Electrons, Bonding & Structure

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  • Created by: ernily
  • Created on: 09-04-15 17:11

Electronic Structure

  • Electrons move around the nucleus in shells. Shells are given principal quantum numbers.
  • Shells contain sub-shells. Sub-shells have orbitals which can hold up to 2 electrons.
  • An orbital is the bit of space that an electron moves in.
  • Electrons in orbitals have to spin in opposite directions.
  • 's' orbitals are spherical, 'p' orbials are dumbbell shaped.
  • Electron fill up the lowest energy sub-shells first.
  • They fill singly before sharing.
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Ionisation Energies

  • First Ionisation Energy: The energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions.
  • Nuclear Charge: The more protons in the nucleus, the more positively charged the nucleus, so the stronger the attraction for electrons.
  • Distance From Nucleus: Attraction falls rapidly with distance. The closer to the nucleus, the stronger the attraction.
  • Shielding: Number of electrons between outer electrons and nucleus increases, attraction decreases.
  • Second Ionisation Energy: The energy needed to remove 1 electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions.
  • Within each shell, successive ionisation energies increase. Electrons are being removed from an increasingly positive ion. There's less repulsion amongst electrons, so they're held more strongly by the nucleus.
  • The big jumps in ionisation energy happen when a new shell is broken into. An electron closer to the shell is being removed
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Ionic Bonding

  • Ions are formed when electrons are transferred from one atom to another.
  • An ionic bond is an electrostatic attraction between two oppositely charged ions.
  • Sodium Chloride & Magnesium Oxide are ionic compounds.
    • In NaCl, the single + charge on the Na+ ion balances the single - charge on the Cl- ion.
    • In MgCl2, the 2+ charge on the Mg2+ ion balances the two - charges on the two Cl- ions.
  • Not all ions are made from single atoms:
    • NO3- : Nitrate
    • CO3 2- : Carbonate
    • SO4 2- : Sulfate
    • NH4+ : Ammonium
  • In NaCl, the ions are packed together in a giant ionic lattice, which has very strong ionic bonds. Lots of energy is needed to overcome the bonds.
  • Ionic compounds conduct electricity when molten or dissolved as the ions are free to move in a liquid and they carry a charge.
  • Ionic compounds have high melting points as lots of energy is needed to overcome the strong electrostatic forces.
  • Ionic compounds dissolve in water because water molecules are polar - they pull the ions away from the lattice, causing it to dissolve.
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Covalent Bonding

  • Molecules are groups of atoms bonded together.
  • A covalent bond is a shared pair of electrons.
  • Some covalent compounds, like sulfur hexafluoride, use 'd' orbitals to expand the octet.
  • Some atoms share more than one pair of electrons.
  • Oxygen has a double bond.
  • Nitrogen can triple bond.
  • Dative Covalent Bonding is where both electrons come from one atom.
  • Ammonium (NH4+) is formed when the nitrogen atom donates a pair of electrons to a proton (H+)
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Giant Covalent Lattices & Metallic Bonding

  • Giant covalent lattices are huge networks of covalently bonded atoms.
  • Diamond is made up of carbon atoms. Due to strong covalent bonds, it:
    • Has a very high melting point
    • Is hard
    • Is a good thermal conductor
    • Can't conduct electricity
    • Is insoluble
  • Graphite has a different structure to diamond:
    • There are weak bonds between the layers; graphite is slippery.
    • There are delocalised electrons between the layers, so graphite can conduct electricity.
    • The layers are far apart, so graphite is less dense.
    • It is insoluble and has a very high melting point.
  • Metals have giant metallic lattice structures.
  • Delocalised 'sea' of electrons, and positive metal ions.
  • More delocalised electrons = higher melting point.
  • There are no bonds, so metal ions slide past eachother. This makes them malleable.
  • Delocalised electrons carry kinetic energy, so they are good thermal conductors.
  • Delocalised electrons carry a current, so they are good electrical conductors.
  • Metals are insoluble due to the strength of the bonds.
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Shapes Of Molecules

  • Molecular shape depends on electron pairs around the central atom.
  • Electron pairs repel each other
    • They are all negatively charged, so they repel as far as possible.
    • Lone pair/Lone pair bond angles are the biggest.
    • Lone pair/Bonding pair bond angles are the second biggest.
    • Bonding pair/Bonding pair bond angles are the smallest.
  • VALENCE SHELL ELECTRON-PAIR REPULSION THEORY.
  • 2 electron pairs = Linear, 180 degrees.
  • 3 electron pairs = Trigonal Planar, 120 degrees.
  • 4 electron pairs and no lone pairs = Tetrahedral, 109.5 degrees.
  • 4 electron pairs and 1 lone pair = Pyramidal, 107 degrees.
  • 4 electron pairs and 2 lone pairs = Non-Linear, 104.5 degrees.
  • 6 electron pairs = Octahedral
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Electronegativity & Intermolecular Forces

  • Covalent bonds may be polarised by differences in electronegativity.
  • In a covalent bond between atoms of different electronegativities, bonding electrons are pulled towards the more electronegative atom.
  • This bond is polar.
  • Diatomic gases are non-polar because of the equal electronegativities.
  • In a polar bond, the difference in electronegativities causes a dipole.
  • Dipole: A difference in charge between two atoms caused by a shift in electron density in the bond.
  • The delta + and delta - charges on polar molecules cause weak electrostatic forces between molecules; permanent dipole-dipole interactions
  • Intermolecular forces are very weak.
  • Hydrogen bonding is the strongest. It only happens when H is covalently bonded to N,O, or F.
  • Hydrogen has a high charge density because it is small, so N,O, and F are very electronegative.
  • This bond is polarised.
  • Hydrogen bonding causes substances to be soluble and have high boiling & melting points.
  • Ice is less dense than liquid water. The water molecules are held together in a lattice by hydrogen bonds. The hydrogen bonds are long, which causes ice to be less dense.
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Van der Waals Forces

  • Electrons in charge clouds are always moving really quickly.
  • When electrons in an atom are more to one side, the atom has a temporary dipole.
  • The dipole can cause another temporary dipole. The dipoles are then attracted to each other.
  • This causes another dipole, and so on.
  • Stronger van der Waals forces means higher boiling points.
  • Covalent bonds don't break during melting and boiling.
  • To melt/boil a covalent compound, you only have to overcome the van der Waals, or the hydrogen bonds.
  • You don't need to break the covalent bond.
  • Covalent compounds have low melting and boiling points.
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