1. Overview
Period 3 runs from sodium to argon. Across it, the elements change from metals with giant metallic lattices, through silicon's giant molecular structure, to non-metals made of small molecules or single atoms. This section describes how atomic radius, ionic radius, melting point and electrical conductivity change across the period, and explains melting point and conductivity from the structure and bonding of each element. Similar patterns repeat in other periods, most clearly in Period 2, and that repetition is what periodicity means.
Key Definitions
- Periodicity: the repeating pattern in the physical and chemical properties of the elements across each period of the Periodic Table, seen when a property is plotted against proton number.
- Atomic radius: half the distance between the nuclei of two neighbouring atoms of the same element that are bonded together (covalently in a non-metal, by metallic bonding in a metal).
- Ionic radius: the radius of an ion in an ionic lattice, worked out from the distance between the nuclei of neighbouring ions.
- Melting point: the temperature at which a solid turns into a liquid at a stated pressure; it is high when strong bonds or forces must be overcome before the particles can leave their fixed positions.
- Electrical conductivity: a measure of how well a substance carries an electric current; a substance conducts only if it contains charged particles, electrons or ions, that are free to move.
- Metallic bonding: the electrostatic attraction between positive metal ions and delocalised electrons.
- Giant molecular structure: a structure, also called giant covalent, in which a very large number of atoms are joined by strong covalent bonds into a continuous network with no separate molecules; in silicon the network is three-dimensional.
- Simple molecular structure: a substance made of separate small molecules, such as P₄, S₈ and Cl₂, with strong covalent bonds inside each molecule but only weak intermolecular forces between the molecules.
- Instantaneous dipole-induced dipole (id-id) forces: weak attractions between molecules, or between noble gas atoms, that arise when an uneven spread of electrons at one instant gives a particle a temporary dipole, which induces a dipole in a neighbouring particle; they are one type of van der Waals' force and are generally stronger when the particles have more electrons.
- Semiconductor: a solid, such as silicon, that conducts electricity far less well than a metal at room temperature but better when heated, because heating frees some electrons from their covalent bonds.
Content
Period 3 at a glance
The melting point and conductivity trends both come from the structure of each element, so learn these two tables first. In them, id-id forces means instantaneous dipole-induced dipole forces, the type of weak van der Waals' force that acts between these non-polar molecules and atoms (topic 3.6).
| Element | Structure | Particles |
|---|---|---|
| Na | giant metallic | Na⁺ ions and delocalised electrons |
| Mg | giant metallic | Mg²⁺ ions and delocalised electrons |
| Al | giant metallic | Al³⁺ ions and delocalised electrons |
| Si | giant molecular | Si atoms |
| P | simple molecular | P₄ molecules |
| S | simple molecular | S₈ molecules |
| Cl | simple molecular | Cl₂ molecules |
| Ar | single atoms | Ar atoms |
| Elements | Held together by | Conducts? |
|---|---|---|
| Na, Mg, Al | metallic bonding | well |
| Si | covalent bonds throughout | very slightly (semiconductor) |
| P, S, Cl | id-id forces between molecules | no |
| Ar | id-id forces between atoms | no |
- At room temperature, sodium to sulfur are solids; chlorine and argon are gases.
- Phosphorus and sulfur each exist in more than one form. This page uses white phosphorus, made of P₄ molecules, and the usual forms of sulfur, which are made of S₈ rings.
- Argon atoms do not bond to each other, so argon exists as single atoms.
Atomic radius across Period 3
Atomic radius decreases from sodium to chlorine.
Topic 1.1 explains why. In short: the nuclear charge rises from +11 to +17, every outer electron is in the third shell and the shielding by the inner shells stays similar, so the outer electrons are pulled closer to the nucleus.
Argon is left out: its atoms do not bond to each other, so its radius has to be measured between atoms that are not bonded. That gives a larger value, so it cannot be compared fairly with the other seven.
At the start of the next period the radius jumps up, because potassium's outer electron is in a fourth shell. Then the fall across the period begins again.
Ionic radius across Period 3
Ionic radius does not follow the atomic radius pattern. It falls from Na⁺ to Al³⁺, jumps up to P³⁻, then falls again to Cl⁻.
