module three: periodic table and energy Flashcards
blocks of periodic table
s d p f
factors that affect ionisation energy
atomic radius
nuclear charge
electron sheilding
trends in ionisation energy
increases across a period
decreases down a group
how does atomic radius affect ionisation energy
e- in shells further away from nucleus are less attracted to the nucleus = easier to lose e-
the further the outer e- from nucleus = the lower the ionisation energy
trends in atomic radius across a period
decreases across a period due to nuclear charge
e.g. Na and Ar
both elements have same number of shells, outer shells: Na = 1 e- Ar = 8 e-
Ar nucleus = more protons = more positive = stronger force of attraction between outer shell and nucleus
how does nuclear charge affect ionisation energy
the more protons in the nucleus, the stronger the attraction between nucleus and electrons = more energy req. to overcome forces when removing e-
trends across Period 3
atomic radius decreases
ionisation energy generally increases
electronegativity increases
trends down group 2
atomic radius increases - more shells
shielding increases
reactivity increases - outer electron lost more easily
ionisation energy decreases
describe + explain reactivity down group 2
reactivity increases down group - elements react by losing electrons to form 2+ ions, req. 2 I.Es
I.Es decreases down group
le chatelier’s principle
the equilibrium shifts to minimise the effect of any change
effect of decreasing the pressure on the rate of a reaction
decreases the rate of reaction
decreased concentration of molecules
fewer molecules per unit volume
less frequent successful collisions
trend in atomic radius across period 3
as the number of protons increases, the force of attraction in the nucleus also increases
the e- are closer to nucleus
all e- are in same shell but as atomic charge increases, e- in their shell are pulled closer to nucleus = distance decreases
outer e- are in same shell, more protons, same amount of shielding, stronger attraction
periodicity
pattern in properties across a row which is repeated in each row
trends in first ionisation energy across period 3
p orbital is further away than s orbital
period increases number of protons in nucleus = greater attraction
properties of metallic substances
due to delocalised ‘sea’ of electrons:
high melting and boiling point: lots of energy req. to overcome strong electrostatic forces of attraction between positive ions and the ‘sea’ of delocalised electrons
solubility: metals do not dissolve. some interaction between polar solvents and charges in the metallic lattice lead to reactions, rather than dissolving e.g. sodium and water
electrical conductivity: conduct electricity in both solid and liquid states due to the delocalised electrons which are free to move & carry charge around the structure
structure of diamond
giant covalent lattice of C atoms
each carbon atom is covalently bonded to four others in a tetrahedral arrangementwith a bond angle of 109.5
result is a giant lattice structure with strong bonds in all directions
diamond is the hardest substance known = used in drills and glass-cutting tools
structure of graphite
giant covalent lattice of C atoms
each carbon atom is bonded to three others in a layered structure
bond angle of 120
spare electrons are delocalised and occupy the space between the layers
all atoms in the same layer are held together by strong covalent bonds
layers are held together by weak intermolecular forces = allow the layers to slide over each other
structure of graphene
contain an infinite lattice of covalently bonded atoms in two dimensions only to form layers
made of a single layer of carbon atoms that are bonded together in a repeating pattern of hexagons
considered two dimensional
structure of silicon (IV) oxide
networks of atoms bonded by strong covalent bonds
known as silicon dioxide
giant covalent lattice
structure made of tetrahedral units all bonded by strong covalent bonds
each silicon is shared by four oxygens and each oxygen is shared by two silicons
an empirical formula of SiO2
explanaton of metallic bonding
metal atoms are tightly packed together in lattice structures
in lattice structures, the e- in outer shells are free to move throughout the structure = delocalised electrons - not bound to their atom
when the e- are delocalised, the metal atoms become positively charged ions
positive charges repel each other and keep the lattice in place
very strong forces between positive metal centres & ‘sea’ of delocalised electrons
explanaton of the variaton in melting points across periods 2 and 3
melting point increases from group 1 to group 4
groups 1 to 3 have metallic bonding = increased forces of attraction between more electrons in the outer shell that are released to the sea of electrons and a smaller positive ion = increases in strength
group 4 has a giant covalent structure = req. lots of energy to overcome
sharp decrease in melting point from group 4 to group 5
group 5 to group 0 have simple molecular structures with weak London forces between molecules = little energy req. to overcome
