Periodicity

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  • Periodicity
    • Periodic Table
      • Elements arranged in order of atomic number
      • Groups = Same number of valence electrons
      • Bromine and mercury only liquids
      • Nobles gases (inert gases), nitrogen, oxygen, chlorine, hydrogen and fluorine only gases
      • Carbon is a solid at room temperature
      • Periods = Same number of energy levels
        • Which is main energy level
      • Lanthanoids and actinoids omitted - actinoids start with actinium
      • Metals formed from elements with lower ionisation energies (on the left and bottom of groups) to readily form positive ions and a sea of delocalised electrons
      • Metals = Low ionisation energies, less exothermic elecron affinities, large atomic radii and low electronegativity
      • Periodic table arranged according to the highest energy subshell occupied by electrons
    • Physical Properties
      • Atomic radius
        • Increases down group
          • Number of electron shells increases
          • Nuclear charge increases down a group but so does electron repulsion to counter
        • Decreases across period
          • Nuclear charge increases to attract outer electrons closer to nucleus
            • Increase in nuclear charge but no real change in shielding because same number of shells
        • Half the internuclear distance
      • Ionic Radii
        • Positive ions are smaller than element
          • More electrons repelling same nuclear charge in elemental form
        • Negative ions are bigger than element
          • More electrons for same nuclear charge expands electron cloud as more repulsion
        • Metals and non-metals decrease across period due to increase in nuclear charge
      • Ionisation Energy
        • Decreases down group
          • Increase in electron shells so outer electrons further from nucleus and less attraction.  More shielding
            • Increase in nuclear charge balanced by shielding
        • Increases across period
          • Nuclear charge increases so electrons more strongly attracted to the nucleus
          • Shielding the same.  Electrons in same electron shell do not shield each other very well
      • Electron Affinity
        • First electron affinity is exothermic
        • Data incomplete as hard to measure experimentally
        • Affinity decreases down the group (mainly from chlorine down)
          • Less exothermic down the group
          • Increase in number of electron shells and atomic radii means attraction is weaker to added electrons
        • Across a period
          • More exothermic across period
          • Increases across a period
            • Increase in nuclear charge and smaller atomic radius means stronger attraction to added electron
      • Electronegativity
        • Decreases down a group
          • Increase atomic radius due to more electron shells and more shielding so less attraction
        • Increases across a period
          • Nuclear charge increases across a period so atomic radius is smaller so more attraction
            • No significant change in shielding due to same number of inner shells
        • The attraction of an atom in a molecule for the electron pair in the covalent bond for which it is part of
      • Properties of elements in group 1 and group 17
        • Group 1
          • Soft, reactive, low melting points
          • Reactivity increases down the group
            • Atomic radius and number of electron shells increases and more shielding so less attraction for outer electron
            • Ionisation energy decreases down group
            • React vigorously in oxygen and tarnish rapidly in air
            • Reactions with water
              • Sodium fizzes, turns into a ball (melts) and moves on the surface
              • Potassium bursts into a lilac flame
              • Caesium explodes on contact
              • Strong bases and ionise completely
          • Melting point decreases down the group
            • Increases in atomic radius and larger atoms means less strong attraction between positive nuclei and delocalised electrons so less energy to break apart the lattice
          • Liquid sodium used as a coolant in some nuclear reactors
        • Same group = SIMILAR chemical properties because all have the same number of valence electrons
        • Group 17
          • Reactivity decreases down group
            • Atomic radius increases so larger atom so less strong attraction to outer electron to be added
          • Melting point increases down the group
            • London forces increase as atomic mass increases so more energy required to break intermolecular forces between molecules
          • React with alkali metals to form white or colourless salts
            • Alkali metal chlorides, bromides and iodides are soluble and produce colourless, neutral solutions
          • Displacement reactions
            • Bromine displaced gives orange solution
            • Iodine displaced produces dark red-brown solution
            • Chlorine is pale yellow-green solution
            • Redox reaction and more reactive halogen oxidises less reactive halide ion
              • Stronger affinity for electrons so removes them
      • Oxides
        • Metallic oxides = Basic
        • Non-metallic oxides = acidic
        • Alkali metal + water --> metal hydroxide
        • Metal oxide + acid --> salt + water
        • Aluminium oxide is amphoteric
