- 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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- 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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- 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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- 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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