6.7% hydrogen, and 53.3% oxygen.

Questions:

Which element is highest in percentage?
Why is percentage composition useful?

Answers:

Oxygen
It helps determine empirical formula and purity.

Case 7: Empirical Formula

A compound contains 27.3% carbon and 72.7% oxygen.

Questions:

Convert percentages to simplest ratio.
Determine empirical formula.

Answers:

C : O ≈ 1 : 2
CO₂

Case 8: Limiting Reagent

Hydrogen reacts with oxygen:


2H_2 + O_2 → 2H_2O

If 2 moles H₂ react with 2 moles O₂:

Questions:

Which is the limiting reagent?
How many moles of water form?

Answers:

Hydrogen
2 moles

Case 9: Significant Figures in Calculation

Multiply 2.5 × 3.42.

Questions:

Calculate the result.
How many significant figures should the answer have?

Answers:

8.55
2 significant figures → 8.6

Case 10: Stoichiometry in Daily Life

In a laboratory, 4 g hydrogen reacts with excess oxygen to produce water.

Questions:

Write the balanced equation.
Calculate moles of hydrogen used.
Calculate moles of water formed.

Answers:

moles H₂
2 moles H₂O

💫Multiple Choice Questions

1. Matter is anything that:

A. has weight only
B. occupies space and has mass ✔
C. has volume only
D. produces energy

2. The SI unit of mass is:

A. gram
B. kilogram ✔
C. pound
D. liter

3. The SI unit of temperature is:

A. Celsius
B. Fahrenheit
C. Kelvin ✔
D. Joule

4. Scientific notation is used to express:

A. chemical reactions
B. large and small numbers ✔
C. temperature
D. volume

5. 1 nanometer equals:

A. m
B. m ✔
C. m
D. m

6. Law of conservation of mass was proposed by:

A. John Dalton
B. Joseph Proust
C. Antoine Lavoisier ✔
D. Ernest Rutherford

7. A pure substance is:

A. mixture
B. compound ✔
C. solution
D. suspension

8. The law of definite proportions was given by:

A. Dalton
B. Proust ✔
C. Lavoisier
D. Thomson

9. Atomic mass is based on:

A. hydrogen atom
B. oxygen atom
C. carbon-12 atom ✔
D. nitrogen atom

10. The value of Avogadro’s number is:

A.
B.
C. ✔
D.

11. One mole of any substance contains:

A. 12 particles
B. particles ✔
C. 1 particle
D. 100 particles

12. The molar mass of water is:

A. 16 g/mol
B. 17 g/mol
C. 18 g/mol ✔
D. 20 g/mol

13. Which has the highest mass?

A. 1 mole H₂
B. 1 mole O₂ ✔
C. 1 mole He
D. 1 mole N₂

14. Empirical formula represents:

A. actual number of atoms
B. simplest ratio ✔
C. molecular mass
D. valency

15. Molecular formula of glucose is:

A. CH₂O
B. C₂H₄O₂
C. C₆H₁₂O₆ ✔
D. CHO

16. Significant figures indicate:

A. color
B. precision ✔
C. density
D. weight

17. Zeros at the beginning of a number are:

A. significant
B. not significant ✔
C. sometimes significant
D. always counted

18. 0.00560 has how many significant figures?

A. 2
B. 3 ✔
C. 4
D. 5

19. The gram molecular mass of CO₂ is:

A. 22
B. 32
C. 44 ✔
D. 48

20. Stoichiometry deals with:

A. structure of atoms
B. reaction quantities ✔
C. colors of compounds
D. states of matter

21. Limiting reagent is the reactant that:

A. is left over
B. is in excess
C. is consumed first ✔
D. slows reaction

22. The number of atoms in 1 mole of helium is:

A. 1
B. ✔
C.
D.

23. Percentage composition is based on:

A. number of atoms
B. mass ✔
C. volume
D. density

24. Which is a derived unit?

A. meter
B. mole
C. kilogram
D. density ✔

25. Balanced chemical equations follow:

A. law of motion
B. law of conservation of mass ✔
C. law of gravity
D. law of inertia

26. Molecular mass is expressed in:

A. grams
B. kilograms
C. atomic mass units ✔
D. liters

27. The simplest formula of benzene (C₆H₆) is:

A. CH ✔
B. CH₂
C. C₂H₂
D. C₃H₃

28. 2 moles of H₂ contain:

