Periodic trends, bonding

Electronegativity

The measure of how strongly an atom attracts electrons towards itself.

Electronegativity change across a period

It increases across a period due to increasing nuclear charge, pulling valence electrons closer.

Electronegativity change down a group

It decreases down a group because more electron shells mean weaker attraction between nucleus and electrons.

First ionisation energy

The energy required to remove one electron from a neutral atom.

First ionisation energy change across a period

It increases across a period because electrons are more strongly attracted to the nucleus.

First ionisation energy change down a group

It decreases down a group because electrons are further from the nucleus and easier to remove.

Atomic radius

The distance from the nucleus to the outermost electron shell.

Atomic radius change across a period

It decreases across a period due to increased nuclear attraction pulling electrons closer.

Atomic radius change down a group

It increases down a group because additional shells make the atom larger.

Melting point

The temperature at which a substance changes from solid to liquid.

Melting point of metals change across a period

It generally increases due to stronger metallic bonding from more delocalised electrons.

Melting point of metals change down a group

It generally decreases as atomic radius increases, weakening metallic bonds.

Exceptions to melting point trend

Non-metals (groups 13-18) and elements 104-118 do not follow this trend.

Bohr's model of the atom

Electrons orbit the nucleus in fixed circular paths with specific energy levels.

Schrodinger's model of the atom

Electrons occupy orbitals in a 3D space with wave-like behaviour; atoms have shells and subshells.

Metallic bonding

Positive metal ions are surrounded by a sea of delocalised electrons held together by electrostatic attraction.

Metal lattice

A structure of cations and delocalised electrons in a repeating pattern.

Why are metals lustrous?

Free electrons reflect light, making the metal shiny.

Why are metals malleable and ductile?

Layers of cations can slide over each other without breaking the metallic bond.

Why do metals conduct electricity?

Delocalised electrons move freely through the metal towards positive terminals.

Ionic bond

The electrostatic attraction between oppositely charged ions formed when electrons transfer from metal to non-metal.

Ionic lattice

A 3D structure of alternating cations and anions held by strong ionic bonds.

Why do ionic compounds have high melting points?

Strong electrostatic forces require lots of energy to break.

Why are ionic compounds brittle?

Applying force can shift ions and cause repulsion between like charges, breaking the lattice.

Do ionic compounds conduct electricity?

Yes, but only when molten or dissolved in water due to free-moving ions.

Covalent bonding

Bonding between non-metals involving shared electrons.

Single, double, and triple covalent bonds

Single: one shared pair of electrons; Double: two shared pairs; Triple: three shared pairs.

Intramolecular forces

Strong covalent bonds within a molecule.

Intermolecular forces

Weaker attractions between molecules that affect physical properties.

Types of intermolecular forces

Dispersion forces, Dipole-dipole interactions, Hydrogen bonding.

Molecule shape determination

By the number of electron groups and lone pairs on the central atom.

Shape of a molecule with 4 electron groups and no lone pairs

Tetrahedral.

Shape of a molecule with 3 bonding pairs and 1 lone pair

Trigonal pyramidal.

Shape of a molecule with 2 bonding pairs and 2 lone pairs

Bent (e.g. water).

Bond polarity determination

The difference in electronegativity between two atoms.

Polar and nonpolar bonds

Polar: unequal sharing of electrons; Nonpolar: equal sharing of electrons.

Can a molecule have polar bonds and still be nonpolar?

Yes, if the molecule is symmetrical.

Solubility determination

The ability of a solute to dissolve in a solvent to form a solution, depending on their nature, temperature, and pressure.

Solute

The substance that is dissolved.

Solvent

The substance that dissolves the solute.

Solution

A homogeneous mixture of solute and solvent.

What are subshells in an atom?

Regions within shells where electrons are likely to be found (s, p, d, f).

How many electrons can each shell hold?

Given by 2n², where n is the shell number.

What is the maximum number of electrons in each subshell?

s: 2, p: 6, d: 10, f: 14

In what order do subshells fill with electrons?

Subshells fill in the order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.

How do valence electrons affect an element's reactivity?

Elements with a full or nearly full/empty valence shell are more reactive due to their desire to gain, lose, or share electrons.

Why do polar substances dissolve in polar solvents?

Because polar molecules interact with polar solvents through dipole interactions, allowing them to mix.

Why do nonpolar substances dissolve in nonpolar solvents?

Nonpolar molecules dissolve in nonpolar solvents due to similar dispersion forces that allow them to mix.

What is the shape of a molecule with 2 electron groups and no lone pairs?

Linear (e.g. carbon dioxide).

Why do molecular compounds usually have low melting and boiling points?

Because they have weak intermolecular forces that require little energy to break.

How does hydrogen bonding affect boiling point?

Hydrogen bonding increases the boiling point because it's a strong type of intermolecular attraction.

What is the difference between intramolecular and intermolecular forces?

Intramolecular forces hold atoms together within a molecule, while intermolecular forces occur between molecules and are generally weaker.

How does molecular size affect dispersion forces?

Larger molecules have more electrons, leading to stronger dispersion forces.

How does the shape of a molecule affect its polarity?

If the molecule is asymmetrical and has polar bonds, it will be polar overall.