Atoms and Molecules — Study Notes

1. Introduction

Everything in the universe is made of matter, and matter is made up of tiny particles. In this chapter we study the building blocks of matter - atoms and molecules. The idea that matter is composed of extremely small particles was proposed long ago by philosophers, but it was John Dalton who, in 1808, put forward a scientific atomic theory that could explain the laws of chemical combination.

The study of atoms and molecules allows us to write chemical formulae, calculate molecular and formula unit masses, and use the mole concept to count particles by weighing them. Understanding atoms and molecules is the first step toward mastering chemistry, because every chemical reaction is essentially a rearrangement of atoms.

In this chapter we cover the laws of chemical combination, Dalton's atomic theory, atomic and molecular masses, writing chemical formulae using valency, and the mole concept with Avogadro's number.

2. Laws of Chemical Combination

Antoine Lavoisier and Joseph L. Proust established two fundamental laws which helped Dalton frame his atomic theory.

Law of Conservation of Mass: This law states that mass can neither be created nor destroyed in a chemical reaction. In any chemical reaction, the total mass of the reactants equals the total mass of the products. For example, when 100 g of calcium carbonate is heated, it gives 56 g of calcium oxide and 44 g of carbon dioxide. The total mass of products, 56 + 44 = 100 g, equals the mass of the reactant.

Law of Constant Proportions: This law, given by Joseph Proust, states that in a chemical compound, the elements are always present in a definite proportion by mass. For example, water always contains hydrogen and oxygen in the ratio 1:8 by mass, whether it is obtained from rain, a river or the sea. Similarly, ammonia (NH3) always contains nitrogen and hydrogen in the ratio 14:3 by mass.

3. Dalton's Atomic Theory

John Dalton published his atomic theory in 1808 to explain the laws of chemical combination. The main postulates of the theory are:

1. All matter is made up of extremely small particles called atoms.
2. Atoms are indivisible particles that cannot be created or destroyed in a chemical reaction. This explains the law of conservation of mass.
3. All atoms of a given element are identical in mass and chemical properties.
4. Atoms of different elements have different masses and chemical properties.
5. Atoms combine in the ratio of small whole numbers to form compounds. This explains the law of constant proportions.
6. The relative number and kinds of atoms in a given compound are constant.

Dalton's atomic theory successfully explained the laws of chemical combination, but later discoveries showed that atoms are divisible and can be further split into protons, neutrons and electrons.

4. Atoms and Molecules

Atom: An atom is the smallest particle of an element that can take part in a chemical reaction. Atoms are extremely small; their radius is measured in nanometers. The radius of a hydrogen atom is about 10^-10 m (0.1 nm). Atoms may or may not exist independently.

Atomic mass unit: The atomic mass of an atom is its mass relative to the carbon-12 standard. One atomic mass unit (u) is defined as one twelfth (1/12) of the mass of one carbon-12 atom. Using this unit, the atomic mass of hydrogen is about 1 u, carbon is 12 u, oxygen is 16 u, nitrogen is 14 u, and sodium is 23 u.

Molecule: A molecule is the smallest particle of an element or a compound which can exist independently and shows all the properties of that substance. Molecules are formed by the combination of two or more atoms.

Molecules of elements contain atoms of the same kind, such as oxygen (O2), hydrogen (H2), nitrogen (N2) and ozone (O3). Molecules of compounds contain atoms of different elements combined in a fixed ratio, such as water (H2O) and ammonia (NH3).

Atomicity: The number of atoms present in a molecule of an element is called its atomicity. For example, helium (He) is monoatomic, oxygen (O2) is diatomic, and sulphur (S8) is polyatomic.

5. Ions

Compounds made of metals and non-metals contain charged particles called ions. A positively charged ion is called a cation, and a negatively charged ion is called an anion.

• Cations: Na+, Ca2+, Mg2+, Al3+
• Anions: Cl-, O2-, S2-

Some ions are made up of groups of atoms which carry a net charge; these are called polyatomic ions. Examples are ammonium (NH4+), sulphate (SO4 2-), nitrate (NO3-), carbonate (CO3 2-) and hydroxide (OH-).

The valency of an ion is equal to the charge on the ion. For example, the valency of Na+ is 1, of Mg2+ is 2, and of Al3+ is 3.

