9701 · 9.1
Periodicity of physical properties of the elements in Period 3 flashcards
Revision flashcards for Cambridge 9701 Periodicity of physical properties of the elements in Period 3 (syllabus 9.1). Flip, recall, then mark a real past-paper question.
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What is periodicity?
The repeating pattern of physical and chemical properties of elements as you move across a period in the periodic table.
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Why does atomic radius decrease across Period 3?
The number of protons (nuclear charge) increases, but electrons are added to the same principal shell (n=3). This increases the nuclear attraction on the outer electrons, pulling them closer, while shielding remains relatively constant.
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Why does the first ionisation energy generally increase across Period 3?
Increasing nuclear charge and decreasing atomic radius mean the outer electron is held more strongly, requiring more energy to remove.
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Explain the drop in first ionisation energy from Magnesium (Group 2) to Aluminium (Group 13).
Aluminium's outermost electron is in a 3p orbital, which is at a slightly higher energy level and more shielded than the 3s orbital of Magnesium. Therefore, it requires less energy to remove.
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Explain the drop in first ionisation energy from Phosphorus (Group 15) to Sulfur (Group 16).
In Sulfur, the electron being removed is from a paired 3p orbital. Electron-pair repulsion makes it easier to remove this electron compared to removing an electron from a half-filled, more stable 3p sub-shell in Phosphorus.
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Why does the melting point increase from Na to Mg to Al?
These are all metals with metallic bonding. The strength of the metallic bond increases due to an increasing number of delocalised electrons per atom (1, 2, 3) and increasing positive charge on the metal ions, leading to stronger electrostatic attraction.
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Why does Silicon (Si) have the highest melting point in Period 3?
Silicon has a giant covalent (macromolecular) structure. A large amount of energy is needed to break the many strong covalent bonds throughout the lattice.
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Why do P, S, and Cl have low melting points?
They exist as simple discrete molecules (P₄, S₈, Cl₂). Melting only requires overcoming weak intermolecular van der Waals' forces, not strong covalent bonds.
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Why is the melting point of sulfur (S₈) higher than that of phosphorus (P₄)?
Both are simple molecular. However, S₈ is a larger molecule with more electrons than P₄. This leads to stronger temporary dipole-dipole interactions (van der Waals' forces) between S₈ molecules, requiring more energy to overcome.
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Why does electrical conductivity decrease dramatically after Aluminium?
Conductivity requires mobile charge carriers. Metals (Na, Mg, Al) have delocalised electrons. Silicon is a semiconductor. Phosphorus, sulfur, and chlorine are non-metals where electrons are localised in covalent bonds, so they cannot move and carry charge.
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Compare the ionic radius of Na⁺ and Mg²⁺.
Mg²⁺ is smaller than Na⁺. Both are isoelectronic (have the same electron configuration, 1s²2s²2p⁶), but Mg has 12 protons while Na has 11. The greater nuclear charge of Mg pulls the same number of electrons in more tightly.