- Within each set, the ions are isoelectronic (they have the same number of electrons). Na⁺, Mg²⁺ and Al³⁺ each have 10 electrons in two shells. P³⁻, S²⁻ and Cl⁻ each have 18 electrons in three shells. More protons pull the same electrons closer, so the radius falls across each set.
- The jump from Al³⁺ to P³⁻ comes from the extra occupied shell. Each of the three negative ions is larger than each of the three positive ions.
- Compared with the atoms (the dashed line): each positive ion is smaller than its atom, because it has lost its whole third shell. Each negative ion is larger than its atom, because the added electrons increase the repulsion between the electrons while the nuclear charge stays the same.
- Silicon is usually left out, because it bonds covalently rather than forming simple ions. If a data table does give a value for silicon, check the charge: Si⁴⁺ (10 electrons) fits the positive set as its smallest member, and Si⁴⁻ (18 electrons) fits the negative set as its largest member.
Topic 1.1 gives the full explanation of ion sizes.
Melting point: the shape of the trend
Reading the graph from left to right:
- The melting point rises from sodium to magnesium. Aluminium is only slightly higher than magnesium.
- Silicon has the highest melting point of the Period 3 elements.
- There is a very large fall from silicon to phosphorus.
- Sulfur melts at a higher temperature than phosphorus.
- The melting point then falls to chlorine and is lowest for argon. If a graph stops at chlorine, chlorine is the lowest point.
Explaining the melting points
A solid melts when its particles gain enough energy to leave their fixed positions. So for each element, ask: what must be overcome, and how strong is it?
Sodium, magnesium and aluminium: giant metallic lattices. Melting loosens the metallic bonding enough for the ions to move past one another. From sodium to aluminium:
- the charge on the ion rises: Na⁺, Mg²⁺, Al³⁺;
- the number of delocalised electrons per atom rises: 1, 2, 3;
- the ions get smaller.
All three make the attraction between the ions and the delocalised electrons stronger, so more energy is needed and the melting point rises. Topic 3.3 explains this model in full, and why the boiling points of magnesium and aluminium show the trend more clearly than their melting points.
Silicon: a giant molecular structure. Each silicon atom is covalently bonded to four others in a tetrahedral arrangement, and the network continues through the whole crystal, as in diamond. There are no separate molecules, so there are no intermolecular forces to overcome. To melt silicon, many strong Si-Si covalent bonds must be broken. That needs a very large amount of energy, so silicon has the highest melting point in the period.
Phosphorus, sulfur, chlorine and argon: small molecules and single atoms. The covalent bonds inside P₄, S₈ and Cl₂ are strong, but melting does not break them: the molecules stay whole. Only the weak id-id forces between the molecules (or between argon atoms) are overcome, and that takes little energy. This is why the melting point falls so sharply from silicon to phosphorus.
The order within this group depends on the strength of the id-id forces, which increases with the number of electrons in each particle:
| Particle | Electrons in the particle | id-id forces and melting point |
|---|---|---|
| S₈ | 8 × 16 = 128 | strongest and highest of the four |
| P₄ | 4 × 15 = 60 | second |
| Cl₂ | 2 × 17 = 34 | third |
| Ar | 18 | weakest; argon has the lowest melting point in Period 3 |
More electrons give larger temporary dipoles, so the id-id forces are stronger and more energy is needed to separate the particles. (The large S₈ ring also has more surface in contact with its neighbours.) Always compare whole molecules: a sulfur atom has only one more electron than a phosphorus atom, but an S₈ molecule has more than twice as many electrons as a P₄ molecule.
Worked example 1: why bond energies do not explain these melting points
Question: The Data Booklet gives these bond energies: P-P 200, S-S 265, Cl-Cl 242 and Si-Si 225 kJ mol⁻¹. A student writes: "Sulfur melts at a higher temperature than phosphorus because the S-S bond is stronger than the P-P bond." (a) Use the bond energies to show that this reasoning fails. (b) Give the correct explanation.
Step 1: test the idea on another pair. The Cl-Cl bond (242 kJ mol⁻¹) is stronger than the P-P bond (200 kJ mol⁻¹). If bond strength decided the melting point, chlorine would melt above phosphorus. In fact chlorine is a gas at room temperature and phosphorus is a solid, so the idea predicts the wrong order.
