Introduction to Analytical Chemistry — Class 11 MCQs with Answers

Introduction to Analytical Chemistry — important questions & MCQs with answers (Class 11 Chemistry)

4 solved questions from this chapter's difficulty levels, each with its answer and explanation. Tap an option to check yourself.

0 of 4 answered

1. Q1 of 4Easy

   One mole of any substance contains how many elementary particles?

   A6.022 × 10²²B3.011 × 10²³C6.022 × 10²³D1.008 × 10²³
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   Answer: 6.022 × 10²³

   Explanation: The Avogadro constant, 6.022 × 10²³ particles per mole, is the fixed number of particles in one mole of any substance.

2. Q2 of 4Easy

   Molarity of a solution is defined as the number of moles of solute dissolved in

   Aone kilogram of solventBone litre of solventCone kilogram of solutionDone litre of solution
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   Answer: one litre of solution

   Explanation: Molarity (M) counts moles of solute per litre of the final solution, not per litre of solvent taken.

3. Q3 of 4Medium

   The percentage of nitrogen by mass in ammonium nitrate, NH₄NO₃ (molar mass 80 g mol⁻¹), is

   A17.5 %B28.0 %C45.0 %D35.0 %
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   Answer: 35.0 %

   Explanation: NH₄NO₃ contains two N atoms, a mass of 28 g per 80 g of compound, so % N = (28 ÷ 80) × 100 = 35 %.

4. Q4 of 4Medium

   2 g of H₂ and 32 g of O₂ are allowed to react according to 2H₂ + O₂ → 2H₂O. The limiting reagent is

   AO₂BH₂Cneither, both are exactly consumedDH₂O
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   Answer: H₂

   Explanation: 2 g of H₂ is 1 mol and 32 g of O₂ is 1 mol, but 1 mol of O₂ needs 2 mol of H₂, so hydrogen runs out first.

Introduction to Analytical Chemistry notes: key concepts (Class 11 Chemistry)

Analytical chemistry answers two questions about a sample: which constituents are present (qualitative analysis) and how much of each is present (quantitative analysis). This chapter builds the counting tools that quantitative work rests on — the mole, molar mass and the Avogadro constant — and then applies them to percentage composition, empirical and molecular formulae, the standard concentration terms (molarity, molality, mole fraction, ppm, % w/w, % v/v), equivalent mass, and stoichiometric calculations that hinge on identifying the limiting reagent.

Mole

The SI unit of amount of substance. One mole contains 6.022 × 10²³ elementary entities, the same number of atoms as there are in 12 g of carbon-12, and it is the bridge between the mass a chemist weighs and the number of particles that actually react.

Molar mass

The mass of one mole of a substance, expressed in g mol⁻¹ and numerically equal to the atomic or molecular mass in unified mass units. Moles are obtained from it as n = mass ÷ molar mass, which is the first line of nearly every calculation in this chapter.

Molar volume of a gas

One mole of any ideal gas occupies 22.4 L at STP (273 K, 1 atm), regardless of which gas it is. This lets a measured gas volume be converted directly into moles, and then into mass, without weighing the gas at all.

Percentage composition

The mass of each element expressed as a percentage of the mass of the whole compound: % of element = (mass of element in one mole ÷ molar mass) × 100. It is what combustion and gravimetric analyses actually report, and it is the starting point for deriving a formula.

Empirical formula

The simplest whole-number ratio of atoms in a compound, obtained by dividing each element's mass percentage by its atomic mass and then dividing all the results by the smallest one. Benzene C₆H₆ and ethyne C₂H₂ share the empirical formula CH.

Molecular formula

The actual number of atoms of each element in one molecule. It is found from molecular formula = n × empirical formula, where n = molar mass ÷ empirical formula mass, so an independent measurement of molar mass is always needed.

Molarity (M)

Moles of solute per litre of solution, unit mol L⁻¹. Because it is defined on the volume of solution and volume expands on heating, molarity is temperature-dependent, which matters whenever a standard solution is prepared or stored at a different temperature.

Molality (m)

Moles of solute per kilogram of solvent, unit mol kg⁻¹. Since only masses appear in the definition and mass does not vary with temperature, molality is temperature-independent and is preferred for work over a range of temperatures.

Mole fraction

The moles of one component divided by the total moles of all components in the solution. It is a dimensionless ratio, and the mole fractions of all components of a solution necessarily add up to exactly one.

ppm and percentage concentrations

ppm expresses parts of solute per million parts of solution and is used for trace amounts such as fluoride in drinking water; % w/w uses masses, % v/v uses volumes of two liquids, and % w/v mixes mass of solute with volume of solution.

Equivalent mass

The mass of a substance that combines with or displaces one mole of H⁺ or OH⁻. For an acid it is molar mass ÷ basicity (H₂SO₄ gives 98 ÷ 2 = 49), and for a base it is molar mass ÷ acidity (Na₂CO₃ gives 106 ÷ 2 = 53).

Stoichiometry

The quantitative study of reactants and products using a balanced equation, whose coefficients give mole ratios — and, for gases at the same temperature and pressure, volume ratios — but never mass ratios directly. Masses must first be converted to moles.

Limiting reagent

The reactant that is completely used up first and therefore caps the amount of product. It is found by dividing the moles available of each reactant by its coefficient in the balanced equation; the smallest quotient identifies the limiting reagent.

Qualitative vs quantitative analysis

Qualitative analysis identifies which species are present in a sample, whereas quantitative analysis measures how much of each is present, typically by volumetric (titration) or gravimetric (weighing a precipitate) methods.