3.1.1 Periodicity
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how was the periodic table arranged before
• elements arranged in order of atomic mass • lines up elements in groups with similiar properties
how is the periodic table arranged now
• elements arranged in order of increasing atomic number (proton) • elements arranged in vertical columns called groups • elements arranged in horizontal rows called periods
what do elements in the same group have
• atoms with the same number of outer-shell electrons and similiar properties
what do elements in the same period have
• the same number of electron shells (energy levels) or highest energy electron shells
what are the 4 blocks in the periodic table
• s, p, d, f
what is the periodic trend in electron configurations across period 2
• 2s sub-shell fills with 2 electrons • 2p sub-shells fill with 6 electrons
what is the periodic trend in electron configurations across Period 3
• 3s sub-shell fills with 3 electrons • 3p sub-shells fill with 6 electrons
what does ionisation energy measure
• measures how easily an atom loses electrons to form positive ions
define first ionisation energy
• the energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions
define second ionisation energy
• energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions
what are the 3 factors that affect ionisation energy
• atomic radius • nuclear charge • electron shielding
why do successive ionisation energies increase
• nuclear attraction on remaining electrons increases • the electrons are closer to nucleus + so harder to remove them > therefore inc ionisation energy
what is the trend in first ionisation energy down a group
• decreases: > atomic radius inc > more inner shells so shielding inc > nuclear attraction to outer electrons dec > first ionisation energy dec > although nuclear charge inc, the effects are outweighed by inc atomic radius
what is the trend in first ionisation energy across a period
• increases: > nuclear charge inc (more protons in nucleus) > same shell - similiar shielding > nuclear attraction inc > atomic radius dec > first ionisation energy inc
why is there a decrease in ionisation energy from Br to B in period 2
• this marks the start of filling the 2p sub-shell • 2p sub-shell in B has higher energy than 2s sub-shell in Br > therefore it’s easier to remove the the 2p electron in B than one of the 2s electrons in Br
why is there a decrease in ionisation energy from N to O in period 2
• in N and O, the highest energy electrons are in 2p sub-shell • in O, the paired electrons in one of the 2p orbitals repel one another > making it easier to remove an electron from O, atom than N atom > so first ionisation energy of O is less than that of N
where are the two changes in first ionisation energy found going across the periods
• from G2-G13 • from G15-G16
what is periodicity
• repeating trend in properties of elements across each period
what do simple molecular structures consist of
• few atoms covalently bonded + held by weak IMF (London forces)
what do giant covalent structures consist of
• billions of atoms held by strong covalent bonds to form a regular lattice
what are some properties of simple molecules
• low mp/bp due to weak imf that require little energy to break • don’t conduct electricity because no delocalised electrons or mobile charge carriers
when melting/boiling simple molecules, what are you breaking
• the weak imf
describe the solubility of giant covalent lattices
• insoluble in almost all solvents • covalent bonds holding atoms together in lattice are too strong to be broken down by interaction with solvents
when carbon forms 4 covalent bonds to other carbon atoms, what is the structure known as
• diamond