6. Writing Chemical Formulae

The chemical formula of a compound gives the symbolic representation of its composition. To write a chemical formula, we need the symbols of the elements and their valencies. The rules are:

1. The valencies or charges on the ions must balance.
2. When a compound is made of a metal and a non-metal, the metal (or its symbol) is written first, as in NaCl and CuO.
3. In compounds with polyatomic ions, the ion is placed in brackets followed by the required number, as in Mg(OH)2. If the number is one, brackets are not needed, as in NaOH.

We use the criss-cross method to write formulae. For example, calcium (Ca, valency 2) and hydroxide (OH, valency 1): crossing the valencies gives Ca(OH)2. For aluminium (Al, valency 3) and sulphate (SO4, valency 2), the formula is Al2(SO4)3.

7. Molecular Mass and Formula Unit Mass

The molecular mass of a substance is the sum of the atomic masses of all the atoms in a molecule of the substance. It is expressed in atomic mass units (u). For example, the molecular mass of water (H2O) is 2 x 1 (for H) + 16 (for O) = 18 u. The molecular mass of carbon dioxide (CO2) is 12 + 2 x 16 = 44 u.

The formula unit mass is used for substances whose constituent particles are ions, such as NaCl. It is calculated in the same way as molecular mass but uses the formula unit of the substance. The formula unit mass of NaCl is 23 + 35.5 = 58.5 u.

8. The Mole Concept

A mole is the amount of a substance that contains as many particles (atoms, molecules or ions) as there are atoms in exactly 12 g of carbon-12. This fixed number of particles is called Avogadro's number, which has the value 6.022 x 10^23.

The molar mass of a substance is the mass of one mole of the substance in grams. It is numerically equal to the atomic or molecular mass of the substance but has the unit g/mol. For example, the molar mass of water is 18 g/mol, and the molar mass of carbon is 12 g/mol.

The number of moles in a given mass of a substance is calculated by the formula:

Number of moles (n) = mass of substance (m) / molar mass (M)

The number of particles can be found using Avogadro's number:

Number of particles = Number of moles x 6.022 x 10^23

For example, 18 g of water contains 1 mole of water molecules, which is 6.022 x 10^23 molecules of water. Each water molecule contains 2 hydrogen atoms and 1 oxygen atom.

Quick Revision Tables

Mind Map

Important Diagrams (SVG)

Diagram 1: Laws of Chemical Combination

Diagram 2: Mole Concept and Avogadro's Number

Common Mistakes

1. Believing that atoms are divisible; Dalton said atoms are indivisible, though later experiments showed they contain subatomic particles.
2. Forgetting to balance the valencies while writing a chemical formula, such as writing CaOH instead of Ca(OH)2.
3. Using the wrong ratio for water; hydrogen and oxygen combine in the ratio 1:8 by mass, not 1:2.
4. Confusing molecular mass (in u) with molar mass (in g/mol); they are numerically equal but have different units.
5. Writing the metal after the non-metal in a formula; in a compound of a metal and non-metal, the metal is written first.
6. Forgetting brackets when a polyatomic ion appears more than once, as in Al2(SO4)3.
7. Confusing atomicity with atomic mass; atomicity is the number of atoms in a molecule, while atomic mass is the mass of an atom.

Exam Tips

1. State the law of conservation of mass and law of constant proportions with one example each.
2. List the six postulates of Dalton's atomic theory and mention which one was later found to be incorrect (atoms are divisible).
3. Practise writing formulae using the criss-cross method for common compounds like Ca(OH)2, Al2(SO4)3 and MgCl2.
4. Learn to calculate molecular mass and formula unit mass step by step, as in H2O = 18 u and NaCl = 58.5 u.
5. Use the mole formulae n = m/M and N = n x NA to solve numerical problems, and remember that NA = 6.022 x 10^23.
6. Memorise the atomicities: He monoatomic, O2 diatomic, O3 triatomic, P4 tetra-atomic, S8 polyatomic.
7. Remember examples of polyatomic ions and their charges: NH4+, SO4 2-, NO3-, CO3 2-, OH-.

Conclusion

In this chapter we learned that matter is made of atoms and molecules, and that the laws of chemical combination - the law of conservation of mass and the law of constant proportions - govern how elements combine. Dalton's atomic theory provided a scientific explanation for these laws. We studied the concepts of atomic mass unit, molecular mass and formula unit mass, and learned to write chemical formulae using valency and the criss-cross method. Finally, the mole concept introduced us to Avogadro's number, enabling us to count the number of particles in a given mass of a substance. These fundamental ideas are the building blocks of chemistry and are used in balancing chemical equations, stoichiometry and quantitative analysis.