2. Atoms, elements and compounds

A compound is a new pure substance with properties different from its elements. In a mixture, each component keeps its chemical identity and no single formula describes the whole mixture.

In a particle diagram, an element contains only one atom symbol, even if its atoms are paired. A compound has particles containing different atom symbols bonded together in the same pattern. A mixture shows more than one kind of particle.

An atom is electrically neutral because it has equal numbers of positively charged protons and negatively charged electrons.

Because an electron's relative mass is extremely small, nearly all atomic mass comes from the protons and neutrons in the nucleus.

Forming an ion changes only the number of electrons. It does not change the proton number, so the ion remains the same element.

In nuclide notation ᴬ_ZX, the upper number is mass number A and the lower number is proton number Z. Both are whole-number counts for one atom.

Charge = number of protons − number of electrons. Particles with the same electron configuration are isoelectronic, but they can have different proton numbers and charges.

2,8,1 has three occupied shells and one outer electron, so it is Period 3, Group I. The configuration 2,8,7 is Period 3, Group VII. The configuration 2,8,8 is Period 3, Group VIII.

Helium is a noble gas with configuration 2. Its first shell is full with two electrons, so it belongs to Group VIII despite not having eight outer electrons.

To test a pair, compare proton numbers first: if they differ, the atoms are different elements. If proton numbers match, compare neutron or mass numbers; a difference confirms isotopes.

To build a symbol from counts, use the proton number to identify X, add protons and neutrons for A, then compare protons with electrons to find the charge. More protons than electrons gives positive charge; more electrons gives negative charge.

For ³⁷₁₇Cl⁻: protons = 17, neutrons = 37 − 17 = 20 and electrons = 18. The charge changes electron count, not A or Z.

A complete explanation links the same number of electrons to the same electronic configuration, especially the same outer-shell arrangement that controls bonding and reactions.

Aᵣ = Σ(isotope relative mass × abundance) ÷ Σ(abundance). For percentage abundances, the denominator is 100; for a ratio, it is the sum of the ratio parts.

The answer must lie between the lightest and heaviest isotope masses and be closer to the mass of the more abundant isotope. Round only as requested after completing the weighted sum.

Ion charge = number of protons − number of electrons. Losing two electrons gives a 2+ ion; gaining two electrons gives a 2− ion.

A magnesium atom loses two outer electrons to form Mg²⁺. An oxygen atom gains two electrons to form O²⁻. In both cases the nucleus and element identity are unchanged.

Electrostatic attraction acts between positive and negative charge. In an ionic compound, this attraction holds cations and anions together.

In a dot-and-cross diagram, use one symbol for each atom's original outer electrons and the other symbol for the transferred electron. Put each ion in brackets, show its full outer shell and write the charge outside the bracket.

For sodium chloride, Na transfers one electron to Cl, forming Na⁺ and Cl⁻ in a 1:1 ratio. The diagram must show the chloride outer shell as seven original electrons plus the transferred sodium electron.

Molten means melted into a liquid. Aqueous means dissolved in water. State matters when describing conductivity.

Oppositely charged ions alternate through the lattice so that each ion is surrounded by ions of opposite charge. The pattern extends in all directions rather than ending as separate molecules.

Mg loses two electrons: Mg²⁺ pairs with O²⁻ in MgO, but with two Cl⁻ ions in MgCl₂. Al³⁺ needs three F⁻ ions in AlF₃. The ion ratio is set by charge balance.

Show transferred electrons with a different dot/cross symbol from the receiving atom's original electrons. Include all required ions in the simplest charge-balanced ratio.

Both cations and anions move when the ionic substance conducts. The mobile charge carriers are ions, not electrons released from the ionic bonds.

The shared pair is attracted to both nuclei and counts in the outer shell of both bonded atoms. Hydrogen reaches a full first shell with two electrons; many other simple non-metal atoms reach eight outer electrons.

Show outer-shell electrons only when instructed. Every H should count two electrons around it; C, N, O and Cl should count eight from bonding pairs plus lone pairs.

Start from the displayed connectivity, place the required shared pair(s) between bonded atoms, then add lone pairs until each H has two outer electrons and each C, N or O has eight.

Keep dots and crosses consistent with electron origin. A double bond must show four shared electrons between the same two atoms; a triple bond must show six.

The general name for the weak attractions between separate molecules is intermolecular forces. Specific types of intermolecular force are not required by this syllabus objective.

Covalent bonds within molecules are strong; intermolecular forces between molecules are much weaker. This difference is why a molecular substance can melt without its molecules decomposing.

Neither structure consists of separate carbon molecules. The covalent network extends through a giant number of atoms; a drawn model shows only a small repeating part.

Strong covalent bonding throughout both structures also gives high thermal stability, so the giant framework is not easily separated by heating.

The structure contains no separate SiO₂ molecules. The formula gives the simplest atom ratio in the giant network, and the covalent bonds continue throughout the solid.

All outer electrons are held in covalent bonds, so there are no mobile ions or delocalised electrons in the structure.

The networks contain different atoms—only carbon in diamond, silicon and oxygen in SiO₂—but both connect atoms by strong covalent bonds in a giant framework.

The electrons remain inside the metal, so the structure is electrically neutral overall. The metal is not a lattice of neutral atoms or alternating positive and negative ions.

Layer movement does not mean metallic bonding is weak. The mobile electron sea continues to attract the positive ions as their positions change.