A. molecules
B. molecules
C. molecules ✔
D. 2 molecules

29. If mass is constant during a reaction, it supports:

A. definite proportions
B. conservation of mass ✔
C. multiple proportions
D. periodic law

30. Which quantity connects microscopic particles with measurable mass?

A. atomic number
B. mole concept ✔
C. valency
D. density

Sample Paper 📜

Below is a Sample Question Paper for Class 11 Chemistry – Some Basic Concepts of Chemistry.
It follows the pattern used in school exams and is useful for revision and practice.

Sample Question Paper

Some Basic Concepts of Chemistry (Class 11)

Time: 3 Hours
Maximum Marks: 70

Section A — Very Short Answer (1 Mark Each)

Q1. Define matter.
Q2. Write the SI unit of amount of substance.
Q3. What is atomic mass unit (u)?
Q4. Write Avogadro’s number.
Q5. Define molar mass.
Q6. What is a significant figure?
Q7. State the law of definite proportions.
Q8. Write the empirical formula of benzene (C₆H₆).
Q9. What is stoichiometry?
Q10. Define limiting reagent.

Section B — Short Answer Questions (2 Marks Each)

Q11. Distinguish between physical and chemical properties.

Q12. Write any two base SI units used in chemistry.

Q13. How many significant figures are present in 0.00450?

Q14. Calculate the number of moles in 11 g of CO₂.

Q15. Write the formula to calculate percentage composition.

Section C — Short Answer Questions (3 Marks Each)

Q16. State the law of conservation of mass with an example. (Law proposed by Antoine Lavoisier)

Q17. Explain the mole concept.

Q18. Differentiate between empirical and molecular formula.

Q19. Calculate the molecular mass of H₂SO₄.
(H = 1, S = 32, O = 16)

Q20. What are significant figures? State the rules.

Section D — Long Answer Questions (5 Marks Each)

Q21. Explain the laws of chemical combination.

Q22. Describe SI units and their importance in scientific measurement.

Q23. A compound contains 40% carbon, 6.7% hydrogen, and 53.3% oxygen. Determine its empirical formula.

Q24. Explain limiting reagent with a suitable example.

Q25. Write the steps involved in calculating molecular formula from empirical formula.

Section E — Numerical Problems (5 Marks Each)

Q26. Calculate the mass of 0.5 mole of sodium carbonate (Na₂CO₃).
(Na = 23, C = 12, O = 16)

Q27. How many molecules are present in 2 moles of water?

Q28. Calculate the percentage of oxygen in CaCO₃.
(Ca = 40, C = 12, O = 16)

Section F — Case-Based Question (5 Marks)

A student heated 20 g of calcium carbonate and obtained 11.2 g calcium oxide and carbon dioxide gas.

Answer the following:

a) Calculate the mass of CO₂ formed.
b) Name the law verified in this experiment.
c) Why is this law important in chemistry?

Section G — Multiple Choice Questions (1 Mark Each)

Q29. 1 mole of a substance contains:
A)
B) ✔
C)
D)

Q30. Molecular mass of CO₂ is:
A) 28
B) 32
C) 44 ✔
D) 48

Solution 💫

Some Basic Concepts of Chemistry

Section A — Very Short Answers

1. Matter is anything that has mass and occupies space.

2. Mole (mol)

3. One atomic mass unit (u) is defined as 1/12 of the mass of a carbon-12 atom.

4.6.22×10²³

5. Molar mass is the mass of one mole of a substance expressed in g/mol.

6. Significant figures are digits in a measurement that indicate its precision.

7. A compound always contains the same elements in a fixed ratio by mass.
(Proposed by Joseph Proust)

8. CH

9. Stoichiometry is the quantitative relationship between reactants and products.

10. The limiting reagent is the reactant that is completely consumed first.

Section B — Short Answers

11. Physical vs Chemical Properties

Physical Property	Chemical Property
Observed without changing identity	Involves chemical change
Example: color, density	Example: rusting, burning

