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Bonding (Born-Haber Cycles (Enthalpy (Electron affinity (First electron…
Bonding
Born-Haber Cycles
Hess's law: The enthalpy change accompanying a chemical change is independent of the route by which the chemical change occurs.
Ionic bonding
Forming ionic bonds is usually extremely endothermic because the electrostatic attractions between the cations and anions are very strong
*Low ionisation energy
Ionisation energy is the energy required for an isolated gaseous atom, to remove the outermost electron, forming a positive cation. The lower the ionisation energy, the more favourable the formation of the positive cation, which is needed to attract the negative anion.
High electron affinity
Electron affinity is the amount of energy released, when an isolated gaseous atom accepts an electron to form a negative anion. The higher the electron affinity, the more favourable the formation of the anion, which is needed to attract the positive cation
Large lattice energy
The electrostatic forces of attraction between cations and anions occur because of the coulombic forces of attraction. These forces release a certain amount of energy when an ionic bond is formed, also known as the lattice energy. If the coulombic attractions are stronger, then more energy gets released, the lattice energy is greater, and a more stable ionic bond is formed.
High melting/boiling points
Because of the many strong electrostatic attractions between cations and anions, ionic crystal lattices have high melting and boiling points, which results in a high lattice energy as a lot of energy is released when these bonds are formed.
Brittle
Because of the many strong electrostatic forces, ionic crystals are hard. However, they are also brittle because with a large force one layer of ions is forced to shift relative to its neighbour. When that happens, it brings ions of the same charge next to each other, causing them to repel and shatter the crystal. When an ionic crystal breaks, it tends to do so along smooth planes because of the regular arrangement of the ions 
Conduction of electricity
Ionic crystals cannot conduct electricity whilst solid as the ions are unable to move and carry charge. In liquid/molten form however, frees the ions to move and carry charge, conducting electricity.
Lattice energy
Ionic charge
The higher the charge, the higher the lattice energy as the effective nuclear charge increases, requiring more energy to break the bonds
Ionic radius
As ionic radius increases, the attractive force between the nucleus and outermost electron decreases, requiring less energy to remove the outermost electron, lowering the lattice energy
Measuring lattice energy
Lattice energy cannot be measured directly because it is impossible for gaseous ions to form a solid crystal directly
Enthalpy
Electron affinity
First electron affinity
Standard enthalpy change when a mole of gaseous atoms is converted to a mole of gaseous ions, each with a single negative charge.
Second electron affinity
Standard enthalpy change when a mole of electrons is added to a mole of gaseous ions each with a single charge to form ions each with two negative charges.
Ionisation energy
First ionisation energy
Standard enthalpy change when one mole of gaseous atoms is converted into a mole of gaseous ions each with a single positive charge.
Second ionisation energy
Standard enthalpy change when one mole of electrons is lost from a mole of singly charged positive ions.
Atomisation
Atomisation
The enthalpy change which accompanies the formation of one mole of gaseous atoms from the element under its standard state conditions.
The lower the enthalpy of formation (for the ionic compound), the more stable it is
Enthalpy of formation for MgCl2 is much lower than that of MgCl, making it more stable and hence being the formula for magnesium chloride.
Polarity
Electronegativity
is a measure of the ability of an atom to attract electrons. The most electronegative elements are on the top right of the period table.
Electronegativity down groups decreases because with each group an extra shell is added, increasing shielding and resulting in weaker electrostatic forces.
Electronegativity across periods increases because the nucleus gets bigger since a proton is added each time, but there is no extra shielding because they are being added to the outer shell, so the electrostatic forces become stronger.
Bonds
The bigger the electronegativity difference, the more polar the bond.
Non-polar covalent: 0-0.4
no dipole, therefore low electronegativity difference
Polar covalent: 0.5-1.9
has a dipole so relatively high electronegativity difference
Ionic: 2.0-4.0
completely separated, very high electronegativity difference
Dative covalent:
a bond where both of the electrons from one atom are being shared across two atoms
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Polarisation
is the distortion of the electron cloud in a molecule or ion by a nearby charge. To achieve maximum distortion, the cation must have a small atomic radius and a string charge, and the anion must have a large atomic radius and a strong charge.
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