Chemistry (9701) | Periodicity of physical properties of the elements in Period 3

The Period 3 Elements: Unlocking the Patterns in Physical Properties (Syllabus 9.1)

Welcome! This chapter is all about spotting and explaining patterns. The Periodic Table is organized beautifully, and Period 3 (Sodium, Na, through Argon, Ar) gives us a perfect snapshot of how structure and bonding fundamentally change across a row.

Understanding these trends will not only boost your grades but also help you predict the behaviour of elements you haven't even studied yet. Don't worry if the explanations seem complex at first; we will break them down using simple structure and bonding concepts!

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Section 1: Trends in Atomic and Ionic Radius

When we move across Period 3, we are filling electrons into the same principal quantum shell (n=3). This means the amount of electron shielding remains relatively constant.

1.1 Atomic Radius (Size of the Atom)

Trend: Atomic radius decreases from Na to Cl.
(Argon is often excluded from the trend line as its radius is usually defined differently, or it's larger due to its weak van der Waals forces).

Explanation (The Domino Effect):

• As you move from Na (\( Z=11 \)) to Cl (\( Z=17 \)), the number of protons in the nucleus (the Nuclear Charge) increases steadily.
• Each element gains one more proton, making the nucleus more positive.
• The outer electrons are in the same shell (n=3), meaning the shielding effect by the inner 10 electrons (\( 1s^2 2s^2 2p^6 \)) stays constant.
• Result: The stronger positive nuclear charge exerts a much greater attractive force on the outer electron shell, pulling it inwards.
• Key Takeaway: Stronger effective nuclear charge = smaller atom.

1.2 Ionic Radius (Size of the Ion)

The trend for ionic radius is more complex because we must distinguish between cations (positive ions, Na to Al) and anions (negative ions, P to Cl).

Trend:

1. Cations (\( \text{Na}^+ \), \( \text{Mg}^{2+} \), \( \text{Al}^{3+} \)) decrease in size.
2. There is a large jump in radius when moving from cations to anions.
3. Anions (\( \text{P}^{3-} \), \( \text{S}^{2-} \), \( \text{Cl}^- \)) then decrease in size.

Explanation:
A. Cations (\( \text{Na}^+ \) to \( \text{Al}^{3+} \)):

• These ions all lose their 3s electrons, meaning they all have the electron configuration of Neon (\( 1s^2 2s^2 2p^6 \)). They are isoelectronic (same number of electrons: 10).
• However, the nuclear charge is still increasing (11+ for Na to 13+ for Al).
• The stronger nuclear charge in \( \text{Al}^{3+} \) pulls the smaller, remaining electron shells in much tighter than in \( \text{Na}^+ \). Therefore, \( \text{Al}^{3+} \) is the smallest cation.

B. Anions (\( \text{P}^{3-} \) to \( \text{Cl}^- \)):

• These ions gain electrons (P gaining 3, S gaining 2, Cl gaining 1) to achieve the Argon configuration (\( 1s^2 2s^2 2p^6 3s^2 3p^6 \)). They are also isoelectronic (same number of electrons: 18).
• Since they have more electrons than protons, the electron-electron repulsion increases significantly, causing the electron cloud to expand (this is why anions are much larger than their parent atoms).
• Moving from \( \text{P}^{3-} \) (15 protons) to \( \text{Cl}^- \) (17 protons): the nuclear charge increases while the number of electrons (18) remains the same.
• This stronger nuclear pull across the anions causes the ionic radius to decrease again (making \( \text{Cl}^- \) the smallest anion).

Quick Review: Atomic and ionic radii decrease across Period 3 because of the steadily increasing nuclear charge acting on electrons in the same principal quantum shell.

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Section 2: Trends in Melting Point (MP)

The trend in melting point is not smooth; it depends entirely on the structure and bonding type, which changes dramatically across Period 3. This section requires careful explanation of the structure for each group of elements.

2.1 Group 1, 2, 13: Metals (Na, Mg, Al)

Structure:
Giant Metallic Lattice (Positive metal ions surrounded by a sea of delocalised electrons).

Trend: Melting point increases from Na to Al (Na < Mg < Al).

Explanation:

• To melt a metal, you must overcome the metallic bonds (the electrostatic attraction between the cations and the delocalised electron sea).
• Moving from Na to Al, three factors increase the bond strength:
  1. Cation Charge: Increases from \( \text{Na}^+ \) (1+) to \( \text{Mg}^{2+} \) (2+) to \( \text{Al}^{3+} \) (3+).
  2. Number of Delocalised Electrons: Increases from 1 to 2 to 3 electrons per atom.
  3. Cation Size: Decreases (as established in 1.1).
• A smaller ion with a higher charge pulling on a larger electron sea results in a much stronger attraction. Aluminum, with its strong \( \text{Al}^{3+} \) core and 3 delocalised electrons, has the strongest metallic bonds and thus a higher melting point.

