Module 2 - [ch5,6] (atoms) Module 3 - [ch7] (periodicity)
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Why does first ionisation energy decrease going down a group [3]
atomic radius increases, increased distance between outermost electron and nucleus more electron shells more electron shielding less nuclear attraction
How are shells filled with electrons?
orbitals are filled in order of energy (lowest energy shells are filled first) electrons are assigned individually before being paired
what is an atomic orbital
a region around the nucleus holding max two electrons with opposite spin (cloud of negative charge) it’s impossible to know where exactly the electron is
ammonium ion (charge)
NH4+
what is the shape and bond angle of an NH4+ ion
tetrahedral , 109.5
Discuss the conductivity of ionic compounds [2]
when solid ions are in fixed positions and are not free to carry charge when molten or in aqueous solution ions are mobile and can carry charge
define first ionisation energy [3]
energy required to remove one electron [1] from each atom in 1 mole [1] of gaseous atoms [1] to form 1 mole of gaseous 1+ ions
why is second ionisation energy generally greater than the first
since an electron has been removed the remaining electrons experience less repulsion and are pulled closer to the nucleus experience more nuclear attraction harder to remove proton : electron ratio is greater
define isotopes [1]
atoms of an element with the same number of protons but different number of neutrons [1]
what would be evidence for shells in an atom?
jump in first ionisation energy suggests the electron is much harder to remove since is much closer to the nucleus
Explain why the first ionisation energies show a general increase across Period 2 [2]
increasing atomic number ( number of protons) more nuclear attraction similar shielding across period since same shell
Explain why the first ionisation energy of B is less than that of Be. [2]
In B, an electron is being removed from a higher energy level An s electron lost in Be and a p electron lost in B
What determines which block an element is in?
the highest energy level shell being filled
what is meant by ionic bonding [1]
electrostatic force of attraction between oppositely charged ions [1]
predict the bond angle of an F2O molecule, explain your choice [3]
2bp , 2lp bond angle 104.5 bent (non-linear) lone pairs repel more than bonded pairs
Describe and explain two anomalous properies of water which results from hydrogen bonding [4]
when solid, water is less dense than when liquid ice has an open lattice relatively high mp and bp hydrogen bonds are strong, require lots of energy to overcome
what is metallic bonding
giant metallic lattice electrostatic attraction between fixed positive ions and delocalised electrons
describe the structure and bonding shown by cl, how does this explain it’s difference in melting point compared to Mg ?
simple molecular lattice Cl has london forces between molecules london forces weaker than metallic bonds hence a lower melting point than Mg
Explain, in terms of bonding and structure, the properties of graphite [6]
graphite is maleable [1] , has a very high melting point [1] and a good conductor [1] maleable - layers can slide over one another, weak london forces between layers [1] high mp - strong covalent bonds require energy to overcome (giant covalent lattice) [1] good conductor - delocalised electrons free to carry charge [1]
shape and bond angle of an ammonia molecule
NH3 3bp 1lp 107 pyramidal lone pairs repel more than bonded pairs
diagram of hydrogen bonding?
H+ - O:~ —— H+ - O: ~
Suggest why H2S has a much lower boiling point than H2O
no hydrogen bonding, weaker IM forces, less energy required to overcome forces
why is chlorine is a gas at room temperature but carbon does not boil until well over 4500 °C. Explain this difference, in terms of bonding and structure.
to boil Cl2 weak london forces have to be broken to boil carbon covalent bonds need to be broken covalent bonds are much stronger than london forces, more energy required
what is meant by hydrogen bonding
electrostatic attraction between a hydrogen atom in one polar molecule (e.g. water) and a small electronegative atom (O,N,F) in another molecule