Lesson 12 of 12 · 16 min
Chapter review
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Must-know facts
18 facts
- 01Seven SI base units: m, kg, s, A, K, mol, cd.
- 02K = °C + 273.15; °F = (9/5)(°C) + 32.
- 031 L = 1000 mL = 1000 cm³ = 1 dm³.
- 04Leading zeros are never significant; trailing zeros are significant only when a decimal point is present.
- 05Multiplication/division: fewest significant figures; addition/subtraction: fewest decimal places.
- 06Conservation of mass: Lavoisier; definite proportions: Proust; multiple proportions: Dalton; gaseous volumes: Gay Lussac; equal volumes, equal molecules: Avogadro.
- 07Dalton's theory could not explain Gay Lussac's law of gaseous volumes.
- 081 u = 1/12 of the mass of a carbon-12 atom = 1.66056 × 10⁻²⁴ g.
- 09Nᴀ = 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
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.