2.2 Group 14: Silicon (Si)

Structure:
Giant Molecular Structure (Giant Covalent Lattice).
(Think of it like diamond—each Si atom is bonded to four others by strong covalent bonds).

Trend: Silicon has the highest melting point in Period 3.

Explanation:

• To melt Silicon, all the strong covalent bonds linking the vast network of atoms must be broken.
• This requires a massive amount of energy (analogous to dismantling an entire skyscraper brick by brick!). This is why its melting point is drastically higher than the metals and non-metals.

2.3 Group 15, 16, 17, 18: Non-Metals (P, S, Cl, Ar)

Structure:
Simple Molecular Structure.
(e.g., \( \text{P}_4 \), \( \text{S}_8 \), \( \text{Cl}_2 \), \( \text{Ar} \)).

Trend: Melting points are all very low, dropping significantly after Si.

Explanation:

• Within these molecules, the bonds are strong covalent bonds (e.g., the P–P bonds inside the \( \text{P}_4 \) molecule).
• However, when you melt these substances, you are not breaking the strong covalent bonds; you are only breaking the weak intermolecular forces (van der Waals forces) that hold the separate molecules together.
• Because these intermolecular forces require very little energy to overcome (analogous to separating sticky notes rather than ripping the notes themselves), their melting points are low.

The Sub-Trend (P, S, Cl, Ar):

The melting points of these simple molecular elements reflect the size of the molecule, which influences the strength of the van der Waals forces:

• \( \text{S}_8 \) (Octa-sulfur, largest molecule) has the highest MP among this group.
• \( \text{P}_4 \) (Tetra-phosphorus, next largest) is lower than S.
• \( \text{Cl}_2 \) (Diatomic molecule) is lower.
• \( \text{Ar} \) (Monatomic, smallest) has the lowest MP.

Memory Aid: For the non-metals, remember the order S (sulfur) > P (phosphorus) > Cl (chlorine) > Ar (argon). This generally follows increasing molecular mass/size.

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Section 3: Electrical Conductivity

Electrical conductivity is determined by the presence of mobile charge carriers (either delocalised electrons or free ions).

3.1 Metals (Na, Mg, Al)

Trend: Conductivity is high and increases from Na to Al.

Explanation:

• Metals contain a sea of delocalised valence electrons.
• When a voltage is applied, these delocalised electrons are free to move and carry charge.
• Conductivity increases because Aluminum provides three delocalised electrons per atom, compared to two for Mg and one for Na. More mobile charge carriers mean higher conductivity.

Did you know? Aluminum is an excellent conductor and is widely used in high-voltage transmission lines because it is lighter and cheaper than copper.

3.2 Semi-metal (Silicon, Si)

Trend: Conductivity is very low at room temperature, but increases significantly upon heating or doping (a semiconductor).

Explanation:

• Silicon is a giant covalent structure. At absolute zero, all four valence electrons of each Si atom are held firmly in strong covalent bonds. There are no delocalised electrons.
• However, due to the energy gap between the bonding electrons and the conduction band, a few electrons gain enough energy (e.g., from heat) to break free and act as charge carriers.
• This makes Silicon a semiconductor, essential for modern electronics (e.g., computer chips).

3.3 Non-Metals (P, S, Cl, Ar)

Trend: These elements are all insulators (do not conduct electricity).

Explanation:

• These elements form simple molecular structures or are monatomic (Ar).
• All their valence electrons are localized, either fixed in the covalent bonds within the molecules (\( \text{P}_4 \), \( \text{S}_8 \), \( \text{Cl}_2 \)) or tightly held by the nucleus (\( \text{Ar} \)).
• Since there are no mobile charge carriers (no delocalised electrons or free ions), they cannot conduct electricity.

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Summary Table of Period 3 Physical Properties

This table provides a quick reference for revision, focusing on the explanations (9.1.2) derived from structure and bonding.

Element	Structure Type	Bonding Type	Melting Point Trend	Electrical Conductivity Trend
Na, Mg, Al	Giant Metallic	Metallic	Increases (Na < Mg < Al) due to stronger metallic bonding (higher charge, more delocalised electrons).	High (Increases Na < Mg < Al) due to increasing number of mobile delocalised electrons.
Si	Giant Molecular	Covalent	Maximum, very high, due to breaking numerous strong covalent bonds.	Low (Semiconductor) - electrons fixed in covalent bonds, few break free at room temp.
P, S, Cl, Ar	Simple Molecular	Covalent (within molecule) + van der Waals (between molecules)	Very low, falls sharply, due to breaking only weak intermolecular forces.	Non-conductors (Insulators) - all electrons are fixed/localised within molecules or atoms.

You've successfully mapped the physical landscape of Period 3! The key to mastering this topic is linking every observed property (size, MP, conductivity) back to the underlying nuclear charge, electron shells, structure, and bonding. Keep practicing those explanations!