Bonding

• Ionic bonding

  •Bonding between metals and nonmetal. The metal loses its electron to the nonmetal, creating ions.

  •Oppositely charged ions attract through electrostatic force to form a giant, ionic lattice

  E.g. $SO_4^{2-},OH^-,NO_3^-,CO_3^{2-},NH_4^+$
• Covalent bonding

  • bonding between 2 or more non metals

  • electrons are shared between the atoms
• dative bonding

  • forms when both the electrons in a shared pair comes from a single atom

  • shown using an arrow and reacts like a covalent bond
• Metallic bonding

  • it’s a lattice of positively charged ions surrounded by a sea of delocalised electrons

  • the greater the charge on the positive ions, the stronger the force of attraction as there’s more electrons released in the sea

  • larger ions in size have weaker attraction due to greater atomic radius
• what determines physical properties of compounds?

  Type of bonding, crystal structures.

  Examples of physical properties include boiling/melting point, solubility and conductivity
• Ionic crystal structure

  •high melting/boiling point due to strong electrostatic forces holding the lattice structure together

  •can conduct electricity when molten/in solution as ions are free to move

  • often brittle as when alternating layers distort, like charges repel (fragnentation)
• Metallic crystal structure

  •good conductor as the sea of delocalised electrons can carry charge

  •it’s malleable as the layers of positive electrons slide over each other.

  •high melting point and it’s almost always solid at room temp due to the electrostatic force of attraction between positive ions and delocalised electrons
• Simple molecular crystal structure


  •low melting/boiling points as simple molecular structures are covalently bonded, held together with weak van der waal forces

  •poor conductors due to lack of charged particles
• Macromolecular crystal structures 


  •very high melting points as atoms have multiple covalent bonds due to giant lattice structure. This makes them rigid.

  •DIAMOND: made of carbons covalently bonded to 4 other carbons

  •GRAPHITE: each carbon is bonded to 3 others in a flat sheet. Electrons not used in bonding are released as free electrons, meaning it can conduct electricity.
• Molecule shapes

  
• Molecule shapes 2

  
• Molecule Shapes 3

  
• What is electronegativity?

  •The power of an atom to attract negative charges towards itself within a covalent bond.

  •power depends on an atoms size and nuclear charge

  •it increases along a period as atomic radius decreases

  •it decreases down a group as shielding increases
• permanent dipole forces

  •when 2 atoms with different electro negativities bond, a polar bond is formed

  •the more electronegative atom draws the negative charge towards itself and away from the other atom, creating a $\delta+$and $\delta-$region.
• Induced dipole

  •can be formed when the electron orbitals around a molecule are influenced by another charged particle
• Van der Waals

  •weakest intermolecular force. It acts as an induced dipole between molecules.

  •The strength depends on a molecules mr/shape

  • larger mr = stronger van der waals

  • straight chains have stronger VDW than branched as they can line up and pack close together, reducing the distance over which the force acts.
• Permanent dipole


  •Acts between molecules with a polar bond.

  •$\delta+$and $\delta-$attract and hold molecules in a lattice like structure.
• Hydrogen bonding


  •strongest intermolecular force. Only form between hydrogen and the three most electronegative atoms (N,O,F)

  •lone pairs form a bond with a hydrogen from another molecule (shown with a dotted line)

  •molecules with hydrogen bonding have much higher melting /boiling points compared to similar sized molecules without hydrogen bonding.