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Chem
Electrons, bonding, and structure
| Question | Answer |
|---|---|
| Number of electrons in first 4 shells (shell number = n) | 2n^2 |
| Orbital definition | A region of space around the nucleus which can hold up to 2 electrons with opposite spins |
| S-orbitals | Spherical Each shell from n=1 has 1 s-orbital Radius increases as n increases Maximum 2 electrons in subshell |
| P-orbitals | Dumb-bell shaped, each at a right angle to each other Each shell from n=2 has 3 p-orbitals Distance from nucleus increases as n increases Maximum 6 electrons per subshell |
| D-orbitals and f-orbitals | Each shell from n=3 has 5 d-orbitals (max 10 electrons per subshell) Each shell from n=4 has 7 f-orbitals |
| Filling of orbitals | Fill in order of increasing energy levels (e.g 1s 2s 2p 3s 3p) Important note that 4s is lower energy than 3d so filling is 3p 4s 3d Same energy orbitals are filled singularly first |
| D block ions electron configuration | 4s orbital empties first |
| Ionic bonding | Electrostatic attraction between positive and negative ions |
| Ionic compound structure | Giant ionic lattice as a result of each ion being surrounded by oppositely charged ions in all directions |
| Ionic compounds melting/boiling points | High melting/boiling points - many ionic bonds which require lots of energy to overcome the strong electrostatic forces between ions in the lattice Ions with greater charges have stronger attraction so higher melting/boiling points |
| Ionic compounds solubility | Soluble in polar solvents Polar solvent molecules break down the lattice and attract and surround the ions Solubility decreases as ionic charge increases |
| Ionic compounds electrical conductivity | Does not conduct when solid but does when molten/aqueous Solid - ions are fixed in lattice, no charge carriers can move Molten/aqueous - lattice breaks down and ions which can carry charge are free to move |
| Covalent bonding | Strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms |
| Average bond enthalpy (covalent bonds) | As average bond enthalpy increases, the stronger the covalent bond |
| Molecule shapes and bond angles (from four electron pairs) | Tetrahedral - 4 bonding pairs, 0 lone pairs, 109.5 Pyramidal - 3 bonding pairs, 1 lone pair, 107 Non-linear - 2 bonding pairs, 2 lone pairs, 104.5 Electron pairs around a central atom repel as far as possible, lone pairs repel more than bonding |
| Molecule shapes and bond angles (from other than four pairs) | Linear - 2 bonding pairs, 180 Trigonal planar - 3 bonding pairs, 120 Octahedral - 6 bonding pairs, 90 |
| Electronegativity | The ability of an atom to attract the bonding electrons in a covalent bond Increases as nuclear charge increases Increases as atomic radius decreases |
| Polar bond | Formed when an electron pair is shared unequally due to a difference in electronegativity between bonded atoms One side will be slightly positive, other will be slightly negative, the separation of opposite charges is a dipole |
| Polar molecules | Formed when there is a dipole in the bonds of a molecule and the molecule is not symmetrical causing the dipoles to cancel out |
| Intermolecular forces categories | Permanent dipole-dipole interactions Induced dipole-dipole interactions (London forces) Hydrogen bonding |
| London forces | Movement of electrons in a molecule create a changing instantaneous dipole - this instantaneous dipole induces a dipole in neighbouring molecules causing them to attract More electrons in the molecule, the larger the dipole + greater the attraction |
| Simple molecular substances structure | Simple molecular lattice when solid, molecules held by weak intermolecular forces |
| Simple molecular substances properties (melting point, solubility) | Low boiling/melting point - little energy needed to break IM forces Solubility - non-polar solutes dissolve in non-polar solvents (IM forces between solute and solvent break lattice), polar solutes dissolve in polar solvents (polar bonds attract) |
| Simple molecular substances properties (electrical conductivity) | Electrical conductivity - no free moving charged particles so no conductivity |
| Hydrogen bonding | Permanent dipole-dipole interaction between lone pair on electronegative atom and a hydrogen atom attached to an electronegative atom (e.g -OH, -NH, HF) |
| Anomalous properties of water | Ice less dense than water - H bonds hold molecules apart in a lattice in ice further than they do in water, the lattice also has holes decreasing density Relatively higher boiling/melting point - more energy needed to break H bonds than London forces |
Created by:
silver54331