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4.1 Ionic Bond

4.1.1 : Ionic bond

+ve (cations) and -ve (anions) ions are attracted to each other and form a continuous ionic lattice

4.1.2

Group 1 metals form +1 ions, group 2 metals form +2 ions, metals in group 3 for +3 ions...funny that :)

Examples: Li⁺, Mg²⁺, Al³⁺

Greater ease of ionisation Li --> Cs is due to the increased electron shielding of the nuclear attraction caused by additional inner shells of electrons. The easier atoms are to ionise, the more reactive they will be because less energy is required to ionise them, and so they react faster.

4.1.3

Group 6 ions will form 2- ions, Group 7 ions will form - ions. Examples : O^-2, Cl^-.

4.1.4

The transitions metals (elements from Ti to Cu, ignore Sc and Zn) can form multiple ions (ie Fe²⁺, Fe³⁺) (due to proximity of 4s and 3d shells)

4.1.5

The ionic or covalent nature of the bonding in a binary compound is a result in the difference between their electronegativity.

NaCl_(s) is ionic, HCl_(g) is (polar) covalent (also, covalent molecules tend to be gases/liquids, ionic tends to be solid - except network covalent which will be solid). In general, if the difference between electronegativities is greater than 1.7, the bond will be more than 50% ionic.

4.1.6

Take the name of the group 1,2, or 3 metal and add fluoride, chloride, bromide, iodide etc , oxide, sulfide etc - Nitride and phosphide - how exciting :)

4.2 Covalent Bond

4.2.1

Covalent bonds are where two atoms each donate 1 electron to form a pair held between the two atoms...Such bonds are generally formed by atoms with little difference in electronegativity...ie C, H and O in organic chemistry.

4.2.2

All electrons must be paired...Lewis diagrams are the element symbol with the outer (valence) shell of electrons left over and spare electrons pair up...in general C forms 4 bonds, N forms 3, O forms 2, halogens form 1, H forms 1...(Li would form 1, Be 2, and B 3 but they don't usually...metallic or ionic bonding)

4.2.3

Electronegativity values range from 0.7 to 4...from bottom left to top right respectively (hydrogen falls B and C with a electronegativity of 2.1...

4.2.4

When covalent molecules have a difference in electronegativity (between the two bonding atoms) then the pair will be held closer to the more electronegative atom...resulting in a small -ve charge on the more electronegative atom, and a small +ve charge on the other...results in polar bonds

4.2.5

Shape of molecule with 4 electron pairs depends on number of lone pairs.

3 lone pairs --> linear, 2 lone pairs --> bent, 1 lone pair --> triagonal pyramid, No lone pairs --> tetrahedral

4.2.6

The polarity of a molecule depends on both the shape and the polarity of the bonds...1) if there are no polar bonds, it's not polar. 2) if there are polar bonds, but the shape is symmetrical, it's not polar (think about it like 3D vector addition...if they add to zero, then it's not polar). 3) if there are polar bonds, and it's not symmetric, then the molecule is polar

4.3 Intermolecular forces

4.3.1

Van der Waal's forces - Electrons will not be evenly spread around an atom/molecule at any given time, meaning the molecule will have a slight +ve charge on one end, and a -ve at the other. this temporary state may cause attraction between two molecules, pulling them together (also known as london dispersion forces).

Dipole-dipole forces - Polar molecules, when properly oriented, will attract each other as a result of this. Stronger than van der Waal's forces.

Hydrogen bonding - When hydrogen is bonded to nitrogen, oxygen or fluorine, a very strong dipole is formed, making the hydrogen very strongly positive. This hydrogen is then attracted to the lone pairs on other similar molecules (nitrogen, oxygen and fluorine all have lone pairs) forming a hydrogen bond, which is stronger than van der Waal's or dipole-dipole, but weaker than covalent bonding.

4.3.2 : Structural features

Nonpolar molecules...van der Waal's forces only - also present in all other molecules, though it's strength is insignificant compared to the others. Polar molecules...dipole-dipole forces arise from polar bonds and asymmetry in molecules.

Hydrogen bonds result from hydrogen bonded as described above. This results in molecules with hydrogen bonding exhibiting stronger intermolecular forces, ie higher boiling/melting points etc. ie H₂O has a higher bp then H₂S due to hydrogen bonding, and so on down the strength list. (nonpolar molecules don't conduct electricity, polar +hydrogen bonding ones will...I suppose this goes here)

4.4 Metallic bond

4.4.1 : Metallic bonding

The metal atoms lose their outer electrons which then become delocalized, and free to move throughout the entire metal. These -ve delocalized electrons hold the metal cations together strongly. Since these electrons can flow, atoms with metallic bonding exhibit high electrical conductivity. Unlike ionic bonding, distorting the atoms does not cause repulsion so metallic substances are ductile (can be stretched into wires) and malleable (can be made into flat sheets). The free moving electrons also allow for high thermal conductivity, and the electrons can carry the heat energy rather than it being transferred slowly through atoms vibrating.

4.5 Physical Properties

• Melting and Boiling point - High with Ionic and metallic bonding (and network covalent), Low with covalent molecular bonding.

• Volatility - Covalent molecular substances are volatile, others aren't

• Conductivity - Metallic substances conduct. Polar molecular substances conduct, non-polar ones don't. Ionic substances don't conduct when solid, do conduct when molten or dissolved in water.

• Solubility - Ionic substances -> generally dissolve in polar solvents (like water). Metallic substances...(Generally not soluble ?). Non-polar molecules are generally soluble in non-polar solvents, and polar in polar. Organic molecules with a polar head : Short chain molecules are solubility in polar solvents, long chains can eventually outweigh the polar 'head' and will dissolve in non-polar solvents.

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