Step 2: ask what melting actually breaks. Melting P₄, S₈ or Cl₂ breaks no covalent bonds, so their bond energies do not matter here. Only the id-id forces between the molecules are overcome.
Step 3: compare the id-id forces. Count the electrons in each molecule: P₄ has 4 × 15 = 60 and S₈ has 8 × 16 = 128.
Answer: (a) Cl-Cl (242 kJ mol⁻¹) is stronger than P-P (200 kJ mol⁻¹), yet chlorine has the lower melting point, so bond strength does not decide the order. (b) Melting sulfur or phosphorus breaks no covalent bonds; it overcomes the id-id forces between molecules. An S₈ molecule has 128 electrons and a P₄ molecule has 60, so the id-id forces between S₈ molecules are stronger, more energy is needed to overcome them, and sulfur has the higher melting point.
Silicon shows the other side. Its Si-Si bond (225 kJ mol⁻¹) is weaker than the S-S and Cl-Cl bonds, yet silicon melts at a far higher temperature than sulfur or chlorine. Melting silicon does break covalent bonds, and a great many of them, because every atom is bonded to four others throughout the giant structure.
Electrical conductivity: the shape of the trend
- Conductivity increases from sodium to magnesium to aluminium. The rise from sodium to magnesium is small; aluminium is clearly the best conductor of the Period 3 elements.
- Silicon conducts only very slightly: it is a semiconductor.
- Phosphorus, sulfur, chlorine and argon do not conduct.
Explaining the conductivity
To conduct, a substance needs charged particles that are free to move. In these elements no ions are free to move, so conduction depends on electrons that can.
Sodium, magnesium and aluminium. The delocalised electrons belong to no single ion. When a potential difference is applied, they move through the lattice towards the positive terminal and carry the current. Sodium gives one delocalised electron per atom, magnesium two and aluminium three, so there are more mobile charge carriers from left to right and the conductivity increases.
The measured rise is smaller than these numbers suggest: magnesium has twice as many delocalised electrons per atom as sodium, but conducts only a little better. How freely the electrons move also matters, and that is beyond this course. In an answer, explain the trend by the number of delocalised electrons per atom.
Silicon. Each silicon atom uses all four of its outer electrons in covalent bonds with its four neighbours. These electrons are localised: each bonding pair is held between two particular atoms. With almost no electrons free to move, silicon conducts only very slightly. When it is heated, a few electrons gain enough energy to break free from their bonds, so its conductivity increases with temperature. That is the behaviour of a semiconductor.
Phosphorus to argon. In P₄, S₈ and Cl₂ every outer electron is in a covalent bond or a lone pair belonging to one atom. Argon's electrons are held in the full shells of single atoms. There are no delocalised electrons and no ions, so these elements do not conduct, as solids or as liquids.
Worked example 2: comparing aluminium and silicon
Question: Aluminium and silicon are next to each other in Period 3. Silicon has the higher melting point, but aluminium is by far the better electrical conductor. Explain both differences in terms of structure and bonding.
Step 1: name each structure and its particles. Aluminium is a giant metallic lattice of Al³⁺ ions and delocalised electrons, three per atom. Silicon is a giant molecular structure: each Si atom is covalently bonded to four others, throughout the crystal.
Step 2: melting point: what must be overcome? Melting aluminium only loosens the metallic bonding, so the ions can move past one another. Melting silicon means breaking many strong Si-Si covalent bonds, and that needs more energy than loosening the metallic bonding in aluminium.
Step 3: conductivity: what can move? In aluminium, the delocalised electrons can move through the lattice. In silicon, the outer electrons are localised in covalent bonds, each pair held between two particular atoms.
Answer:
- Melting point. Silicon has a giant molecular structure, so many strong covalent bonds must be broken to melt it. Aluminium has metallic bonding between Al³⁺ ions and delocalised electrons, and melting only loosens it. More energy is needed to melt silicon, so silicon has the higher melting point.
- Conductivity. Aluminium has three delocalised electrons per atom, which are free to move through the lattice and carry charge. In silicon the outer electrons are held in covalent bonds between pairs of atoms, so almost none can move. So aluminium conducts far better than silicon.
Each point names the structure and particles of both elements, says what is overcome or what can move, and ends with a comparison. That is the pattern for any "compare" or "explain the difference" question in this section.
Periodicity: the patterns repeat
Plot these properties against proton number across Periods 2 and 3, and broadly the same shape appears in each period.
- Atomic radius falls from lithium to fluorine, jumps up at sodium (a new shell), then falls again to chlorine.
- Melting point: both periods have the same sequence of structures, giant metallic, then giant molecular, then simple molecular, then single atoms. In Period 2, lithium and beryllium are metallic, boron and carbon are giant molecular, nitrogen, oxygen and fluorine are simple molecular, and neon exists as single atoms. So the melting points rise to very high values in the middle of each period and drop sharply at Group 15. (In Period 2 the giant molecular structures start one group earlier, at boron.)
- Electrical conductivity: in both periods the metals on the left conduct and the elements from Group 15 onwards do not. Group 14 differs: silicon is a semiconductor, but carbon as graphite conducts along its layers, where delocalised electrons can move (diamond does not conduct).
- Ionic radius: in each period the positive ions near the start are much smaller than the negative ions near the end.
Key Equations
This section is mostly descriptive. These counting rules are the ones you need.
Electrons in a molecule of an element (to compare id-id forces):
electrons in the molecule = number of atoms in the molecule × proton number
For example, S₈ has 8 × 16 = 128 electrons.
Delocalised electrons per atom for sodium, magnesium and aluminium: equal to the number of outer-shell electrons, which is also the charge on the ion (Na⁺ 1, Mg²⁺ 2, Al³⁺ 3).
Common Mistakes to Avoid
- Mistake: sketching the melting point graph with phosphorus as the lowest point, or with the peak at aluminium. Fix: silicon is the peak. Phosphorus is low, but sulfur is above it. From sodium to chlorine the lowest point is chlorine; from sodium to argon it is argon.
- Mistake: saying intermolecular forces are overcome when silicon melts. Fix: silicon has a giant molecular structure with no separate molecules. Write "many strong covalent bonds must be broken", not "strong intermolecular forces".
- Mistake: saying covalent bonds break when phosphorus, sulfur or chlorine melts. Fix: the molecules stay whole and only id-id forces are overcome. Bond energies predict the wrong order: Cl-Cl (242 kJ mol⁻¹) is stronger than P-P (200 kJ mol⁻¹), yet chlorine melts at a lower temperature (Worked example 1).
- Mistake: comparing electrons per atom to explain why sulfur melts above phosphorus. Fix: compare whole molecules. A sulfur atom has only one more electron than a phosphorus atom, but S₈ has 128 electrons and P₄ only 60.
- Mistake: writing "metals conduct because they have delocalised electrons" and stopping there. Fix: say that the electrons move: "the delocalised electrons are free to move through the lattice and carry charge".
- Mistake: giving silicon delocalised electrons, for example "four delocalised electrons per atom". Fix: silicon's outer electrons are localised in covalent bonds between pairs of atoms, which is why it conducts only very slightly.
- Mistake: drawing ionic radius as one smooth fall across the period, or making Al³⁺ larger than Mg²⁺. Fix: draw two falling sets with a jump between them. Al³⁺ is the smallest of Na⁺, Mg²⁺ and Al³⁺, because it has 13 protons for the same 10 electrons, and P³⁻ is the largest of the six ions.
Exam Tips
- Sketch a trend by plotting one point per element, in order of proton number, and joining the points with straight lines. Label both axes. No numbers are needed, but the relative heights must be right: for melting point, aluminium only slightly above magnesium, silicon the highest, sulfur above phosphorus and argon the lowest; for conductivity, sodium below magnesium below aluminium, silicon low but above zero, and phosphorus to argon at zero.
- Describe means give the direction of the trend and where it changes, with no reasons: "melting point increases from sodium to silicon, falls sharply to phosphorus, rises a little to sulfur, then falls to argon".
- Explain a melting point in four links: the structure, the particles, what is overcome (metallic bonding, covalent bonds or id-id forces) and how much energy that needs. Use the formulas P₄, S₈ and Cl₂.
- Explain conductivity by naming the charge carriers and saying whether they can move. For a non-conductor, state that there are no delocalised electrons or ions free to move.
- Compare two elements by naming both and using a comparative in every link: "S₈ has more electrons than P₄, so the id-id forces between S₈ molecules are stronger, so more energy is needed to overcome them".