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CHEMISTRY (PERIODIC TABLE (TRENDS (IONIC RADIUS (IONIC RADIUS OF ANION IS…

- - - - ATOMIC RADIUS INCREASES DOWN THE GROUP BECAUSE the number of filled principle quantum shells increases. This results in the valence electrons being increasingly further and less strongly attracted to the nucleus.
      - ATOMIC RADIUS DECREASES ACROSS THE PERIOD BECAUSE valence electrons are added to the same principal quantum shell, thus the increase in screening effect is negligible. The increase in nuclear charge outweighs the negligible increase in screening effect, causing effective nuclear charge to increase. The resulting stronger electrostatic forces of attraction between the nucleus and valence electrons cause the valence electrons to be pulled closer to the nucleus, effectively decreasing the atomic radius.
    - - IONIC RADIUS OF ANION IS LARGER THAN ATOMIC RADIUS OF THE SAME ELEMENT BECAUSE as valence electrons are added to the atom to form anions, there is greater interelectronic repulsion between valence electrons, while nuclear charge remains the same due to their being same number of protons in the atom. Valence electrons are hence further away from the nucleus. Weaker electrostatic forces of attraction between the nucleus and valence electrons cause anion radius to be larger than its atomic radius
      - IONIC RADIUS OF ISOELECTRONIC IONS DECREASES ACROSS PERIOD BECAUSE isoelectronic ions have a similar number of electrons. Hence, screening effect is similar. However, nuclear charge increases across the period due to the increasingly larger number of protons in each atom. The increase in effective nuclear charge causes there to be stronger electrostatic forces of attraction between the nucleus and valence electrons, causing valence electrons to be pulled closer to the nucleus, decreasing ionic radius.
      - WITHIN SAME GROUP, IONIC RADIUS OF CATIONS ARE SMALLER THAN THAT OF ANIONS because anions have an additional electronic shell when compared to cations. This results in a smaller screening effect for cations, outweighing the negligible decrease in nuclear charge, causing effective nuclear charge to increase. Stronger electrostatic forces of attraction between nucleus and electrons pulls valence electrons closer to the nucleus, causing the ionic radius of cations to be smaller than the anions.
      - IONIC RADIUS OF CATION IS SMALLER THAN ATOMIC RADIUS OF THE SAME ELEMENT BECAUSE as valence electrons are removed from the atom to form a cation, screening effect decreases as there is one less quantum shell of electrons, while nuclear charge remains the same due to the same number of protons in the atom. Effective nuclear charge increases, hence, the stronger electrostatic forces of attraction between the nucleus and valence electrons in the cation causes valence electrons to be pulled closer to the nucleus.
    - - FIRST IONISATION ENERGY DECREASES DOWN THE GROUP BECAUSE the number of filled quantum shells increases. Thus electrons are increasingly further away from the nucleus and are less strongly attracted to the nucleus. Smaller amount of energy required to remove the valence electrons.
      - FIRST IONISATION ENERGY INCREASES ACROSS THE PERIOD BECAUSE across the period, electrons are added to the same principal quantum shell, hence, there is a negligible increase in screening effect. As increase in nuclear charge outweighs the negligible increase in screening effect, effective nuclear charge increases. There is then stronger electrostatic forces of attraction between the nucleus and valence electrons. More energy is hence required to remove the valence electrons.
    - - ELECTRONEGATIVITY DECREASES DOWN THE PERIOD BECAUSE the number of filled quantum shells increases. Hence, distance between the nucleus and the bonding electrons decrease. Thus, the ability of the atom to attract the bonding electrons decreases, decreasing electronegativity down the group.
      - ELECTRONEGATIVITY INCREASES ACROSS THE PERIOD because across the period, electrons are added to the same principal quantum shell, resulting in a negligible increase in screening effect. The increase in nuclear charge outweighs the negligible increase in screening effect. Hence, effective nuclear charge increases, increasing an atom's ability to attract bonding electrons, increasing electronegativity.
    - - GIANT METALLIC STRUCTURE
        
        Large amount of energy required to overcome strong metallic bond
        
        Melting point increases as metallic bond strength increases
        
        number of delocalised valence electrons per atom increases across the period
        
        Cationic charge increases, cationic radius decreases
        
        Stronger electrostatic forces of attraction between ion and sea of delocalised valence electrons which require larger amount of energy to overcome.
      - GIANT MOLECULAR STRUCTURE
        
        Very large amount of energy required to overcome the strong and extensive covalent bonds between atoms in the 3 - dimensional structure
      - SIMPLE MOLECULAR STRUCTURE
        
        Small amount of energy is required to overcome the weak instantaneous dipole - induced dipole moments between molecules
  - - - IONIC OXIDES
        
        LARGE ELECTRONEGATIVITY DIFFERENCE results in electron transfer (aka ionic bond)
        
        HIGH MELTING POINTS: Large amounts of energy required to overcome the strong electrostatic forces of attraction between oppositely charged ions
        
        SOME COVALENT CHARACTER for AL2O3 as AL has a high charge density, Hence, AL3+ polarises the electron cloud of O2-
        
        ELECTRICAL CONDUCTIVITY: Presence of mobile ions in the liquid state hence are good conductors of electricty
      - GIANT MOLECULAR OXIDES
        
        SMALL ELECTRONEGATIVITY DIFFERENCE results in electron sharing (aka covalent bond)
        
        HIGH MELTING POINTS: Large amount of energy required to overcome the strong and extensive covalent bonds in the 3 - dimensional structure.
        
        POOR ELECTRICAL CONDUCTIVITY: absence of mobile charge carriers
      - SIMPLE MOLECULAR OXIDES
        
        SMALL ELECTRONEGATIVITY DIFFERENCE results in electron shring (aka covalent bonding)
        
        LOW MELTING POINTS: Small amount of energy required to overcome the weak intermolecular forces of attraction between molecules
        
        NO ELECTRICAL CONDUCTIVITY: absence of mobile charge carriers