Some Basic Concepts of Chemistry

Q1. What is the importance of chemistry in agriculture and food?

Chemistry plays a crucial role in agriculture and food by:

• Providing chemical fertilizers like urea, calcium phosphate, and ammonium phosphate to enhance crop yield.
• Developing insecticides, fungicides, and pesticides to protect crops from harmful bacteria and insects.
• Using preservatives to extend the shelf life of food products such as jam, butter, and squashes.

Q2. How does chemistry contribute to health and sanitation?

Chemistry aids health and sanitation by:

• Producing life-saving drugs like penicillin, sulpha drugs (for dysentery and pneumonia), cisplatin (for cancer), and AZT (for AIDS).
• Supplying disinfectants like phenol to kill microorganisms in drains and toilets.
• Using chlorine (0.2–0.4 ppm) to sterilize drinking water.

Q3. What role does chemistry play in environmental conservation?

Chemistry helps save the environment by:

• Developing eco-friendly substitutes like CNG (Compressed Natural Gas) to reduce pollution from automobiles.
• Researching alternatives to industrial chemicals that cause pollution.

Q4. How is chemistry applied in industries?

Industrial applications of chemistry include:

• Manufacturing fertilizers, acids, alkalis, dyes, polymers, soaps, and detergents.
• Producing metal alloys and new materials that boost the economy.

Q5. Define matter and give examples.

Matter is anything that has mass and occupies space (e.g., book, water, air).

Q6. How is matter classified physically?

Matter exists in three physical states:

1. Solids: Definite shape and volume; particles are tightly packed.
2. Liquids: Definite volume but no fixed shape; particles can move.
3. Gases: No definite shape or volume; particles are far apart.

Q7. What is the chemical classification of matter?

Matter is chemically classified into:

1. Pure substances: Elements (e.g., sodium) and compounds (e.g., water).
2. Mixtures: Homogeneous (uniform composition, e.g., sugar water) and heterogeneous (non-uniform, e.g., iron + sulphur).

Q8. Differentiate between compounds and mixtures.

Compounds:

• Chemically combined elements in fixed ratios.
• Homogeneous; constituents lose their properties.
• Cannot be separated physically.
  Mixtures:
• Physically mixed; variable ratios.
• May be homogeneous or heterogeneous.
• Constituents retain properties; separable physically.

Q9. What are physical and chemical properties?

Physical properties: Observable without changing composition (e.g., color, melting point).
Chemical properties: Involve chemical changes (e.g., acidity, combustibility).

Q10. List the seven basic SI units.

The SI units are:

1. Metre (length)
2. Kilogram (mass)
3. Second (time)
4. Kelvin (temperature)
5. Ampere (current)
6. Candela (luminous intensity)
7. Mole (amount of substance).

Q11. Define the SI unit ‘mole’.

A mole is the amount of substance containing as many entities (atoms, molecules) as there are in 12g of carbon-12 (6.022 × 10²³ entities).

Q12. Differentiate between mass and weight.

Mass: Constant quantity of matter (SI unit: kg).
Weight: Force due to gravity (varies with location).

Q13. How is volume measured in chemistry?

Volume is measured in:

• SI unit: m³.
• Common unit: Litre (1L = 1000 cm³ = 1 dm³).
• Tools: Burette, pipette, measuring flask.

Q14. Convert 25°C to Kelvin and Fahrenheit.

• Kelvin: 25 + 273.15 = 298.15 K.
• Fahrenheit: (25 × 9/5) + 32 = 77°F.

Q15. What is density? Give its SI unit.

Density = mass/volume. SI unit: kg/m³; common unit: g/cm³.

Q16. Explain scientific notation with an example.

Scientific notation expresses numbers as N × 10ⁿ (1 ≤ N < 10).
Example: 232.508 = 2.32508 × 10².

Q17. What are significant figures? State the rules.

Significant figures are certain digits in a measurement. Rules:

1. All Non-zero digits are significant.

2. Zeros preceding to first non-zero digit are not significant. Such zero indicates the position of decimal point. [ 0.03 has 1 significant figure ]

3. Zeros between non-zero digits are significant.

4. Trailing zeros after a decimal are significant.

Q18. Add 3.52, 2.3, and 6.24 with correct significant figures.

Sum = 12.06 → Reported as 12.1 (least decimal places: 1).

Q19. Multiply 2.2120 by 0.011 with correct significant figures.

Product = 0.024332 → Reported as 0.024 (2 significant figures in 0.011).

Q20. What is dimensional analysis?

A method to convert units using conversion factors (e.g., 1 km = 1000 m).

Q21. State the Law of Conservation of Mass.

Mass is neither created nor destroyed in a chemical reaction (total mass of reactants = total mass of products).

Q22. Explain the Law of Definite Proportions.

A compound always contains elements in a fixed ratio by mass (e.g., H₂O is always 1:8 by mass).

Q23. What is the Law of Multiple Proportions?

When two elements form multiple compounds, the mass ratios are simple whole numbers (e.g., CO and CO₂: 12:16 and 12:32).

Q24. State Gay Lussac’s Law of Gaseous Volumes.

Gases combine in simple whole-number volume ratios under the same conditions (e.g., 2H₂ + O₂ → 2H₂O; ratio 2:1:2).

Q25. What is Avogadro’s Law?

Equal volumes of gases at the same temperature and pressure contain equal molecules.

Q26. Summarize Dalton’s Atomic Theory.

Key points:

1. Matter consists of indivisible atoms.
2. Atoms of an element are identical.
3. Compounds form in fixed ratios.
4. Atoms rearrange in reactions; not created/destroyed.

Q27. Define atomic mass and atomic mass unit (amu).

Atomic mass: Relative mass of an atom compared to carbon-12 (12u).
1 amu = 1/12th mass of a carbon-12 atom.

Q28. What is average atomic mass?

The weighted mean of atomic masses of an element’s isotopes (e.g., chlorine: 35.5u).

Q29. Calculate the molecular mass of CH₄.

CH₄ = C + 4H = 12.011u + 4(1.008u) = 16.043u.

Q30. What is formula mass? Give an example.

Formula mass is the sum of atomic masses in an ionic compound (e.g., NaCl = 23u + 35.5u = 58.5u).

Q31. Define Avogadro’s number.

Avogadro’s number (NA) = 6.022 × 10²³ entities per mole.

Q32. Calculate the percentage composition of hydrogen in H₂O.

%H = (2 × 1.008 / 18.016) × 100 = 11.19%.

Q33. What is an empirical formula? Give an example.

The simplest whole-number ratio of atoms (e.g., H₂O₂ → HO).

Q34. How is molecular formula related to empirical formula?

Molecular formula = (Empirical formula) × n, where n is an integer (e.g., C₆H₁₂O₆ = (CH₂O) × 6).

Q35. What is stoichiometry?

The calculation of reactant/product masses in a chemical reaction (e.g., CH₄ + 2O₂ → CO₂ + 2H₂O).

Q36. Define limiting reactant.

The reactant completely consumed in a reaction, limiting the product amount.

Q37. What is molarity? Give its formula.

Moles of solute per litre of solution. Formula:
Molarity (M) = moles of solute / volume (L).

Q38. Define molality.

Moles of solute per kilogram of solvent. Formula:
Molality (m) = moles of solute / mass of solvent (kg).

Q39. What is mass percent?

Mass of solute per 100g of solution:
Mass % = (mass of solute / mass of solution) × 100.

Q40. Explain mole fraction.

Ratio of moles of a component to total moles in solution:
Mole fraction (X) = n₁ / (n₁ + n₂).

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