Atoms and Molecules: Unit Review

Introduction

This unit began with a simple question: what is matter made of? Along the way you met Dalton's atomic theory, the laws that supported it, the difference between atoms and molecules, and a way of comparing the masses of particles far too small to weigh one at a time. This review ties those threads into one connected picture. The aim is not to memorise a list, but to see how each idea depends on the one before, so that you can tackle unfamiliar questions with confidence.

Core explanation

Strand 1: atoms and Dalton's ideas. An atom is the smallest particle of an element that can take part in a chemical reaction. Atoms are extraordinarily small, around 0.1 nm across, so even a grain of salt contains roughly 10¹⁸ of them. John Dalton, in the early 1800s, proposed that each element is made of its own kind of atom, that atoms combine in simple whole-number ratios to form compounds, and that chemical reactions rearrange atoms without creating or destroying them.

Strand 2: the laws that Dalton explained. Three experimental laws fit neatly into this picture:

Law What is observed Particle explanation --- --- --- Conservation of mass Total mass is unchanged in a reaction Atoms are rearranged, not created or destroyed Constant composition A compound always has the same elements in the same proportion by mass Every particle of the compound contains the same fixed ratio of atoms Multiple proportions When two elements form more than one compound, the masses of one that combine with a fixed mass of the other are in a simple whole-number ratio Atoms combine only as whole units, such as CO and CO₂

Strand 3: what later science corrected. Atoms are not indivisible: they contain protons, neutrons and electrons. Atoms of one element are not all identical in mass, because isotopes exist. Some compounds, such as many metal oxides, do not have perfectly fixed ratios. Yet the core of Dalton's theory — atoms keep their identity and are only rearranged in reactions — remains the foundation of chemistry.

Strand 4: how atoms are arranged. A molecule is two or more atoms chemically bonded as a separate unit. Molecules of elements contain one kind of atom (O₂, P₄, S₈); molecules of compounds contain more than one kind (H₂O, CO₂). Atomicity counts the atoms in one molecule. Noble gases exist as single atoms. Allotropes, such as diamond and graphite, are different structures of the same element. Ionic compounds such as NaCl, and giant covalent substances such as silicon dioxide, have no discrete molecules; their formulae give the simplest ratio of particles, called a formula unit.

Strand 5: the mass scale. Since atoms cannot be weighed individually in ordinary experiments, chemists compare them with carbon-12, which is defined as exactly 12. Relative atomic mass, Ar, is an average that includes isotopes, which is why chlorine's value is 35.5 rather than a whole number. Adding the Ar values of every atom in a formula gives the relative formula mass, Mr. Relative masses are ratios, so they have no units.

Formulae

Mr = sum of (number of each atom × its Ar)

percentage by mass of an element = (number of atoms of the element × Ar ÷ Mr) × 100

Step-by-step reasoning

To answer an unseen question that mixes these ideas:

1. Identify the type of substance: atoms, simple molecules, ions or a giant structure. 2. Read the formula carefully, expanding any brackets so that every atom is counted. 3. Look up the Ar values and multiply each by the number of atoms. 4. Add the totals to obtain Mr, then use it for percentages or comparisons. 5. For a reaction, check that each kind of atom is conserved and that the masses of reactants and products balance.

Visual explanation

Imagine the unit as a staircase. The bottom step is the atom. The next step is Dalton's rules for how atoms combine. Above that sit molecules, ions and giant structures, built from atoms. The top step is the relative mass scale, which turns a formula into a number. Each step rests on the one below; remove the atom and the whole staircase collapses.

Real-world analogy

Think of a construction set with a few types of brick. Each brick type is an element, and each finished model is a compound. Taking a model apart and rebuilding it into something new is like a chemical reaction: the bricks are all still there, just arranged differently. Weighing the bricks tells you the total mass, whatever shape you build.