          • Reacts with acids and bases to form salts
            • Forms Al3+ with acid and aluminium hydroxide with bases
          • Does not react with water
          • Insoluble
        • Magnesium oxide sparingly soluble and only slightly alkaline
        • NO is produced in an internal combustion engine, is insoluble and forms neutral solutions
          • Contributes to photochemical smog
        • NO2 forms nitric (V) acid with water
        • N2O also a neutral oxide
    • D-block elements
      • Transition element:  An element that forms at least one stable ion with a partially filled d subshell
      • Zinc has a full d subshell, does not exhibit some of the properties of transition elements, has one stable ion and does not form coloured compounds
      • Scandium is regarded as a transition element but has no d subshell and nearly always forms a 3+ ion and occasionally a 2+ ion but the bonding is complicated and not always a 2+ ion is involved
      • Atomic radius and first ionisation energy minimal change across the period
      • Properties
        • Formed coloured compounds/ complexes
        • Compounds can be magnetic
        • Metals and their complexes can act as catalysts
          • Finely divided iron acts as a catalyst in the Haber process
          • Normally homogeneous catalysts (same phase as reactants) but iron is a heterogeneous catalyst in the Haber process (different phase to reactants)
          • Depends on transition metal having varying oxidation numbers and coordinating with other molecules/ ions
        • Can form complex ions
        • Can exhibit multiple oxidation numbers in compounds/ complexes
          • 4s and 3d subshells close in energy so no large jump in ionisation energy when they are removed
            • Number of electrons lost depends on lattice enthalpy, hydration enthalpy and ionisation energy
        • High melting points and densities - typical metals
      • 4s electrons removed before 3d electrons to form ions
      • Maximum oxidation number increases as number of electrons in 4s and 3d subshells increases
      • Always have oxidation state +2
        • 4s electrons
      • Magnetism
        • Paramagnetism = Unpaired electrons and are attracted by a magnetic field.  More dominant
        • Diamagnetism = Paired electrons repel a magnetic field slightly
        • More unpaired or paired electrons, the greater the magnetic moment
        • Compounds of copper (I) are diamagnetic only
      • Complex Ions
        • Coordinate covalent bond where ligand is Lewis base and ion is Lewis acid
        • Ligand:  Negative ions or neutral molecules that have a lone pair of electrons
        • Hexaaquairon(II) ion
        • All transition metals apart from titanium can form octahedral complexes with the water ligand with a 2+ charge
        • Can undergo substitution reactions to replace ligands (add concentrated HCl to add Cl- ions)
          • Reversible reactions
          • In the presence of concentrated HCl, copper (II) sulfate solution goes from blue to green
        • CO is a neutral ligand.  Br- is a negatively charged ligand (all halogens)
    • Coloured Complexes
      • In the gaseous state, all d orbitals are degenerate in transition metal
      • When surrounded by ligands, 2 orbitals in transition metal raise to higher energy and 3 to lower energy
        • Repulsion between ligand lone pairs and electrons in d orbitals
        • 2 d orbitals face ligands so are raised.  3 are between ligands so are lowered relative to the other 2 orbitals
      • Absorption of light energy promotes electron from lower to higher orbitals so light emitted shows everything but the colour of the light absorbed (shows complementary colour)
        • MUST BE IN VISIBLE SPECTRUM
      • REQUIRES A PARTIALLY FILLED D SUBSHELL
        • Need a partially filled d subshell to allow an electron to be promoted (to have electrons to promote and for space for them to be promoted to)
        • Otherwise compounds will be colourless
      • Factors affecting colour complexes
        • Greater difference in energy between lower and higher orbitals means higher frequency of light absorbed
        • Transition metal ion
          • Manganese 2+ is pale pink/ colourless but Fe2+ is pale green
          • Different electron configurations and different levels of repulsion and therefore splitting
          • Higher nuclear charge means ligands pulled in more closely, more repulsion and greater splitting
        • Oxidation number
          • Greater oxidation number, greater splitting
          • Different electron configurations
          • Greater charge on ion means ligands closer, more repulsion, more splitting
        • Nature of ligand
          • Spectrochemicals series of how much ligands split orbitals
          • Stronger field ligands = Greater splitting ligands
          • Neutral molecules generally cause more splitting than ions apart from CN- which is the strongest with CO
          • Charge density = Charge per unit volume
            • F- and I- have same charge but F- is smaller so more repulsion and greater splitting of d orbitals
          • Pi bonding between lone pair and d orbitals reduces splitting so more pi bonding, less splitting
            • Pi electron donation effect
            • CO has electron density in cloud from metal ion donated to it, meaning there is a greater chance of splitting

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