Internal Exam Oriented Chemistry | Semester 1 | Questions and Answers

Sreehari K August 20, 2025

Internal Exam Oriented Chemistry | Semester 1 | Questions and Answers

 Questions


1. What is the difference between molecular mass and formula mass?

2. Calculate the molecular mass of sulphuric acid and calcium

carbonate.

3. Define isotopes with an example. How are isobars different from

isotopes?

4. What are isotones? Give two examples of isotones.

5. Define mole and Avogadro's number.

6. What is the molar volume of an ideal gas at STP?

7. How many moles are present in 88 g of CO₂?

8. Calculate the number of molecules in 2 moles of water.

9. Define oxidation and reduction in terms of electron transfer.

10. What is oxidation number? Give an example.

11. Differentiate between valency and oxidation number.

12. State the oxidation number of sulfur in H₂SO₄.

13. What is the variable valency shown by iron?

14. Define equivalent mass. Calculate the equivalent mass of sulphuric

acid and sodium hydroxide.

15. Define molality and molarity.

Answers

Sure! Here are the answers to your questions:

1. Difference between molecular mass and formula mass

  • Molecular mass: Sum of atomic masses of all atoms present in a molecule (used for covalent compounds).

  • Formula mass: Sum of atomic masses of atoms in the formula unit of an ionic compound (used for ionic compounds which don’t exist as discrete molecules).

2. Molecular mass calculations

  • Sulphuric acid (H₂SO₄):
    = 2×1 + 32 + 4×16 = 2 + 32 + 64 = 98 u

  • Calcium carbonate (CaCO₃):
    = 40 + 12 + 3×16 = 40 + 12 + 48 = 100 u

3. Isotopes
Atoms of the same element with same atomic number but different mass numbers.
Example:

  • ¹H, ²H, ³H are isotopes of hydrogen.

Isobars: Atoms of different elements having same mass number but different atomic numbers.
E.g., ⁴⁰Ar and ⁴⁰Ca.

4. Isotones
Atoms having same number of neutrons but different proton numbers.
Examples:

  • ¹⁴C and ¹⁵N (both have 8 neutrons)

  • ³⁶S and ³⁷Cl (both have 20 neutrons)

5. Mole & Avogadro’s number

  • Mole: Amount of substance containing as many entities as there are atoms in 12 g of C-12.

  • Avogadro’s number: 6.022 × 10²³ particles per mole.

6. Molar volume at STP
= 22.4 L for 1 mole of an ideal gas at STP (0 °C and 1 atm).

7. Moles in 88 g of CO₂
Molecular mass CO₂ = 44 g/mol
Moles = mass / molar mass = 88 / 44 = 2 moles

8. Number of molecules in 2 moles of H₂O
2 × 6.022 × 10²³ = 1.2044 × 10²⁴ molecules

9. Oxidation & reduction (electron transfer)

  • Oxidation: Loss of electrons

  • Reduction: Gain of electrons

10. Oxidation number
The apparent charge an atom appears to have when electrons are assigned according to rules.
Example: Oxidation number of Na in NaCl is +1.

11. Valency vs Oxidation number

Property Valency Oxidation Number
Meaning Combining capacity Apparent charge on atom
Positive/Negative Always positive Can be positive, negative or zero
Fixed? Usually fixed Can vary depending on compound

12. Oxidation number of S in H₂SO₄
Let x = oxidation number of S
2(+1) + x + 4(−2) = 0
2 + x – 8 = 0 → x = +6

13. Variable valency of iron
Iron shows valency +2 (ferrous) and +3 (ferric).

14. Equivalent mass
Mass of a substance that combines with or displaces 1 mole of hydrogen (1.008 g), 8 g oxygen, or 35.5 g chlorine.

  • H₂SO₄ (sulphuric acid):
    Equivalent mass = molecular mass / basicity = 98 / 2 = 49 g eq⁻¹

  • NaOH (sodium hydroxide):
    Equivalent mass = molecular mass / acidity = 40 / 1 = 40 g eq⁻¹

15. Molality & Molarity

  • Molarity (M): Moles of solute per litre of solution.

    M=moles solutevolume of solution in LM = \frac{\text{moles solute}}{\text{volume of solution in L}}

  • Molality (m): Moles of solute per kg of solvent.

    m=moles solutemass of solvent in kgm = \frac{\text{moles solute}}{\text{mass of solvent in kg}}