12. Two SI Base Units

Mass → kilogram (kg)
Temperature → kelvin (K)

13. Significant figures in 0.00450

= 3 significant figures
(Leading zeros not significant; trailing zero after decimal is significant)

14. Moles in 11 g CO₂

Molar mass CO₂ = 44 g/mol


\text{Moles} = \frac{11}{44} = 0.25 \text{ mol}

15. Formula for Percentage Composition


\% \text{ element} = \frac{\text{mass of element}}{\text{molar mass}} \times 100

Section C — Answers

16. Law of Conservation of Mass

Mass is neither created nor destroyed during a chemical reaction.

Example:


CaCO_3 → CaO + CO_2

Total mass before reaction = total mass after reaction.
Proposed by Antoine Lavoisier.

17. Mole Concept

A mole represents particles.

It connects microscopic particles with measurable mass.

Example:

1 mole H₂O = 18 g
1 mole O₂ = 32 g

18. Empirical vs Molecular Formula

Empirical Formula	Molecular Formula
Simplest ratio	Actual number
CH₂O	C₆H₁₂O₆

19. Molecular Mass of H₂SO₄

H₂ = 2 × 1 = 2
S = 32
O₄ = 4 × 16 = 64

Total = 98 u

20. Rules of Significant Figures

Non-zero digits are significant
Zeros between digits are significant
Leading zeros are not significant
Trailing zeros in decimals are significant

Section D — Long Answers

21. Laws of Chemical Combination

Law of Conservation of Mass — mass remains constant.
Law of Definite Proportions — fixed ratio of elements.
Law of Multiple Proportions — simple whole number ratios.

22. SI Units and Importance

SI system provides universal measurement standards.

Examples:

meter (length)
kilogram (mass)
kelvin (temperature)
mole (amount)

Importance:

uniformity worldwide
accurate scientific communication
avoids confusion

23. Empirical Formula Calculation

Given:
C = 40%, H = 6.7%, O = 53.3%

Step 1: Divide by atomic masses

C:
H:
O:

Step 2: Divide by smallest value (3.33)

C = 1
H = 2
O = 1

Empirical formula = CH₂O

24. Limiting Reagent

The limiting reagent is consumed first and determines product formed.

Example:


2H_2 + O_2 → 2H_2O

If hydrogen is less, it limits water formation.

25. Steps to Find Molecular Formula

Determine empirical formula.
Calculate empirical formula mass.
Use:


n = \frac{\text{molecular mass}}{\text{empirical formula mass}}

Multiply empirical formula by n.

Section E — Numerical Solutions

26. Mass of 0.5 mole Na₂CO₃

Molar mass:

Na₂ = 46
C = 12
O₃ = 48

Total = 106 g/mol


0.5 × 106 = 53 \text{ g}

Answer: 53 g

27. Molecules in 2 moles of water


2 × 6.022 × 10^{23} = 1.204 × 10^{24}

Answer: molecules

28. Percentage of Oxygen in CaCO₃

Molar mass:

Ca = 40
C = 12
O₃ = 48

Total = 100


\%O = \frac{48}{100} × 100 = 48\%

Answer: 48%

Section F — Case Study

Initial mass = 20 g
Mass of CaO = 11.2 g

a) Mass of CO₂


20 - 11.2 = 8.8 \text{ g}

b) Law verified: Law of Conservation of Mass

c) Importance: ensures mass remains constant in reactions and supports accurate calculations.

Section G — MCQs

29. B ✔
30. C ✔

Additional Sample Question Paper💫

Some Basic Concepts of Chemistry (Class 11)

Time: 3 Hours
Maximum Marks: 70

Section A — Very Short Answer Questions (1 Mark Each)

Define matter.
Write the SI unit of mass.
What is Avogadro’s number?
Define atomic mass.
What is the molar mass of oxygen?
State the law of multiple proportions.
What is the empirical formula of hydrogen peroxide?
Define limiting reagent.
What is the number of significant figures in 0.003405?
Write the SI unit of temperature.

