Intermolecular forces

Question 1

London forces?

Answer 1

aka instantaneous dipole-induced dipole
- ALL atoms and molecules form
- weakest intermolecular forces
*any molecule/atom w electrons can form a dipole when they move near to another atom or molecule

Question 2

London forces step-by-step

Answer 2

*electrons in electron clouds alw moving really quickly/electron density fluctuates over time
1. at any particular moment - electrons in atom likely to be more to 1 side than the other
- at this mo - atom would hv a temporary/instantaneous dipole (LHS & RHS have delta +-)

2. this dipole can induce another temp dipole in opp direction on neighbouring atom as electron clouds in both mols repel e/o - pushing electron density of neighbouring mol to far end
   - (as e.g. delta+ end of mol A closer to mol B so - electron density of mol B pulled to left so - delta- on LHS & delta+ on RHS so - induced dipole) - 2 dipoles then attracted to e/o
   - leaves + and - attracted to e/o - weak attraction = London force
   - when mols move away dipole interaction destroyed

3.2nd dipole can induce another in 3rd atom
4. cos electrons constantly moving - dipoles r being created & destroyed all the time
- even tho dipoles keep changing - overall effect = atoms attracted to e/o

Question 3

London forces hold some molecules in a lattice/crystal structure? e.g. I2

Answer 3

• I atoms held tgt in pairs by strong covalent bonds to form mols of I2
• molecules then held tgt in a molecular lattice arrangement by weak London forces aka simple molecular structure

Question 4

Features of London forces?

Answer 4

• (the larger the mol, the larger the electron cloud) the more electrons there r in a mol, the greater the fluctuation in electron density (more electron drift) & larger the instantaneous & induced dipoles created
• depend on shape & size of mols - the more points of contact between mols, the greater force
• mols w greater SA - stronger London as bigger exposed electron cloud
• boil liquid/melt solids - breaking weak London forces - stronger forces - more energy - higher bp
• longer, straight chain hydrocarbons have more forces so - more energy needed to overcome - bp increases
• branched hydrocarbons can’t pack tgt as close - weakens forces between chains - lowers bp (molecular surface contact small compared to straight)

Question 5

Dipole-Dipole

Answer 5

• delta+ & delta- charges on polar molecules cause - weak electrostatic forces of attraction between molecules aka permanent dipole - permanent dipole bonds
• unlike London, these interactions involves molecules w permanent dipole so - stronger
• mols w this hv London asw
• mols that hv both generally hv higher mp & bp w similar London that can’t form dipole-dipole

Question 6

Testing polar molecules?

Answer 6

• place charged rod bear a steady stream of a polar liquid e.g. water
• shld see liquid bend towards rod as mols align to face opp charged rod
  e.g. O delta- &+ rod attracted & bends towards

Question 7

Hydrogen bonding?

Answer 7

• when H on 1 mol forms bond w lone pair on F, N, O (3 most electronegative elements)
• F, N, O very electronegative so - draw bonding electrons away from H atom
• bond so polarised & H has such high charge density (cos so small) that - H atoms form weak bonds w lone pairs of electrons on F, N, O atoms of other mols
• strongest intermolecular force
• organic mols that form H bonds often contain OH/NH groups
• when comparing: look at how many H bonds each molecule can form

Question 8

Drawing hydrogen bonding?

Answer 8

• dashed lines between lone pairs on F, N, O & H on other mol
• *lone pairs
• *partial charges

Question 9

Features of H bonds?

Answer 9

• substances that form H bonds hv high mp & bp - a lot of energy required to overcome H bonding

Question 10

Why does ice float on water?

Answer 10

• in ice: H2O mols arranged so that there’s max no. H bonds - regular lattice structure
• H bonds hold H2O molecules further apart than liquid - less dense
• as ice melts - some of h bonds broken & lattice breaks down - allowing molecules to ‘fill’ spaces
• so - ice much less dense than H2O - floats

Question 12

How do the no. hydrogens i’m a molecule limit the number of hydrogen bonds it can form?

Answer 12

• when there aren’t enough hydrogens to satisfy all the lone pairs & vice versa
  e.g. HF can only form 1 : as 1 H even tho 3 F lone pairs
• e.g. each N in NH3 has 1 lone pair & 3 Hs so - not enough lone pairs to satisfy Hs