Trends in Physical Properties Across Periods and Down Groups – Trends in Physical Properties of Elements in the Periodic Table (Senior 4 Chemistry) | Ganzaa

Trends in Physical Properties of Elements in the Periodic Table

Lessons

Senior 4 · Chemistry

Trends in Physical Properties of Elements in the Periodic Table

3 lessons

Lesson 3 of 3

Trends in Physical Properties Across Periods and Down Groups

25 min

Eight elements, four kinds of solid

Walk along Period 3 from sodium to argon and the elements change completely: a soft metal that melts in hot water, two harder metals, a shiny grey semiconductor that melts above 1400 °C, a waxy white solid, a yellow powder, a poisonous green gas and an unreactive gas. Each step adds one proton and one electron. This lesson explains how such small steps produce such large changes, and why going down a group changes properties more gradually. The two keys are the forces on the outer electrons (Units 2 and 3) and the type of structure each element forms (Units 3 and 4).

Silicon, the metalloid in the middle of Period 3. Left: a solar cell made of multicrystalline silicon. Right: a rod (top) and chunks (bottom) of polycrystalline silicon, shiny grey like a metal but brittle, and a semiconductor rather than a true conductor. Images: Georg Slickers, Warut Roonguthai and NITOL (compilation), CC BY-SA 3.0, via Wikimedia Commons

Definition

Covalent radius: half the distance between the nuclei of two identical atoms joined by a single covalent bond. Metallic radius: half the distance between neighbouring nuclei in a metal crystal. Ionic radius: the radius of an ion in an ionic crystal.

Metalloid: an element with properties between those of metals and non-metals, such as silicon, a shiny solid that is a semiconductor. Isoelectronic species have the same number of electrons.

1. Atomic radius

The size of an atom is taken from bonded atoms, using the covalent radius for non-metals and the metallic radius for metals. Noble gases form no bonds, so their radii are measured differently and are left out of comparisons.

Across a period the radius decreases. The nuclear charge rises by one each step, but the new electron goes into the same shell, where it shields the other outer electrons very little. The effective nuclear charge increases, and the outer shell is pulled closer: Na 166 pm to Cl 102 pm.
Down a group the radius increases. Each element has one more shell, further from the nucleus, and more inner electrons to shield it: Li 128, Na 166, K 203, Rb 220, Cs 244 pm.

2. Ionic radius

A positive ion is smaller than its atom: it has usually lost its whole outer shell, and the remaining electrons are pulled in by the same nuclear charge (Na 166 pm, $\mathrm{Na^{+}}$ 102 pm).
A negative ion is larger than its atom: the extra electrons repel each other and the outer electron cloud expands (Cl 102 pm, $\mathrm{Cl^{-}}$ 181 pm).
In an isoelectronic series (same number of electrons), the ion with the most protons is the smallest.

Worked example 1: explaining sizes

Explain why (a) the atomic radius of magnesium (141 pm) is smaller than that of sodium (166 pm); (b) $\mathrm{Mg^{2+}}$ (72 pm) is much smaller than a Mg atom; (c) $\mathrm{Mg^{2+}}$ is smaller than $\mathrm{Na^{+}}$ (102 pm).

Solution.

(a) Both have outer electrons in shell 3 with the same inner shells (shielding about the same), but magnesium has 12 protons against 11, so its outer electrons are pulled closer.
(b) $\mathrm{Mg^{2+}}$ has lost both 3s electrons, so its outermost shell is now shell 2, much closer to the nucleus; 12 protons now hold only 10 electrons.
(c) The ions are isoelectronic (10 electrons each), but $\mathrm{Mg^{2+}}$ has 12 protons and $\mathrm{Na^{+}}$ has 11, so the electrons of $\mathrm{Mg^{2+}}$ are pulled in more tightly.

3. Ionisation energy, electronegativity and electron affinity

These follow from the same factors as radius. Across a period, the effective nuclear charge rises and the atom shrinks, so first ionisation energy (Unit 2) and electronegativity (Unit 4) both increase; down a group they decrease. The first electron affinity becomes generally more exothermic across a period, reaching its largest values in Group 17. Down Group 17 it becomes less exothermic from chlorine (−349) to bromine (−325) to iodine (−295 kJ mol⁻¹), because the added electron is further from the nucleus. Fluorine (−328) is an exception: its atom is so small that the incoming electron is strongly repelled by the electrons already there.

Table 3.1: Period 3

Element	Na	Mg	Al	Si	P	S	Cl	Ar
Structure	giant metallic			giant covalent	$\mathrm{P_4}$ molecules	$\mathrm{S_8}$ molecules	$\mathrm{Cl_2}$ molecules	single atoms
Melting point / °C	98	650	660	1414	44	115	−101	−189
Boiling point / °C	883	1090	2519	3265	280	445	−34	−186
Electrical conductivity	good	good	very good	semiconductor	none (insulators)
Electronegativity	0.9	1.3	1.6	1.9	2.2	2.6	3.2	–
First ionisation energy / kJ mol⁻¹	496	738	578	787	1012	1000	1251	1521

4. Melting and boiling points across Period 3

Na, Mg, Al: giant metallic lattices. The melting and boiling points rise as the number of delocalised electrons per atom rises and the ions get smaller and more highly charged (Unit 3).
Si: a giant covalent lattice, like diamond, in which every atom is joined to four others by strong covalent bonds. Melting means breaking many of these bonds, so silicon has the highest melting point in the period.
P, S, Cl: simple molecules held together only by weak London forces. Little energy separates the molecules. The order follows the number of electrons per molecule: $\mathrm{S_8}$ (128 electrons) > $\mathrm{P_4}$ (60) > $\mathrm{Cl_2}$ (34).
Ar: single atoms with only 18 electrons and the weakest London forces: the lowest melting point.

