Dalton's Atomic Theory and Atomic & Molecular Masses

Dalton turned the combination laws into a model of atoms; from it come atomic and molecular masses that feed the mole concept. This topic supplies definitions and small calculations worth easy marks in CBSE and NEET.

Core ideas

Dalton's atomic theory (1808) postulates:

• Matter consists of indivisible atoms.
• All atoms of a given element are identical in mass and properties.
• Compounds form when atoms of different elements combine in fixed whole-number ratios.
• Chemical reactions involve rearrangement of atoms; atoms are neither created nor destroyed.

This explains conservation of mass, definite and multiple proportions. Limitation: atoms are divisible (electrons, protons, neutrons) and isotopes show atoms of an element can differ in mass.

Atomic mass unit: 1 amu (now u or dalton) = 1/12 the mass of one carbon-12 atom = 1.66056 × 10⁻²⁴ g.

Average atomic mass = Σ(isotopic mass × fractional abundance). Example: chlorine = (34.97×0.7577)+(36.97×0.2423) ≈ 35.45 u.

Molecular mass = sum of atomic masses of all atoms in a molecule (H₂O = 2×1 + 16 = 18 u).

Formula mass is used for ionic compounds that have no discrete molecules (NaCl formula mass = 23 + 35.5 = 58.5 u).

Exam Tricks & Tips

• 🎯 Average atomic mass is a weighted mean — multiply each isotope mass by its fractional (not percentage) abundance before adding.
• 🎯 Use "molecular mass" for covalent molecules and "formula mass" for ionic solids like NaCl.
• 🎯 The carbon-12 scale fixes 1 u; a species' mass in u is numerically equal to its molar mass in g/mol.
• 🎯 Dalton's theory fails on isotopes (atoms of one element with different masses) — a common assertion–reason point.
• 🎯 Learn atomic masses H=1, C=12, N=14, O=16, Na=23, Cl=35.5, S=32 for fast molecular-mass sums.
• ❌ Common mistake: adding isotope masses directly without weighting by abundance when computing average atomic mass.

Expected exam pattern

Expect an average-atomic-mass numerical from isotopic abundances, a molecular-mass calculation, and an MCQ on a postulate/limitation of Dalton's theory.

Quick recap

Dalton: indivisible atoms, identical per element, fixed ratios in compounds. 1 u = 1/12 of C-12. Average atomic mass = weighted mean of isotopes; molecular/formula mass = sum of atomic masses.

Cover the answer, recall, then check. 12 cards on Dalton's theory and atomic/molecular masses.

Q1. In what year did Dalton propose his atomic theory?
A1. 1808.

Q2. State two key postulates of Dalton's atomic theory.
A2. Matter is made of indivisible atoms; all atoms of an element are identical in mass and properties.

Q3. How does Dalton's theory explain conservation of mass?
A3. Atoms are only rearranged in reactions, neither created nor destroyed, so mass is conserved.

Q4. Give one limitation of Dalton's atomic theory.
A4. Atoms are divisible (into electrons, protons, neutrons) and isotopes show atoms of one element can differ in mass.

Q5. Define one atomic mass unit (u/dalton).
A5. Exactly 1/12 the mass of one carbon-12 atom (≈ 1.66 × 10⁻²⁴ g).

Q6. How is average atomic mass calculated?
A6. Sum of (isotopic mass × fractional abundance) over all isotopes.

Q7. Why is chlorine's atomic mass 35.5 and not a whole number?
A7. It is a weighted average of Cl-35 (75.77%) and Cl-37 (24.23%).

Q8. Define molecular mass.
A8. The sum of the atomic masses of all atoms in one molecule.

Q9. Calculate the molecular mass of H₂O.
A9. 2(1) + 16 = 18 u.

Q10. When do we use "formula mass" instead of "molecular mass"?
A10. For ionic compounds without discrete molecules (e.g. NaCl = 58.5 u).

Q11. Which element's atom is the standard for atomic masses?
A11. Carbon-12 (assigned exactly 12 u).

Q12. Molecular mass of CO₂?
A12. 12 + 2(16) = 44 u.

