Reactivity

• Formal charge = valence number - number of bonds - number of free electrons
• Sterics = shape
• Electronics = electron density
• We generate structures that there is full octets and that FC is as close to zero as possible
• Charge assigned to an atom within a molecule assumes e- are shared equally in covalent bonds
• Cations are more reactive than anions because theyre attracted to e-
• the more electronegative the greater capacity to accommodate -ve charge, when comparing reactivities compare the capacity to accommodate charge, greater capacity = less reactive.
• SP2 = Trigonal planar
• SP = Linear
• SP3 = tetrahedral
• Reactions occur above and below the plane, due to e- density
• if an molecule has an empty P orbital, it may accept a pair of e- making it a lewis acid eg BF3
• hybridisation does not change through resonance
• Steric number = number of sigma bonds + lone pairs of e-
• Steric number = 4 -> SP3, if also not involved any e- movement (resonance)
• Steric number = 3, SP2
• Steric number = 2 , SP
• In resonance movements it is much better to place the -ve charge on the most electronegative atom and the +ve on the most electropositive
• Maximum bonding is preferred, the most pi bonds possible and least amount of charges possible
• A compound is less reactive when you can delocalize e- due to evenly spread charge
• To weigh up what is more reactive out of a selection, consider resonance and what will create the largest partial positive charge

  • does it have resonance? How many structures (more = more stable)
  • Whats more electronegative (more electronegative = harder to delocalise e- = more reactive)
• Triple bonds are more stable, this is due to it being SP hybridised with 50% S character, which is more than SP2 (67%) and more again than SP3 (75%). Having a higher S character means that the atoms will hold their electrons closer to their nucleus, providing more stable bonds.
• PKa = How acidic a molecule is, the lower the pka the more acidic. the equation for PKa is concentration of conjugates/ concentration of reactants.
• Strong acid -> weak conjugate base
• Weak acid -> strong conjugate base
• Lewis Acid = E- Acceptor
• Lewis base = E- Donor
• Bronstead Acid - h+ donor
• Bronstead base - h+ acceptors
• Steps to identify acid strengths:

  1. Identify all H environments
  2. Draw conjugate bases
  3. Observe resonance
  4. Consider stability
• more resonance = more stability
• delocalization of e- in resonance occurs in parallel P orbitals
• Hockles rule = 4n + 2(pi) e-
• equilibrium will favour direction of the weaker acid
• When an acid donates a proton it becomes its conjugate base
• to draw the conjugate bases of a molecule, find all h environments, draw a structure for each environment and ensure that it misses a proton from one of the h environments.
• To determine which Hydrogen is more acidic on an acid, look at the conjugate bases. The most acidic hydrogen will form the weakest conjugate base. This means the conjugate base with the most resonance structures will be associated with the most acidic hydrogen.
• PKa of a whole molecule is similar to the PKa of the most acidic hydrogen
• Tautomerization: Tautomers are structural isomers of chemical compounds which readily interconvert. The chemical reaction interconverting the two is called tautomerization.
• Tautomers are equal in energy/bonding but have a heirachy based upon steric interactions.