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Topic 3 and 13 - Periodicity (3.2 Periodic Trends (Physical Properties…

- - - - Measured as half the distance between neighbouring nuclei
      - Increases down a group as number of electron shells increases
      - Decreases across the period as increasing nuclear charge creates a larger attraction between nucleus and electrons. Significant difference
    - - Negative ions are larger that parent atoms as electrons are added to the outer shell. Increased electron repulsion caused electrons to move further apart
      - Decreases across the period (from 1-14 for +ve and from 14-17 for -ve) as nuclear charge increases, pulling outer electron closer
      - Positive ions are smaller than negative ions - +ve have two occupied electron principal energy levels -ve have three
      - 5.Increases down a group as no.
        of electron energy levels increases.
      - Positive ions are smaller than parent atoms as outer shell is lost
    - - The energy required to remove one mole of electron from one mole of gaseous atoms in their ground state.
      - Increases across a period - increase in effective nuclear charge = increase in attraction between nucleus and electrons = more difficult to remove electron
      - Decreases down a group - effective nuclear charge remains the same, distance between electron and nucleus increases = easier to remove
      - Exceptions
        Group 13 elements with nS2 nP1 config have lower first ionization energies than Group 2 with nS2 config - p orbitals have higher energy than s orbitals.
        Drop between group 15 & 16, group16 electron is removed from double filled p orbital = easier to remove due to repulsion between electrons
    - - The energy change when 1 mole of gaseous negative ions when 1 mol of electrons is added to 1 mole of gaseous atoms. It is a measure of the electron attracting ability of an atom.
      - First affinity is exothermic (electrons attracted to positive nucleus)
        Second affinity is endothermic (electrons repelled by negative ion)
    - - Relative measure of the attraction that an atom has for a shared pair of electrons when covalently bonded to another atom. Also a measure of attraction between the nucleus and outer electron (bonding electrons)
      - Increases across a group - as nuclear charge increases = increased attraction
      - Decreases down a group - as ionic / atomic radii increases = further away from nucleus = less attraction.
      - Does not apply to group 18 - do not form covalent bonds
      - Increasing reactivity from bottom left to top right
      - Highly electronegative elements = most exothermic electron affinities
    - - Decreases down group 1 - Metallic structures held together by delocalized electrons. Attraction between these electron and positively charged ions decreases down group
      - Increases down group 17 - held together by London dispersion forces which increases with number of electrons.
      - Generally rise across period & reach max at group 14 then fall to min at group 18.
        Bonding changes from metallic --> giant covalent --> van der Waals between simple molecules --> single atoms.
    - - Increases as protons are added to successive elements
        i.e from left to right across the periodic table
      - Outer electrons are shielded from nucleus and repelled by inner electrons
      - Effective charge - experience by outer electrons, not full nuclear charge due to shielding and rellesion
      - Effective Nuclear Charge =
        Nuclear charge - No. Shielding Electrons
      - Down the group - increase in nuclear charge canceled out by increasing orbitals
      - Given by atomic number
    - - Metals have lower ionization energies and electronegativities than non-metals.
      - Valence electrons can move away from nucleus = metals can conduct electricity
      - Diagonal band of metalloids - have similar electronegativities
  - - - don't form anions as extra electrons would have to be added to an empty outer energy level, where they would experience a negligible nuclear pull-up
      - colorless gases
      - very unreactive (because they cannot gain/lose electrons)
      - monoatomic
      - don't form cations as they have the highest ionization energies
    - - Too reactive to be found in nature. Stored in oil to prevent reactions with air/water
      - good conductors of electricity and heat
      - low densities
      - form ionic compound with non-metals
      - Reactions with water
        
        alkiali metal + water -----> hydrogen + metal hydroxide
        
        Resulting solution is Alkaline
        
        Sodium - vigorous release of H, heat produced melts unreacted metal which forms small ball that floats on surface of water
        
        Potassium - more vigorous than sodium, enough heat to ignite hydrogen produced (lilac coloured flame), moves excitedly on surface
        
        Lithium - floats, reacts slowly, releases H, keeps its shape
    - - Reacton with group 1 Metals
        
        form ionic halides
        
        electrostatic force bonds ions together
        
        Most vigorous reaction between elements furthest apart on periodic table
      - Displacement Reactions
        
        Can displace less reactive halogens, as more reactive halogens are better oxidising agents hence they gain electrons from the halogen that they oxidise.
        Cl2 + 2Br ^ - -----------> 2Cl ^ - + Br2
      - The Halides
      - Reactivity decreases down the group because the outer shell is at increasingly higher energy levels and further from the nucleus. This decreases the attractive forces that pull an electron into the valence shell of a halogen, decreasing its reactivity.
      - coloured
      - gradual change from gases (F,Cl) --> liquid (Br) --> solids (I, At)
      - form ionic compound with metals and covalent compounds with non-metals
      - very reactive - very exothermic electron affinities and high effective nuclear charge = pull on electrons from other atoms