Section B — Short Answer Questions (2 Marks Each)

Differentiate between a compound and a mixture.
Give an example of a chemical property and a physical property.
Convert 5 × 10⁻⁷ m into nanometers.
Write the steps to calculate percentage composition.
State the law of conservation of mass with one example.

Section C — Short Answer Questions (3 Marks Each)

Explain the mole concept with one example.
Write the empirical formula of a compound containing 40% C, 6.7% H, and 53.3% O.
Calculate the molecular mass of HNO₃. (H = 1, N = 14, O = 16)
What are significant figures? Give two rules.
Distinguish between empirical and molecular formula with examples.

Section D — Long Answer Questions (5 Marks Each)

Explain the three laws of chemical combination.
Explain the SI units used in chemistry and their importance.
A compound contains 36% C, 6% H, and 58% O. Determine its empirical formula.
Explain limiting reagent with a numerical example.
Write the steps to determine the molecular formula from the empirical formula.

Section E — Numerical Problems (5 Marks Each)

Calculate the mass of 0.25 mole of NaOH (Na = 23, O = 16, H = 1).
How many molecules are present in 0.5 mole of CO₂?
Calculate the percentage of hydrogen in CH₄.
5 g of H₂ reacts with 40 g O₂ to form water. Identify the limiting reagent and calculate mass of water formed.
Calculate the moles of sulfur in 32 g of sulfur atoms.

Section F — Case-Based Questions (5 Marks Each)

Case 1:
A student heats 25 g of magnesium carbonate and obtains magnesium oxide and carbon dioxide gas.

a) Calculate the mass of CO₂ formed if magnesium oxide weighs 15 g.
b) Name the law verified.
c) Why is this law important in chemistry?

Case 2:
A sample contains 12 g of carbon and 32 g of oxygen.

a) Find the empirical formula.
b) If molecular mass of the compound is 44 g/mol, find the molecular formula.
c) Explain the difference between empirical and molecular formula.

Section G — Multiple Choice Questions (1 Mark Each)

One mole of any substance contains:
A)
B)
C)
D)
The molecular mass of H₂SO₄ is:
A) 98
B) 100
C) 96
D) 92
The empirical formula of glucose (C₆H₁₂O₆) is:
A) CH
B) CH₂O
C) C₆H₁₂O₆
D) C₂H₄O₂
The number of significant figures in 0.00780 is:
A) 2
B) 3
C) 4
D) 5
Which of the following is a chemical property?
A) Melting point
B) Density
C) Rusting ✔
D) Color
The SI unit of amount of substance is:
A) gram
B) mole ✔
C) liter
D) kilogram
Percentage composition of oxygen in H₂O₂ is:
A) 50%
B) 53.3%
C) 66.7% ✔
D) 48%
2 moles of H₂ contain:
A) molecules
B) molecules ✔
C) molecules
D) 2 molecules
Limiting reagent is:
A) Leftover reactant
B) Reactant in excess
C) Reactant consumed first ✔
D) Product formed first
Molecular mass of CO is:
A) 12
B) 16
C) 28 ✔
D) 32

Solutions – Additional Sample Paper

Some Basic Concepts of Chemistry

Section A — Very Short Answer (1 Mark Each)

1. Matter is anything that has mass and occupies space.

2. Kilogram (kg)

3. Avogadro’s number =

4. Atomic mass is the relative mass of an atom compared to 1/12 of carbon-12 atom.

5. Molar mass of oxygen = 32 g/mol

6. Law of Multiple Proportions: When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in simple whole-number ratios.

7. Hydrogen peroxide (H₂O₂) → Empirical formula: HO

8. Limiting reagent is the reactant that is completely consumed first in a chemical reaction.

9. 0.003405 → 4 significant figures (3, 4, 0, 5)

10. Kelvin (K)

Section B — Short Answer (2 Marks Each)

11. Compound vs Mixture

Compound	Mixture
Chemically combined	Physically combined
Fixed composition	Variable composition
Properties differ from elements	Retain properties of constituents

12. Examples:

Chemical property: rusting of iron
Physical property: density of water

13. Convert 5 × 10⁻⁷ m to nm


1 \text{ m } = 10^9 \text{ nm}  
5 × 10^{-7} × 10^9 = 500 \text{ nm}

14. Steps to calculate % composition:

Find molar mass of compound
Divide mass of each element by molar mass
Multiply by 100

15. Law of Conservation of Mass Example:


2H_2 + O_2 → 2H_2O  
\text{Mass before = Mass after}

Proposed by Antoine Lavoisier

Section C — Short Answer (3 Marks Each)

16. Mole Concept

1 mole = particles.
Example: 1 mole H₂O = 18 g contains molecules.