In 1808 John Dalton turned the laws of chemical combination into a picture of reality: matter is built from tiny indivisible atoms. That model, refined but never discarded, lets us assign a mass to every atom and molecule — the numbers you plug into every mole calculation.

Definition: Dalton's atomic theory proposes that all matter is composed of atoms — the smallest indivisible particles of an element — which combine in whole-number ratios to form compounds.

Beginner — Dalton's postulates

1. Matter consists of tiny indivisible particles called atoms.
2. Atoms of a given element are identical in mass and properties; atoms of different elements differ.
3. Compounds form when atoms of different elements combine in fixed whole-number ratios.
4. Atoms are neither created nor destroyed in a chemical reaction (only rearranged).

This neatly explains conservation of mass and definite/multiple proportions.

Intermediate — atomic mass and the unified mass unit

Atoms are far too light to weigh individually, so we use relative masses. The standard is carbon-12: one atom of ¹²C is assigned exactly 12 units.

1 unified atomic mass unit (u) = 1/12 the mass of one ¹²C atom ≈ 1.66 × 10⁻²⁴ g.

Average atomic mass accounts for isotopes — it is the weighted average of isotopic masses by natural abundance:

Average atomic mass = Σ (isotopic mass × fractional abundance)

Advanced — molecular, formula, and average masses

• Molecular mass = sum of atomic masses of all atoms in a molecule (e.g. H₂O = 2(1.008) + 16.00 = 18.02 u).
• Formula mass is used for ionic compounds that have no discrete molecule (e.g. NaCl = 23 + 35.5 = 58.5 u).

Modern limits of Dalton's theory: atoms are divisible (electrons, protons, neutrons), and isotopes mean atoms of one element can differ in mass — but the core idea of combining whole-number atomic ratios still holds.

Worked example

Chlorine occurs as ³⁵Cl (mass 34.97 u, 75.77%) and ³⁷Cl (mass 36.97 u, 24.23%). Find its average atomic mass.

Average = (34.97 × 0.7577) + (36.97 × 0.2423)
= 26.49 + 8.96 = 35.45 u.

That is why the periodic table lists chlorine as 35.45, not a whole number — it is a weighted isotopic average.

Real-world / exam application

Mass spectrometry measures isotopic abundances exactly this way. In exams, "average atomic mass from isotopic data" and "molecular mass of a compound" are guaranteed one-markers, and they feed directly into mole and stoichiometry problems.

Exam tricks & mnemonic

Convert percentages to fractions before multiplying (75.77% → 0.7577). For molecular mass, list each element × its subscript to avoid missing an atom. Mnemonic for postulates "I Identical Combine Conserve": Indivisible, Identical, Combine in ratios, Conserved.

Key formulas

• Atomic mass unit: 1 u = 1/12 of a ¹²C atom = 1.66 × 10⁻²⁴ g.
• Average atomic mass: Σ(isotopic mass × fractional abundance).
• Molecular mass: sum of atomic masses of all atoms in a molecule.
• Formula mass: used for ionic compounds (no discrete molecules).
• Dalton's postulates: atoms are indivisible; atoms of an element are identical; compounds form in fixed whole-number ratios.

Example: chlorine (³⁵Cl 75%, ³⁷Cl 25%) → average mass = 35(0.75) + 37(0.25) = 35.5 u.

Worked Example

Problem: Chlorine occurs as two isotopes: Cl-35 (relative mass 35, abundance 75%) and Cl-37 (relative mass 37, abundance 25%). (a) Find the average atomic mass of chlorine. (b) Find the molecular mass of H₂SO₄. (Atomic masses: H = 1, S = 32, O = 16.)

Solution:
Step 1 — Average atomic mass is the abundance-weighted mean of the isotopic masses:
Average = (35 × 0.75) + (37 × 0.25).

Step 2 — Evaluate:
= 26.25 + 9.25 = 35.5 u.

Step 3 — For H₂SO₄, add the atomic masses of all atoms in the formula:
H₂SO₄ = 2(H) + 1(S) + 4(O) = 2(1) + 32 + 4(16).

Step 4 — Compute:
= 2 + 32 + 64 = 98 u.

Answer: (a) Average atomic mass of chlorine = 35.5 u. (b) Molecular mass of H₂SO₄ = 98 u.