Worked example 2: predicting a property down a group

The melting points of the Group 1 metals are Li 181, Na 98 and K 63 °C. Predict whether rubidium melts above or below 63 °C, and explain. The boiling points of the Group 18 gases are He −269, Ne −246, Ar −186 and Kr −153 °C. Predict the trend for xenon.

Solution. Down Group 1, each metal has one delocalised electron per atom, but the ions get larger, so the attraction between the ions and the electron sea weakens and the melting point falls. Rubidium should melt below 63 °C (it melts at 39 °C).

Down Group 18, the atoms have more electrons, so London forces get stronger and boiling points rise. Xenon should boil above −153 °C (it boils at −108 °C). The two groups show opposite trends because their bonding is different.

5. Metallic character and conductivity

Elements on the left and lower down the table lose electrons easily (low ionisation energy and electronegativity): they are metals, with delocalised electrons that conduct. Across a period, metallic character decreases: Na, Mg and Al are metals, silicon is a metalloid (a semiconductor whose conductivity rises with temperature) and P, S, Cl and Ar are non-metals that do not conduct. Down a group metallic character increases: in Group 14, carbon is a non-metal, silicon and germanium are metalloids, and tin and lead are metals.

Warning

Common mistakes. (1) Saying atoms get smaller across a period "because they have more electrons": the cause is the greater nuclear charge with nearly constant shielding. (2) Comparing the melting points of P, S and Cl by their covalent bonds: melting only separates molecules, so compare London forces. (3) Saying silicon conducts like a metal: it is a semiconductor. (4) Using the radius of an atom when the question asks about an ion.

Tip

Two-line answers for any trend: across a period, "nuclear charge increases, electrons enter the same shell, shielding is similar, so the outer electrons are held more tightly"; down a group, "more shells, outer electrons further from the nucleus and more shielded, so they are held less tightly". For melting points, start instead from the structure.

MaChemGuy: Periodic Patterns in Melting Points Across Period 3 (12 min). Watch how each element's structure, metallic, giant covalent or simple molecular, is used to explain its melting point.

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Practice

Define covalent radius. Explain why noble gases are left out when comparing atomic radii.
Explain why the atomic radius decreases from lithium to fluorine.
Put these in order of increasing radius and explain: $\mathrm{S^{2-}}$ , $\mathrm{Cl^{-}}$ , $\mathrm{K^{+}}$ , $\mathrm{Ca^{2+}}$ (all have 18 electrons).
Explain why silicon has a much higher melting point than phosphorus.
Explain why the melting point of sulfur is higher than that of chlorine.
Astatine is below iodine in Group 17. Predict, with reasons, its state at room temperature and whether its electronegativity is greater or less than iodine's.
Describe and explain how the electrical conductivity of the elements changes across Period 3.

Show answersHide answers

Half the distance between the nuclei of two identical atoms joined by a single covalent bond. Noble gases do not form bonds, so their radii are measured from atoms that only touch, which gives larger values that cannot fairly be compared.
The nuclear charge increases from 3 to 9 while the electrons are added to the same shell (n = 2) with the same inner shell for shielding; the effective nuclear charge rises and pulls the outer shell closer.
$\mathrm{Ca^{2+}}$ < $\mathrm{K^{+}}$ < $\mathrm{Cl^{-}}$ < $\mathrm{S^{2-}}$ : the same 18 electrons are held by 20, 19, 17 and 16 protons respectively; more protons pull them closer.
Silicon is a giant covalent lattice: melting breaks many strong covalent bonds. Phosphorus is made of $\mathrm{P_4}$ molecules held by weak London forces, which need little energy to overcome.
Both are simple molecular. $\mathrm{S_8}$ has 128 electrons per molecule against 34 for $\mathrm{Cl_2}$ , so its London forces are much stronger.
Down Group 17 the molecules have more electrons and stronger London forces (F₂ and Cl₂ gases, Br₂ liquid, I₂ solid), so astatine should be a solid. Electronegativity decreases down the group (more shells, more shielding), so astatine's is less than iodine's.
Na, Mg and Al conduct well (delocalised electrons; Al best, with 3 per atom); Si is a semiconductor; P, S, Cl and Ar do not conduct because all their electrons are held in covalent bonds or in single atoms.

Quick check

Why does the atomic radius decrease from sodium to chlorine?

Key points

Across a period: nuclear charge rises, shielding is similar, so atomic radius decreases and ionisation energy and electronegativity increase.
Down a group: more shells and more shielding, so radius increases and ionisation energy and electronegativity decrease.
Cations are smaller and anions larger than their atoms; in an isoelectronic series, more protons means a smaller ion.
Period 3 melting points follow structure: metallic (rising Na to Al), giant covalent Si (highest), simple molecular P, S, Cl (low), monatomic Ar (lowest).

1Eight elements, four kinds of solid

21. Atomic radius

32. Ionic radius

43. Ionisation energy, electronegativity and electron affinity

54. Melting and boiling points across Period 3

65. Metallic character and conductivity