Structure and Bonding in Solids – Lecture 1

Flashcards covering key points from ENMP102B Lecture 1 on atomic structure, periodic trends, and primary/secondary bonding in solids.

Which four fundamental factors determine the properties of a solid material?

1) The types of atoms present, 2) the geometrical arrangement of atoms (often crystalline), 3) the nature of the inter-atomic bonding, and 4) imperfections/defects in the crystal structure.

What particles make up an atomic nucleus?

Protons and neutrons.

What is the magnitude of the charge on a single electron or proton?

±1.60 × 10⁻¹⁹ C (negative for electrons, positive for protons).

Define the atomic number (Z) of an element.

The number of protons in the nucleus of an atom.

How is the atomic mass (A) of an element approximated?

A ≈ Z + N, where N is the number of neutrons.

How many particles are in one mole of a substance?

6.023 × 10²³ atoms or molecules (Avogadro’s number).

According to the Pauli Exclusion Principle, how many electrons can occupy a single orbital (subshell state)?

Two, provided they have opposite spins.

What happens when an electron undergoes a ‘quantum jump’?

It absorbs or emits a discrete amount of energy to move between allowed energy levels.

What are valence electrons?

Electrons occupying the outermost electron shell; they largely control bonding and chemical behaviour.

Why do elements within the same column of the periodic table exhibit similar properties?

They have similar valence-electron configurations.

Why are Group 0 (18) elements chemically inert?

They possess completely filled valence shells, giving them very stable electron configurations.

Why are Group IA elements highly reactive?

They have only one valence electron, which is easily lost to attain a stable configuration.

Define electronegativity in the context of bonding.

A measure of an atom’s tendency to attract (accept) electrons when forming bonds.

What is meant by an ‘electropositive’ atom?

An atom (typically metallic) that readily donates its valence electrons to form positive ions.

At approximately what interatomic distance (r₀) does bonding equilibrium typically occur?

Around 0.3 nm for many atoms.

Name the three primary (chemical) bond types.

Ionic, covalent, and metallic bonds.

Name the three main categories of secondary (physical) bonds.

Fluctuating induced dipole, polar-induced dipole, and permanent dipole (including hydrogen bonds).

What drives the formation of primary bonds at the electronic level?

The attainment of stable (usually filled) outer electron shells for the participating atoms.

Describe the essence of ionic bonding.

Valence electrons transfer from electropositive (metal) atoms to electronegative (non-metal) atoms, producing oppositely charged ions that attract via Coulombic forces.

Give a common solid that is predominantly ionically bonded.

Sodium chloride (NaCl).

Describe the essence of covalent bonding.

Atoms share pairs of valence electrons to achieve stable electron configurations; the bonds are directional.

Provide two examples of materials with strong covalent bonding.

Diamond (carbon) and silicon (Si).

What is metallic bonding?

Valence electrons are delocalised and form an ‘electron cloud’ that glues a lattice of positive metal ions together.

Why are metals typically good electrical and thermal conductors?

Their delocalised (free) valence electrons can move easily, transporting charge and heat.

Why are ionic solids often brittle?

When like-charged ions are forced adjacent under stress, strong Coulombic repulsion causes the crystal to cleave.

Why are metals generally ductile?

Metallic bonding is non-directional, allowing planes of atoms to slide without breaking the overall bonding network.

What is the relative bond energy of secondary (Van der Waals) interactions compared to primary bonds?

Secondary bond energies are one to two orders of magnitude smaller than primary bond energies.

How do fluctuating induced dipole bonds arise?

Momentary distortions in electron density create temporary dipoles that induce complementary dipoles in neighbouring atoms/molecules.

Which physical phenomenon in inert gases relies on fluctuating induced dipole bonding?

Liquefaction (formation of the liquid phase).

What is a polar molecule–induced dipole interaction?

A permanent dipole in one molecule induces a dipole in a neighbouring non-polar molecule, resulting in attraction stronger than that from purely fluctuating dipoles.

What is the strongest type of secondary bond and when does it occur?

The hydrogen bond, occurring when hydrogen is covalently bonded to a highly electronegative atom (e.g., O, N, F), leaving the hydrogen side as a nearly bare proton.

Explain why real materials seldom possess purely ionic or purely covalent bonds.

Because bonding character depends on electronegativity difference; most bonds have both ionic and covalent components to varying degrees.

What property trend can be explained by understanding bonding types?

Electrical conductivity: metals (metallic bonds) conduct well, whereas ionic and covalently bonded materials tend to be insulators.

List the three questions posed in the lecture recap.

1) What main factors affect solid-state properties? 2) What are the main primary and secondary bond types? 3) Provide a solid example for each primary bond type and state what drives these bonds.