Real-world example

A farmer choosing a nitrogen fertiliser uses every idea in this unit. Ammonium nitrate, NH₄NO₃, has Mr = 14 + (4 × 1) + 14 + (3 × 16) = 80. It contains two nitrogen atoms per formula unit, so its percentage of nitrogen is (28 ÷ 80) × 100 = 35%. Comparing such percentages helps decide which product delivers the most nitrogen for each kilogram spread on a field.

Why?

Why do chemists care so much about relative masses when they cannot see atoms? Because reactions happen between particles in fixed ratios, but in the laboratory we measure mass. Relative masses are the bridge: they let a balanced equation, written in atoms and molecules, be translated into masses that can actually be weighed on a balance.

Common misconception

"Because mass is conserved, the number of molecules must also be conserved." This is false. Atoms are conserved, but molecules can change in number. In 2H₂ + O₂ → 2H₂O, three molecules of reactants become two molecules of product, yet every hydrogen and oxygen atom is still present and the total mass is unchanged.

Worked example

Question: Water forms from hydrogen and oxygen: 2H₂ + O₂ → 2H₂O. Using Ar values H = 1 and O = 16, show that mass is conserved, and state the percentage by mass of hydrogen in water.

Reasoning: Mass of reactants = 2 × (2 × 1) + (2 × 16) = 4 + 32 = 36. Mass of products = 2 × [(2 × 1) + 16] = 2 × 18 = 36. The totals are equal because the same four hydrogen atoms and two oxygen atoms appear on both sides. For water, Mr = 18 and hydrogen contributes 2, so the percentage of hydrogen = (2 ÷ 18) × 100 ≈ 11.1%.

Answer: Both sides total 36 relative mass units, so mass is conserved; water is about 11.1% hydrogen by mass.

Quick check

1. Why is the formula NaCl not described as a molecule? Answer: Sodium chloride is a giant ionic lattice with no separate molecules; NaCl is a formula unit showing a 1:1 ratio of Na⁺ to Cl⁻ ions.

Exam focus

Examiners often combine several ideas in one question: identify the particle type, count atoms from a formula with brackets, calculate Mr and then a percentage. Show every multiplication, remember that elements such as oxygen and nitrogen are diatomic, and never give units for Ar or Mr.

Advanced insight

The ideas in this unit lead directly to the mole. Because Mr compares particle masses, a mass in grams equal to Mr always contains the same number of particles, about 6.02 × 10²³. That link turns a formula such as Mg(NO₃)₂, with Mr = 24 + 2 × (14 + 48) = 148, into a practical way of counting particles by weighing.

Summary

Matter is made of atoms, which Dalton described as combining in whole-number ratios and being rearranged, not destroyed, in reactions. This explains conservation of mass, constant composition and multiple proportions. Atoms form molecules, ions or giant structures. Relative atomic masses, based on carbon-12, add up to relative formula masses, which allow percentages by mass and reaction masses to be calculated.

Practice questions

1. State two parts of Dalton's atomic theory that are still accepted and one that has been corrected. Answer: Still accepted: elements are made of atoms, and atoms are rearranged but not created or destroyed in reactions. Corrected: atoms are not indivisible, since they contain protons, neutrons and electrons. 2. Carbon forms CO and CO₂. What is the ratio of the masses of oxygen that combine with 12 g of carbon in the two compounds, and which law does this illustrate? Answer: 16 g and 32 g, a ratio of 1:2; this illustrates the law of multiple proportions. 3. Calculate Mr for calcium hydroxide, Ca(OH)₂ (Ca = 40, O = 16, H = 1). Answer: 40 + 2 × (16 + 1) = 40 + 34 = 74. 4. Give the atomicity of ozone, O₃, and of argon, Ar, and explain the difference. Answer: Ozone has atomicity 3, because each molecule has three oxygen atoms; argon has atomicity 1, because as a noble gas it exists as single atoms. 5. Calculate the percentage by mass of carbon in carbon dioxide, CO₂ (C = 12, O = 16). Answer: Mr = 12 + 32 = 44; percentage of carbon = (12 ÷ 44) × 100 ≈ 27.3%.