Chapter 3: The Periodic Table

1) By increasing atomic number 2) By period (horizontal rows) 3) By group (vertical columns) 4) By metal/metalloid/non-metal - Metal: Alkali metals, metals, transition metals - Metalloids: have properties of both metals and non-metals (demarcated by ‘staircase’ boundary at B) - Non-metals: halogens and noble gases

Metallic structure: lattice of cations in a sea of delocalised electrons - Metals: able to lose electrons easily (low ionisation energies) to form cations • Across a period: ionisation energies increase, elements lose electrons less readily (metals usually formed by elements on left side of periodic table) • Down a group: ionisation energies decrease, elements lose electrons more readily (metals usually formed by elements at the end of a group (eg. group 14) METALS: 1) Large atomic radii 2) Low ionisation energies 3) Low (exothermic) electron affinity 4) Low electronegativity

• Highest filled energy level: 3 • Number of electrons at highest filled energy level: 6 3s²3p⁴

ATOMIC RADIUS: size of an atom (half the internuclear distance of a molecule of that element) Determined by: 1) Nuclear charge: number of protons (creating “attractive force”) 2) Shielding effect: repulsion of valence electrons by inner electrons (number of “core” non-valence electrons) 3) Effective nuclear charge (nuclear charge - shielding effect) Increases down a group: - Atoms have more electron shells down a group - Nuclear charge and shielding effect both increase, counteracting any effect caused by a greater number of protons Decreases across a period: - Nuclear charge increases (no. of protons increase) - Shielding effect remains the same (no. of non-valence electrons remain the same) - Effective nuclear charge increases (outer electrons are more strongly attracted to the nucleus) - Atomic radii decreases

IONIC RADIUS: size of an ion Cationic radii: smaller than its atomic radii - Cations have fewer electron shells than their atoms - Fewer electrons - Less repulsion between electrons, valence electrons are more strongly attracted to the nucleus - Smaller electron cloud Anionic radii: larger than its atomic radii - Constant nuclear charge - More electrons - More repulsion between electrons - Larger electron cloud Across isoelectronic series (Si⁴⁻, P³⁻, S²⁻, Cl⁻): - Increase in nuclear charge - Constant screening effect - Increase in effective nuclear charge - Valence electrons are more strongly attracted to nucleus - Ionic radii decreases

FIRST IONISATION ENERGY: energy required to remove the outermost electron from a gaseous atom M (g) –> M⁺ (g) + e⁻ Decreases down the group: - Increase in atomic radius - Valence electrons are less strongly attracted to the nucleus - Less energy required to remove the outermost electron Increases across a period: - Increase in nuclear charge - Constant screening effect - Increase in effective nuclear charge - Smaller atomic radius - Valence electrons are more strongly attracted to the nucleus - More energy required to remove the outermost electron *Exception across Period 2: 1) Group 2 (Be) and Group 13 (B) - B has a smaller AR, and is supposed to have higher IE. However, because the electron being removed from B is at the 2s sub-shell and the electron removed from Be is at the 2p sub-shell (higher energy level), B has a lower IE than Be 2) Group 15 (N) and Group 16 (O) - O has a smaller AR and is supposed to have higher IE. However, because O has 2 paired electrons in the 2p-orbital generating interelectronic repulsion whereas N only has 1 electron (2p³), O has a lower IE

ELECTRON AFFINITY: energy change when one electron is added to a gaseous atom X (g) + e⁻ –> X⁻ (g) - Usually exothermic for all elements: bond forming is exothermic Decreases down a group: - Increase in AR - Electron will be less strongly attracted - Releases less energy when electron is brought in - Less exothermic *Exception: Group 17 Cl has the most exothermic electron affinity (supposed to be F) - However, due to the small size of the F atom, it generates large amounts of interelectronic repulsion, requiring more energy to attract an electron and releasing less energy Increase across a period: - Decrease in AR - Electron will be more strongly attracted to nucleus when brought in - Releases more energy - More exothermic *Exception: Period 3 - P (group 15) is supposed to have more exothermic electron affinity than Si (group 14) - However, P has 3 electrons in 3 separate p-orbitals (3p³), which means that bringing in an extra electron would generate interelectronic repulsion, weakening the attraction between the nucleus and the electrons, making its electron affinity less exothermic

ELECTRONEGATIVITY: measure of the attraction of an atom in a molecule for the electron pair in the covalent bond of which it is a part Decreases down a group: - Increase in AR - Increase in bonding distance between nucleus and electron pair - Weaker attraction between electron pair and nucleus - Decreased electronegativity Increases across a period - Decrease in AR - Decrease in bonding distance between nucleus and electron pair - Stronger attraction between nucleus and electron pair - Increased electronegativity

All elements down a group: - Similar chemical properties - Gradation in physical properties Group 1 metals: - Metallic bonding: electrostatic attraction between the lattice of cations and the sea of delocalised electrons Melting point decreases down the group: - IR increases - Nucleus is further away from delocalised electrons - Weaker attractive force between nucleus and delocalised electrons - Less energy required to break apart the lattice and overcome the IMFA between particles REACTIONS: - Alkali metals are all highly reactive: 1) With oxygen (to form an oxide) 2) With water (to form a base and hydrogen gas) • Get more reactive down the group: AR increases, IE decreases

• Diatomic molecules • Bonding in halogens: covalent bonding (weak intermolecular forces of id-id) to form simple molecular structure Boiling point increases down the group: - Mᵣ increases - London forces between molecules grow stronger - More energy required to overcome them Reactivity decreases down the group: 1) With alkali metals (to form salts) 2) With other halogens (displacement reaction)

OXIDES: - Basic (react with water to form alkaline solutions IF SOLUBLE + react with acids to form salts) - Acidic (react with water to form acidic solutions IF SOLUBLE + react with bases to form salts) - Amphoteric (react with both acids and bases to form salts + does not react with water) - Neutral Across a period: Basic (metallic oxides) –> Amphoteric (some metal oxides) –> Neutral (some non-metallic oxides) -> Acidic (non-metallic oxides)

Sodium oxide (Na₂O): Basic Na₂O + H₂O –> 2 NaOH Magnesium oxide (MgO): Basic MgO + H₂O –> Mg(OH)₂ Aluminium oxide (Al₂O₃): Amphoteric Al₂O₃ + 6H⁺ –> 2Al³⁺ + 3H₂O Al₂O₃ + 2OH⁻ + 3H₂O –> 2Al(OH)₄⁻ Phosphorous pentoxide (P₄O₁₀): Acidic P₄O₁₀ + 6 H₂O –> 4 H₃PO₄ Sulfur dioxide (SO₂)/sulfur trioxide (SO₃): SO₂ + H₂O –> H₂SO₃ SO₃ + H₂O –> H₂SO₄