17. Empirical formula of C, H, O

C = 40 g → 40/12 = 3.33
H = 6.7 g → 6.7/1 = 6.7
O = 53.3 g → 53.3/16 = 3.33

Divide by smallest (3.33):
C = 1, H = 2, O = 1 → Empirical formula = CH₂O

18. Molecular Mass of HNO₃

H = 1, N = 14, O₃ = 48

Total = 1 + 14 + 48 = 63 u

19. Significant figures rules:

Non-zero digits are significant
Leading zeros are not significant
Trailing zeros after decimal are significant
Zeros between non-zero digits are significant

20. Empirical vs Molecular formula

Empirical Formula	Molecular Formula
Simplest ratio	Actual number of atoms
Example: CH₂O	Example: C₆H₁₂O₆

Section D — Long Answer (5 Marks Each)

21. Laws of Chemical Combination

Law of Conservation of Mass – Mass is neither created nor destroyed.
Law of Definite Proportions – Compounds contain elements in fixed ratio.
Law of Multiple Proportions – Masses combine in small whole numbers.

22. SI Units

Quantity	Unit
Length	meter (m)
Mass	kilogram (kg)
Time	second (s)
Temperature	kelvin (K)
Amount	mole (mol)

Importance: Uniformity in measurements, accurate calculations, scientific communication.

23. Empirical formula calculation

Given: C = 36%, H = 6%, O = 58%

C: 36/12 = 3
H: 6/1 = 6
O: 58/16 = 3.625 ≈ 3.625 → divide by 3 → 1, 2, 1.2 → multiply by 5 → 5, 10, 6 → Empirical formula = C₅H₁₀O₆ (approx)

24. Limiting reagent example


2H_2 + O_2 → 2H_2O

If 2 moles H₂ and 1 mole O₂ react: H₂ is limiting reagent, 2 moles water formed.

25. Molecular formula steps

Determine empirical formula
Find empirical formula mass
Use: n = molecular mass / empirical mass
Multiply subscripts in empirical formula by n

Section E — Numericals

26. Mass of 0.25 mole NaOH

Molar mass: Na = 23, O = 16, H = 1 → 40 g/mol

0.25 × 40 = 10 g

27. Molecules in 0.5 mole CO₂

0.5 × = 3.011 × 10²³ molecules

28. % H in CH₄

H = 4 × 1 = 4
Molar mass = 12 + 4 = 16


\% H = \frac{4}{16} × 100 = 25\%

29. Limiting reagent and water formed

Reaction: 5 g H₂ + 40 g O₂ → H₂O

Moles H₂ = 5/2 = 2.5 moles
Moles O₂ = 40/32 = 1.25 moles

Reaction ratio: 2:1 → H₂ needed = 2 × 1.25 = 2.5 → exact
Limiting reagent = H₂
Water formed = 2 × 1.25 = 2.5 moles → 2.5 × 18 = 45 g

30. Moles of sulfur in 32 g S

Atomic mass S = 32 g/mol


\text{Moles} = \frac{32}{32} = 1 \text{ mole}

Section F — Case-Based Questions

Case 1:

Initial mass = 25 g, MgO = 15 g → CO₂ mass = 25 – 15 = 10 g
Law verified: Law of Conservation of Mass
Importance: Mass remains constant in reactions, essential for calculations

Case 2:

C = 12 g, O = 32 g → Ratio = 12:32 → 3:8 → Empirical formula = CO₂

Molecular mass = 44 g/mol, empirical mass = 12+32=44 → Molecular formula = CO₂

Difference: