Period 3 Physical Properties

The Periodic Table is called periodic because properties repeat: go across Period 3 from sodium to argon and the same pattern of change reappears across Period 2 and Period 4. This note covers the physical side of that pattern for the elements sodium to argon: atomic radius, ionic radius, melting point and electrical conductivity, and explains each one using structure and bonding. It is a favourite source of "describe and explain" questions in Paper 2, often with a graph or table of data to interpret, and of multiple-choice questions in Paper 1.

What periodicity means

Periodicity is the repeating pattern of physical and chemical properties of the elements when they are arranged in order of proton number. Each period starts with a reactive metal, passes through elements with giant structures, and ends with non-metals that exist as small molecules or single atoms.

The reason is electronic. Across a period, electrons are added one at a time to the same outer shell while the nuclear charge rises by one each step. When the outer shell is full, the next element starts a new shell and the pattern begins again.

The Period 3 elements, their outer electrons and their structures are summarised below. Every property in this note can be traced back to this table.

element	Na	Mg	Al	Si	P	S	Cl	Ar
outer electrons	$3s^1$	$3s^2$	$3s^2 3p^1$	$3s^2 3p^2$	$3s^2 3p^3$	$3s^2 3p^4$	$3s^2 3p^5$	$3s^2 3p^6$
structure	giant metallic	giant metallic	giant metallic	giant molecular (giant covalent)	simple molecular, $\ce{P4}$	simple molecular, $\ce{S8}$	simple molecular, $\ce{Cl2}$	monatomic, $\ce{Ar}$
bonding between particles	metallic	metallic	metallic	covalent	id–id forces between molecules	id–id forces between molecules	id–id forces between molecules	id–id forces between atoms

Here "id–id forces" means instantaneous dipole–induced dipole forces (also called London dispersion forces), the weak attractions between all molecules caused by temporary uneven distributions of electrons. Within each $\ce{P4}$, $\ce{S8}$ or $\ce{Cl2}$ molecule the atoms are held by strong covalent bonds; only the forces between molecules are weak.

Atomic radius across Period 3

The atomic radius is measured as half the distance between the nuclei of two bonded atoms of the same element: the metallic radius for Na, Mg and Al and the covalent radius for Si to Cl.

element	Na	Mg	Al	Si	P	S	Cl
atomic radius / nm	0.157	0.136	0.125	0.117	0.110	0.104	0.099

The atomic radius decreases steadily across the period.

Argon is usually left out of the comparison. It forms no bonds, so its radius cannot be measured in the same way; quoted values for noble gases are van der Waals radii, which are not comparable with covalent radii.

Ionic radius across Period 3

The elements on the left of the period form positive ions by losing their outer electrons; those on the right form negative ions by gaining electrons to fill the third shell. (Silicon does not really form a $\ce{Si^4+}$ ion, but a value is quoted for comparison.)

ion	$\ce{Na+}$	$\ce{Mg^2+}$	$\ce{Al^3+}$	$\ce{Si^4+}$	$\ce{P^3-}$	$\ce{S^2-}$	$\ce{Cl-}$
electron configuration	$2,8$	$2,8$	$2,8$	$2,8$	$2,8,8$	$2,8,8$	$2,8,8$
ionic radius / nm	0.095	0.065	0.050	0.041	0.212	0.184	0.181

There are two isoelectronic series (species with the same number of electrons) separated by a large jump.

• Cations, $\ce{Na+}$ to $\ce{Si^4+}$. Each has 10 electrons in two shells. The nuclear charge rises from 11 to 14, so the same number of electrons is pulled in more tightly and the radius falls.
• The jump from $\ce{Si^4+}$ to $\ce{P^3-}$. The anions have an extra shell of electrons (the third shell, now full), so they are much larger than the cations.
• Anions, $\ce{P^3-}$ to $\ce{Cl-}$. Each has 18 electrons in three shells. The nuclear charge rises from 15 to 17, so the radius falls again.

A positive ion is always smaller than its atom (it has lost its outer shell and has more protons than electrons). A negative ion is always larger than its atom (same nuclear charge, more electrons, more repulsion). The general treatment of atomic and ionic radius is in Particles and isotopes; here you only need to apply it to Period 3.

Melting point across Period 3

Melting point measures how much energy is needed to break enough of the forces holding the particles in fixed positions. So the question is always: what forces have to be overcome, and how strong are they? The graph falls into three regions.

Sodium, magnesium and aluminium: giant metallic lattices

The melting point increases from sodium to aluminium.

Metallic bonding is the electrostatic attraction between positive metal ions and delocalised electrons. Across Na, Mg, Al:

• the charge on the metal ion increases ($\ce{Na+}$, $\ce{Mg^2+}$, $\ce{Al^3+}$);
• the number of delocalised electrons per atom increases (1, 2, 3);
• the ionic radius decreases, so the charge is more concentrated and the delocalised electrons are closer to the ion centres.

All three make the metallic bonding stronger, so more energy is needed to overcome it.

Silicon: a giant molecular structure

Silicon has the highest melting point in the period. Each silicon atom is covalently bonded to four others in a tetrahedral arrangement, building a three-dimensional giant molecular (giant covalent) lattice like diamond. To melt silicon, many strong covalent bonds must be broken, which needs a very large amount of energy.

Phosphorus to argon: simple molecules and single atoms

The melting point drops sharply at phosphorus. Phosphorus, sulfur and chlorine exist as small molecules, $\ce{P4}$, $\ce{S8}$ and $\ce{Cl2}$, and argon as single atoms. Melting these only overcomes the weak id–id forces between molecules; the covalent bonds inside the molecules are not broken.

The strength of id–id forces depends on the number of electrons in the molecule (more electrons, larger and more easily distorted electron cloud, larger instantaneous dipoles):

particle	$\ce{S8}$	$\ce{P4}$	$\ce{Cl2}$	$\ce{Ar}$
electrons per particle	$8 \times 16 = 128$	$4 \times 15 = 60$	$2 \times 17 = 34$	18
melting point / K	392	317	172	84

So the order is $\ce{S8} > \ce{P4} > \ce{Cl2} > \ce{Ar}$. Sulfur melts higher than phosphorus because an $\ce{S8}$ molecule is bigger, with more electrons, than a $\ce{P4}$ molecule. The fact that a sulfur atom has more protons than a phosphorus atom is not the reason.

Electrical conductivity across Period 3

To conduct electricity, a substance needs mobile charged particles. In the solid elements, only delocalised electrons can do this.

element	Na	Mg	Al	Si	P	S	Cl	Ar
conductivity	good	better	best	low (semiconductor)	none	none	none	none
reason	delocalised electrons			few electrons free to move; rises with temperature	all electrons held in covalent bonds or atoms

• Na to Al: conductivity increases. The number of delocalised electrons per atom rises from 1 to 3, and the atoms get smaller so there are more delocalised electrons per unit volume. More charge carriers means higher conductivity.
• Si: a semiconductor. Silicon has a giant covalent structure with all four outer electrons in bonds, but a few electrons can be promoted and become free to move. Its conductivity is very low but increases with temperature, unlike a metal.
• P, S, Cl, Ar: non-conductors. Every electron is held either in a covalent bond within a molecule or in a single atom. There are no mobile charged particles.

Periodicity in other periods

The same pattern repeats in Period 2: lithium and beryllium are metals with giant metallic structures, boron and carbon (diamond or graphite) have giant covalent structures with very high melting points, and nitrogen, oxygen, fluorine and neon are simple molecules or atoms with very low melting points. This is what "indicate the periodicity" means: show that the shape of the graph is repeated, and explain it with the same structural changes.

element	Li	Be	B	C	N	O	F	Ne
structure	giant metallic	giant metallic	giant covalent	giant covalent	$\ce{N2}$	$\ce{O2}$	$\ce{F2}$	$\ce{Ne}$
melting point / K (approximate)	454	1560	2350	about 3800 (sublimes)	63	55	53	25

First ionisation energy and electronegativity

Two more Period 3 trends are covered elsewhere but are part of the same story.

