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Bonding (Born-Haber Cycles (Describe quantitatively the energetic factors…
Bonding
Born-Haber Cycles
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Ionic compounds:
- are hard crystalline substances
- have high melting and boiling points
- don't conduct electricity when in solid because their ions cannot move away from fixed positions in the giant lattice.
The lattice energy for an ionic solid is the standard enthalpy change when one mole of the compound forms from free gaseous ions (or the other way around). The lattice enthalpy is always negative (exothermic) when the lattice is being formed as bonds are being made, releasing energy. It is positive (endothermic) when the lattice is being broken as the bond breaking takes in energy.
Two factors influence lattice energy:
- the charges on the ions (higher charge = higher lattice energy
- the ionic radii (which affects the distance between the ions) (smaller ionic radius = higher lattice energy)
Lattice energy cannot be measured directly because it is impossible for gaseous ions to form a solid crystal directly.
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.
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: 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 of the ionic compound, the more stable it is. For example, the enthalpy of formation of MgCl is much higher than the enthalpy of formation of MgCl2, even though they are both exothermic. This is why when magnesium and chlorine react, they form MgCl2 and not MgCl; it is more stable.
Polarity
Electronegativity: ability to attract electrons in a covalent bond
The most electronegative elements are on the top right of the periodic table.
You can determine the type of bond between two atoms by calculating the difference in electronegativity values between the elements. The bigger the electronegativity difference, the more polar the bond.
Non-polar covalent: 0-0.4
Polar covalent: 0.5-1.9
Ionic: 2.0-4.0
Polar covalent bonds: relatively high electronegativity difference, has a dipole (distance between delta + and delta -)
Non-polar covalent bonds: no dipole (no distance between delta + and delta -) therefore very low electronegativity difference
Ionic bonds: no sharing of electrons, very high electronegativity difference
Dative covalent bonds (coordinate bond): a bond where there is a pair of shared electrons between two atoms. The difference relative to a covalent bond is that in a dative covalent bond these electrons both come from one atom.
Electronegativity across periods increases because the nucleus gets bigger (more protons added) but there is no extra shielding so the electrostatic forces are stronger and the electrons are more attracted to the nucleus.
Electronegativity down groups decreases because an extra shell is added each time, increasing shielding and so the electrostatic forces are weaker.
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Polarisation is the distortion of the electron cloud in a molecule or ion by a nearby charge. The polarisation power of a cation depends on its charge and its radius and the polarisation power of an anion depends on its size. The larger the negative ion and the larger its charge, the more polarisation power it has. So, iodide ions have a higher polarisation power than fluoride ions.
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