Lesson 12 of 12 · 16 min

Chapter review

Watch a class

The whole chapter on YouTube

Complete first chapter revision in one go
PW Class 11 Science · Hinglish · Whole chapter · Open on YouTube

Chapter with equations and past questions
PW Class 11 Science · Hinglish · Whole chapter · Open on YouTube

Must-know facts

18 facts

1Seven SI base units: m, kg, s, A, K, mol, cd.
2K = °C + 273.15; °F = (9/5)(°C) + 32.
31 L = 1000 mL = 1000 cm³ = 1 dm³.
4Leading zeros are never significant; trailing zeros are significant only when a decimal point is present.
5Multiplication/division: fewest significant figures; addition/subtraction: fewest decimal places.
6Conservation of mass: Lavoisier; definite proportions: Proust; multiple proportions: Dalton; gaseous volumes: Gay Lussac; equal volumes, equal molecules: Avogadro.
7Dalton's theory could not explain Gay Lussac's law of gaseous volumes.
81 u = 1/12 of the mass of a carbon-12 atom = 1.66056 × 10⁻²⁴ g.
9Nᴀ = 6.022 × 10²³ mol⁻¹.
10Molar mass in g mol⁻¹ is numerically equal to the molecular or formula mass in u.
11Ionic compounds have formula mass, not molecular mass (NaCl = 58.5 u).
12One mole of an ideal gas occupies 22.7 L at 273.15 K and 1 bar.
13Empirical formula of glucose is CH₂O.
14Molecular formula = n × empirical formula, with n = molar mass / empirical formula mass.
15The limiting reagent decides the amount of product, not the reactant present in larger mass.
16Molality is temperature-independent; molarity is not.
17Sum of mole fractions of all components = 1.
18Dilution: M₁V₁ = M₂V₂.

Common traps

Where marks are lost

Counting molecules when the question asks for atoms of one element.

Multiply moles of molecules by Nᴀ and then by the number of that atom per formula; 0.5 mol NH₃ holds 1.5 mol of H atoms.

Treating the reactant with the larger mass as the one in excess.

Compare moles divided by coefficients, never raw masses; the smallest ratio identifies the limiting reagent.

Using mass of solution instead of mass of solvent when computing molality.

Molality is per kg of solvent; subtract the solute's mass from the solution's mass first.

Using volume of solvent instead of volume of solution for molarity.

Molarity divides by litres of the whole solution, obtained from its mass and density when needed.

Counting leading zeros as significant (e.g. saying 0.0050600 has eight).

Leading zeros only locate the decimal point; 0.0050600 has five significant figures (5, 0, 6, 0, 0).

Calling a ratio of 1 : 2 masses in two different compounds the law of definite proportions.

Fixed mass ratio within one compound is definite proportions; a whole-number ratio between compounds of the same two elements is multiple proportions.

Writing molecular mass for NaCl.

NaCl has no discrete molecules in the solid; the correct term is formula mass.

Using 22.4 L per mole at STP without checking the pressure convention.

22.7 L is the molar volume at 273.15 K and 1 bar, the value NCERT uses; 22.4 L is for 273.15 K and 1 atm. Use the conditions the question states.

Stopping at the empirical formula when molar mass or vapour density is given.

Divide molar mass by empirical formula mass to find n, then scale the formula.

Formulas

11 to know

Kelvin and Fahrenheit conversion

K = °C + 273.15 ; °F = (9/5)(°C) + 32

No negative temperatures exist on the kelvin scale.

Density

d = m / V

SI unit kg m⁻³; commonly g cm⁻³ or g mL⁻¹.

Number of moles

n = mass / molar mass = N / Nᴀ

Nᴀ = 6.022 × 10²³ mol⁻¹; molar mass in g mol⁻¹.

Moles of a gas at STP

n = V / 22.7 L

STP taken as 273.15 K and 1 bar.

Average atomic mass

Aᵥ = Σ (fractional abundance × isotopic mass)

Explains non-integer atomic masses such as 12.011 u for carbon.

Mass percent of an element

% element = (mass of element in 1 mol compound / molar mass) × 100

Percentages of all elements add up to 100.

Molecular formula multiplier

n = molar mass / empirical formula mass

Background, not in this NCERT chapter: molar mass ≈ 2 × vapour density (relative to H₂).

Mole fraction

x_A = n_A / (n_A + n_B + ...)

Dimensionless; all fractions sum to 1.

Molarity

M = moles of solute / volume of solution in L

Unit mol L⁻¹; temperature dependent.

Molality

m = moles of solute / mass of solvent in kg

Unit mol kg⁻¹; temperature independent.

Dilution

M₁V₁ = M₂V₂

Moles of solute stay constant on dilution.

Key terms

13 terms

Homogeneous mixture: A mixture whose composition is uniform throughout, like a salt solution.
Compound: A pure substance made of two or more elements in a fixed mass ratio, separable only by chemical means.
Precision: Agreement of repeated measurements with each other.
Accuracy: Agreement of a measurement with the true value.
Significant figures: The certain digits of a measurement plus one estimated digit.
Unified mass (u): One-twelfth of the mass of a carbon-12 atom, used to express atomic and molecular masses.
Formula mass: The sum of atomic masses in the formula unit of an ionic compound.
Mole: The amount of substance holding 6.022 × 10²³ specified entities.
Molar mass: Mass in grams of one mole of a substance.
Empirical formula: Formula showing the simplest whole-number atom ratio.
Limiting reagent: The reactant consumed first, which caps the amount of product.
Molality: Moles of solute per kilogram of solvent.
Molarity: Moles of solute per litre of solution.

Test yourself: 9 questionsExam-style questions on Some Basic Concepts of Chemistry, with full solutions.Start