• First ionisation energy increases across the period (greater nuclear charge, same shell, similar shielding), with small dips at aluminium ($3p$ electron removed, higher in energy and shielded by $3s$) and sulfur (spin-pair repulsion in a paired $3p$ orbital). See Ionisation energy.
• Electronegativity increases across the period from 0.9 (Na) to 3.0 (Cl), for the same reasons. This drives the change from ionic to covalent bonding in the oxides and chlorides. See Electronegativity, ionic and metallic bonding.

Worked examples

Practice

Answers

1. Na, Mg, Al: giant metallic. Si: giant molecular (giant covalent). P, S, Cl: simple molecular ($\ce{P4}$, $\ce{S8}$, $\ce{Cl2}$). Ar: monatomic (single atoms).
2. Magnesium has one more proton (12 against 11), so a greater nuclear charge. Its outer electrons are in the same (third) shell as sodium's, with similar shielding by the inner shells. The outer electrons are attracted more strongly and held closer to the nucleus.
3. $\ce{Mg^2+} < \ce{Na+} < \ce{Cl-} < \ce{S^2-}$. $\ce{Mg^2+}$ and $\ce{Na+}$ both have 10 electrons in two shells; Mg has more protons so it is smaller. $\ce{Cl-}$ and $\ce{S^2-}$ both have 18 electrons in three shells, so they are much larger than the cations; Cl has more protons (17 against 16) so $\ce{Cl-}$ is smaller than $\ce{S^2-}$.
4. Aluminium contributes three delocalised electrons per atom, sodium only one. Aluminium atoms are also smaller, so there are more delocalised electrons per unit volume. More mobile charge carriers give a higher conductivity.
5. Both are simple molecular, so melting overcomes only id–id forces between molecules. An $\ce{S8}$ molecule has 128 electrons, a $\ce{Cl2}$ molecule only 34. The larger electron cloud of $\ce{S8}$ gives larger instantaneous dipoles and stronger id–id forces, which need more energy to overcome.
6. Silicon has a giant covalent lattice in which all four outer electrons of each atom are used in covalent bonds, so very few electrons are free to move and the conductivity is very low. Heating gives some electrons enough energy to become free, so the conductivity increases with temperature (the opposite of a metal, whose conductivity falls slightly as it is heated).
7. Melting phosphorus does not break the P–P covalent bonds; these are strong and remain intact within the $\ce{P4}$ molecules. Phosphorus has a low melting point because only the weak id–id forces between $\ce{P4}$ molecules are overcome, which needs little energy.
8. The graph rises from Li to a maximum at C, then falls sharply at N and stays very low for N, O, F, Ne (the same shape as Period 3). Carbon is at the peak because it has a giant covalent structure (diamond or graphite): many strong covalent bonds must be broken. Nitrogen is very low because it is made of $\ce{N2}$ molecules held by weak id–id forces.
9. The element with the highest melting point is Si, so Q is phosphorus and its neighbour on the right is sulfur. Both are simple molecular non-conductors; $\ce{S8}$ (128 electrons) has stronger id–id forces than $\ce{P4}$ (60 electrons), so sulfur melts at a higher temperature (392 K against 317 K).
10. On melting, the metal ions only become able to move past one another; much of the metallic bonding remains in the liquid. On boiling, the atoms are separated completely, so all of the metallic bonding must be overcome and the boiling point reflects its strength more directly. Aluminium has stronger metallic bonding because $\ce{Al^3+}$ has a higher charge and a smaller radius than $\ce{Mg^2+}$, and each atom contributes three delocalised electrons rather than two, so the electrostatic attraction between the ions and the delocalised electrons is greater.