Chemical Bonding and Molecular Structure

Atoms rarely live alone. They join together to form the molecules and compounds that make up everything around us. The forces that hold atoms together are chemical bonds. To understand bonding, we first need the building blocks: what atoms are, what is inside them, and how the periodic table organises them. Understanding why atoms bond, the main types of bond, and how bonding decides the shape of a molecule is essential to all of chemistry.

Long before atoms were seen, careful weighing of reactions revealed two basic laws:

• The Law of Conservation of Mass: in a chemical reaction, mass is neither created nor destroyed. The total mass of the products equals the total mass of the reactants.
• The Law of Constant (Definite) Proportions: a pure chemical compound always contains the same elements combined in the same fixed proportion by mass. Water, for instance, is always hydrogen and oxygen in a mass ratio of 1 : 8.

These laws were explained by John Dalton's atomic theory. It proposed that all matter is made of indivisible atoms. Atoms combine in fixed whole-number ratios. These ideas turned chemistry from guesswork into an exact science.

An atom is the smallest particle of an element that takes part in a chemical reaction. Atoms are extraordinarily small, far too tiny to be seen even with an ordinary microscope. Each element is made of its own kind of atom and is represented by a symbol (H for hydrogen, O for oxygen, Na for sodium from its Latin name natrium).

Atoms usually join together:

• A molecule is a group of two or more atoms held together by chemical bonds. It is the smallest particle of an element or compound that can exist independently. A molecule can be made of one element (such as O₂, two oxygen atoms) or of a compound (such as H₂O, water).
• The number of atoms in a molecule of an element is its atomicity (oxygen is diatomic, ozone triatomic).
• An ion is a charged particle. It is an atom or group of atoms that has lost or gained electrons. A positive ion is a cation. A negative ion is an anion. Compounds like common salt are made of ions.

Atoms are too small to weigh directly, so chemists use a relative scale. On this scale, atom masses are compared against a fixed standard. That standard is one-twelfth the mass of a carbon-12 atom, defined as one atomic mass unit (u).

• Atomic mass is the mass of an atom on this scale (hydrogen ≈ 1 u, oxygen ≈ 16 u).
• Molecular mass is the sum of the atomic masses of all atoms in a molecule (water = 2×1 + 16 = 18 u).

Reactions involve vast numbers of atoms, so chemists use a counting unit called the mole. One mole of any substance contains Avogadro's number of particles, about 6.022 × 10²³. Its mass in grams equals its atomic or molecular mass.

The atom was once thought to be indivisible. It is in fact built from three subatomic particles:

• Protons: positively charged, found in the central nucleus.
• Neutrons: no charge, also in the nucleus.
• Electrons: negatively charged, very light, moving around the nucleus.

The nucleus at the centre is tiny but holds almost all the atom's mass (the protons and neutrons). The electrons occupy the much larger space around it. An atom is neutral overall because it has equal numbers of protons and electrons.

Our picture of the atom developed step by step:

• Thomson imagined electrons embedded in a positive sphere (the "plum pudding").
• Rutherford, from his gold-foil experiment, showed the atom has a tiny, dense, positive nucleus with electrons around it.
• Bohr proposed that electrons move in fixed orbits (shells) of definite energy, explaining the spectra of atoms.

Modern quantum theory refines this further, but Bohr's shells remain a useful picture.

Two numbers identify an atom:

• The atomic number (Z) is the number of protons. It defines the element. Every carbon atom has 6 protons.
• The mass number (A) is the total of protons plus neutrons.

Atoms of the same element with different numbers of neutrons are called isotopes (for example, carbon-12 and carbon-14). They have the same chemistry but different masses.

The chemistry of an element depends on how its electrons are arranged, called its electronic configuration. Electrons fill shells (energy levels) around the nucleus in a definite order, starting from the innermost. Each shell can hold only a limited number of electrons.

The electrons in the outermost shell are the valence electrons. They decide how an atom bonds and reacts. Chemistry is, at its core, the behaviour of these outer electrons.

With over a hundred elements known, chemists needed a way to organise them. The first great periodic table was made by the Russian chemist Dmitri Mendeleev. He arranged the known elements in order of increasing atomic mass, placing those with similar properties in the same column. His table was so insightful that he left gaps for undiscovered elements and correctly predicted their properties.

Later, when the atom's structure was understood, the table was reorganised by atomic number rather than mass, giving the modern periodic table we use today.

The modern periodic law states that the properties of elements are a periodic function of their atomic numbers. In other words, when elements are arranged by atomic number, similar properties recur at regular intervals.

The table is laid out in:

• Periods: the horizontal rows (there are seven), and
• Groups: the vertical columns (eighteen). Elements in the same group share similar chemical behaviour because they have the same number of outer electrons.

Elements are also grouped into metals, non-metals and metalloids, and into blocks (s, p, d, f).

The real power of the table is the trends it reveals. Properties change in regular ways across a period and down a group:

• Atomic size generally decreases across a period (left to right) and increases down a group.
• Ionisation energy (the energy to remove an electron) generally increases across a period and decreases down a group.
• Electronegativity (the pull on bonding electrons) increases across a period and decreases down a group.

These trends let chemists predict reactivity and bonding without memorising every element. That is why the periodic table is so valuable.

Atoms bond in order to become more stable, reaching a lower-energy state. The key to stability is the outer shell of electrons: atoms are most stable with a complete outer shell (like the unreactive noble gases).

So atoms gain, lose or share electrons until their outer shells are full. This drive towards a stable electron arrangement (often eight electrons, the "octet") is the reason chemical bonds form at all.

The two main types of bond differ in how the electrons are handled:

• An ionic bond forms when one atom transfers electrons to another. The atom that loses electrons becomes a positive ion, the one that gains them a negative ion, and the opposite charges attract. This happens between metals and non-metals, as in common salt (sodium chloride).
• A covalent bond forms when atoms share electrons. This happens between non-metals, as in water (H₂O) and methane (CH₄). Sharing can be single, double or triple bonds.

Ionic compounds tend to be solids with high melting points. Covalent compounds are often gases, liquids or soft solids.

Molecules are not flat blobs but have definite three-dimensional shapes, and these shapes affect their properties. A simple, powerful idea explains them: the VSEPR theory (Valence Shell Electron Pair Repulsion).

VSEPR says that the electron pairs around a central atom repel each other and so spread out as far apart as possible. This determines the shape:

• methane (CH₄) is tetrahedral,
• water (H₂O) is bent, and
• carbon dioxide (CO₂) is linear.

The shape of a molecule influences whether it is polar, how it reacts, and even biological function. For example, the bent shape of water makes it a remarkable solvent.