Questions 

 1.       
What is the difference between
molecular mass and formula mass?


2.       
Calculate the molecular mass of sulphuric
acid and calcium carbonate.


3.       
Define isotopes with an example. How are
isobars different from isotopes?


4.       
What are isotones? Give two examples of
isotones.


5.       
Define mole and Avogadro's number.


6.       
What is the molar volume of an ideal gas at
STP?


7.       
How many moles are present in 88 g of CO₂?


8.       
Calculate the number of molecules in 2 moles
of water.


9.       
Define oxidation and reduction in terms of
electron transfer.


10.   
What is oxidation number? Give an example.


11.   
Differentiate between valency and oxidation
number.


12.   
State the oxidation number of sulfur in H₂SO₄.


13.   
What is the variable valency shown by iron?


14.   
Define equivalent mass. Calculate the
equivalent mass of sulphuric acid and sodium hydroxide.


15.   
Define molality and molarity.


16.   
What is meant by a
primary standard. Give examples.


17.   
Define Normality with
equation and write the terms.


18.   
What are standard
solutions and illustrate with example specifying weight volume and
concentration?


19.   
What is molarity and
write expression to find molarity?


20.   
What is molality and
write expression to find molality?


21.   
How molarity differs
from molality?


22.   
How formality and
molarity differs each other?


23.   
Explain the importance
of mole fraction and parts per million as units for representing concentration.


24.   
What are the
characteristics of a primary standard.?


25.   
What are acid base
indicators. Give examples?


26.   
Why is oxalic acid considered a
primary standard but sodium hydroxide is not?


27.   
A 0.500 g sample of Na₂CO₃ is
dissolved and made up to 250 mL. Calculate the molarity of the solution.


28.   
Why is no external indicator
required in KMnO₄ titrations?


29.   
What are redox titrations, given an
example.


30.   
Write the basic principle of
complexometric titrations.


31.   
In EDTA titration, the endpoint is
detected using Eriochrome Black T. What colour change occurs?

Answers


Sure! Here are clear and concise answers to your chemistry questions:

1. Difference between molecular mass and formula mass:

  • Molecular mass is the mass of one molecule of a covalent compound, calculated by adding the atomic masses of atoms present.

  • Formula mass is the mass of one formula unit of an ionic compound, calculated similarly but used because ionic compounds don't exist as single molecules.

2. Molecular mass

  • H₂SO₄ = (2×1) + 32 + (4×16) = 98 u

  • CaCO₃ = 40 + 12 + (3×16) = 100 u

3. Isotopes: Atoms of the same element with the same atomic number but different mass numbers. e.g.,
¹²C and ¹⁴C.
Isobars are atoms of different elements with different atomic numbers but same mass number. e.g., ⁴⁰Ar and ⁴⁰Ca.

4. Isotones: Atoms with the same number of neutrons but different atomic and mass numbers.
Examples: ¹⁴C (6p,8n) and ¹⁵N (7p,8n), or ³⁷Cl and ³⁸Ar both have 20 neutrons.

5. Mole: Amount of substance containing exactly 6.022×10²³ particles (Avogadro’s number).
Avogadro’s number ≡ 6.022 × 10²³.

6. Molar volume of an ideal gas at STP = 22.4 L per mole.

7. Moles in 88 g CO₂:
Molar mass = 44 g/mol ⇒ Moles = 88 / 44 = 2 moles

8. Molecules in 2 moles of H₂O =
2 × 6.022×10²³ = 1.2044×10²⁴ molecules

9. Oxidation = loss of electrons.
Reduction = gain of electrons.

10. Oxidation number: Apparent charge an atom has in a compound.
E.g., Na in NaCl has oxidation number +1.

11. Valency vs Oxidation number:

  • Valency = combining capacity (usually whole number)

  • Oxidation number = apparent charge (can be +, –, zero, fractions)

12. Oxidation number of S in H₂SO₄ = +6

13. Iron shows +2 and +3 variable valency (Fe²⁺/Fe³⁺).

14. Equivalent mass = molar mass / n-factor.

  • H₂SO₄: molar mass 98, n = 2 ⇒ 49 g/eq

  • NaOH: molar mass 40, n =1 ⇒ 40 g/eq

15.

  • Molarity (M) = moles of solute / litre of solution

  • Molality (m) = moles of solute / kg of solvent

16. Primary standard: Pure, stable substance used to prepare a standard solution directly. Examples: Oxalic acid, K₂Cr₂O₇.

17. Normality (N) = (gram equivalent of solute) / litre of solution

18. Standard solution: Solution of known concentration, e.g., 0.1 N HCl or 1 M NaOH.

19. Molarity = moles of solutevolume of solution in L\dfrac{\text{moles of solute}}{\text{volume of solution in L}}

20. Molality = moles of solutemass of solvent in kg\dfrac{\text{moles of solute}}{\text{mass of solvent in kg}}

21. Molarity depends on volume (temperature-dependent). Molality depends on mass (temperature-independent).

22. Formality is used for ionic salts that don’t exist as molecules in solution—similar to molarity but for ionic species.

23. Mole fraction and ppm are used for very dilute solutions; they are temperature independent and useful in gas phase and environmental studies.

24. Characteristics of a primary standard:

  • High purity

  • Stable

  • High molar mass

  • Non-hygroscopic

  • Easily weighed

25. Acid–base indicators: Compounds that change colour with pH, e.g., Phenolphthalein, Methyl orange.

26. Oxalic acid is pure and stable ⇒ primary standard.
NaOH absorbs CO₂ and moisture ⇒ not a primary standard.

27. Molarity of Na₂CO₃:
Moles = 0.500 g / 106 g mol⁻¹ = 0.00472 mol
Volume = 0.250 L ⇒ M = 0.0189 M

28. KMnO₄ is a self-indicator due to its deep purple colour.

29. Redox titrations involve oxidation–reduction reactions.
Example: KMnO₄ vs FeSO₄ titration.

30. Complexometric titrations are based on formation of stable coloured complexes between metal ions and ligands (e.g., EDTA).

31. In EDTA titration with Eriochrome Black T, endpoint colour changes from wine red → blue.


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