15 short questions for 2 to 3 marks based on the "Some Basic Concepts of Chemistry" chapter
1.Differentiate between homogeneous and heterogeneous mixtures with one example each?
Homogeneous mixtures have a uniform composition throughout. For example, a saltwater solution where salt is evenly dissolved in water. Heterogeneous mixtures have a non-uniform composition with visible boundaries between the components. For example, a mixture of sand and water where the sand particles are distinct from the water.
2. Explain the law of definite proportions with a suitable example?
The law of definite proportions states that a given compound always contains exactly the same proportion of elements by weight. For example, in water (H2O), the ratio of the mass of hydrogen to the mass of oxygen is always 1:8, regardless of the source of water.
3. State the law of multiple proportions?
How does it differ from the law of definite proportions? The law of multiple proportions states that if two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in the ratio of small whole numbers. The law of definite proportions deals with the constant ratio of elements in a single compound, while the law of multiple proportions deals with the ratio of masses of one element that combine with a fixed mass of another element in different compounds.
4. Define the terms 'mole' and 'Avogadro's number'. What is the relationship between them?
A mole is the amount of a substance that contains as many elementary entities (atoms, molecules, ions, etc.) as there are atoms in exactly 12 grams of the carbon-12 isotope. Avogadro's number (NA) is the number of elementary entities in one mole of a substance, which is approximately 6.022×1023. The relationship is that one mole contains Avogadro's number of entities.
5. Calculate the number of moles present in 49 g of sulfuric acid (H2SO4). (Molar mass of H2SO4 = 98 g/mol) Number of moles = Given mass / Molar mass = 49 g / 98 g/mol = 0.5 moles.
6. What is the difference between empirical and molecular formulas?
If the empirical formula of a compound is CH2O and its molar mass is 90 g/mol, what is its molecular formula? The empirical formula is the simplest whole-number ratio of atoms in a compound, while the molecular formula shows the actual number of atoms of each element in a molecule. To find the molecular formula, we first calculate the empirical formula mass (12 + 2(1) + 16 = 30 g/mol). Then, we find the ratio n = Molar mass / Empirical formula mass = 90 g/mol / 30 g/mol = 3. The molecular formula is then (Empirical formula)n = (CH2O)3=C3H6O3.
7. Define 'limiting reactant' and 'excess reactant' in a chemical reaction. How does the limiting reactant govern the amount of product formed?
The limiting reactant is the reactant that is completely consumed in a chemical reaction. The excess reactant is the reactant that is present in a greater amount than required for complete reaction with the limiting reactant. The limiting reactant determines the maximum amount of product that can be formed because once it is used up, the reaction stops, regardless of how much excess reactant is still present.
8. In the reaction 2H2(g)+O2(g)→2H2O(l), if 4 g of hydrogen reacts with 32 g of oxygen, identify the limiting reactant.
Molar mass of H2 = 2 g/mol, so 4 g of H2 is 4 g / 2 g/mol = 2 moles. Molar mass of O2 = 32 g/mol, so 32 g of O2 is 32 g / 32 g/mol = 1 mole. From the balanced equation, 2 moles of H2 react with 1 mole of O2. Here, we have exactly the required stoichiometric ratio. Therefore, neither reactant is limiting, and both are completely consumed.
9. What is the molar volume of a gas at STP? Calculate the volume occupied by 2 moles of nitrogen gas at STP.
The molar volume of a gas at STP (Standard Temperature and Pressure) is 22.4 liters per mole. The volume occupied by 2 moles of nitrogen gas at STP would be 2 moles × 22.4 L/mol = 44.8 liters.
10. Differentiate between molarity (M) and molality (m). Which concentration term is preferred when the temperature changes and why?
Molarity (M) is defined as the number of moles of solute per liter of solution. Molality (m) is defined as the number of moles of solute per kilogram of solvent. Molality is preferred when the temperature changes because it is based on the mass of the solvent, which does not change with temperature, unlike the volume of the solution which can expand or contract with temperature variations.
11. A solution is prepared by dissolving 5.85 g of sodium chloride (NaCl) in 100 mL of water. Calculate the molarity of the solution. (Molar mass of NaCl = 58.5 g/mol)
Number of moles of NaCl = 5.85 g / 58.5 g/mol = 0.1 moles. Volume of solution = 100 mL = 0.1 L. Molarity = Number of moles of solute / Volume of solution (in L) = 0.1 moles / 0.1 L = 1 M.
12. State the rules for determining significant figures in a measurement. How many significant figures are there in the following numbers: (a) 0.0230 (b) 245 (c) 100?
Rules for significant figures include: 1. All non-zero digits are significant. 2. Zeros between non-zero digits are significant. 3. Leading zeros are not significant. 4. Trailing zeros in a number containing a decimal point are significant. 5. Trailing zeros in a number not containing a decimal point may or may not be significant (ambiguous without more context). (a) 0.0230 has 3 significant figures (2, 3, and the trailing zero after the decimal). (b) 245 has 3 significant figures. (c) 100 has 1 significant figure (the '1'). The trailing zeros are ambiguous without a decimal point.
13. Explain the terms 'precision' and 'accuracy' with the help of an example.
Precision refers to the closeness of several measurements of the same quantity to one another. Accuracy refers to the closeness of a single measurement to the true or accepted value. For example, if the true weight of an object is 10.00 g, and three measurements are 9.98 g, 10.01 g, and 9.99 g, these measurements are precise (close to each other) but might not be perfectly accurate if the average (9.993 g) is slightly off from the true value. Conversely, a measurement of 10.02 g could be more accurate than the individual precise measurements, even if less precise on its own.
14. What is the importance of stoichiometry in chemical calculations? Explain briefly.
Stoichiometry is crucial in chemical calculations because it provides the quantitative relationships between the amounts of reactants and products in a balanced chemical equation. It allows us to predict the amount of product formed from a given amount of reactants, determine the amount of reactants needed to obtain a desired amount of product, and identify the limiting reactant in a reaction. This is essential for efficient and safe execution of chemical reactions in laboratories and industrial processes.
15. Define 'pure substance'. Classify pure substances into elements and compounds, giving one example for each?
A pure substance is a substance having a fixed or constant chemical composition and distinct properties. Pure substances can be classified into two types: Elements and Compounds. An element is a substance that cannot be broken down into simpler substances by chemical means. Example: Oxygen (O). A compound is a substance formed when two or more elements are chemically combined in a fixed ratio. Example: Water (H2O).
