1.1 atomic structure- electron config Flashcards
what is an orbital
region of space in which electrons move
what are the s and p orbital shapes
s=spherical
p=dumbbell
how many electrons can each orbital hold
s=2
p=6
d=10
what is chromiums electron structure
1s1 2s2 2p6 3s2 3p6 3d5 4s1
-by having one electron in the 4s subshell, it can have a half full 3d subshell which is more stable when half/completely full
what is coppers electron structure
1s1 2s2 2p6 3s2 3p6 3d10 4s1
-by having one electron in the 4s subshell, it can have a full 3d subshell which is more stable when half/completely full
why is aluminium a p-block element
its highest energy electron is in a p-orbital
what is the first ionisation energy
energy needed to remove 1 mole of electrons from 1 mole of atoms in their gaseous state, to form 1 mole of +1 ions in their gaseous state
X(g)–>X+(g) + e-
what is the second ionisation energy
energy needed to remove 1 mole of electrons from 1 mole of +1 ions in their gaseous state, to form 1 mole of +2 ions in their gaseous state
X+(g)–>X2+(g) + e-
3 factors affecting the size of ionisiation energies
-as atomic radius increases, force of attraction between valence electrons+nucleus decreases
-greater number of protons=greater force of attraction between valence electrons+nucleus
-valence electrons are repelled by inner shell electrons. This shielding effect decreases attraction between valence electrons+nucleus
why do successive ionisation energies increase in each shell (2)
electrons are being removed from an increasingly positive ion
so they’re held more strongly by the nucleus
what are the big jumps in ionisation energy graphs
happens when there’s a new shell broken into- an electron is being removed from a shell closer to the nucleus
how to work out the group of the element from an ionisation energy graph
count how many electrons are removed before the first big jump
how can ionisation energy graphs predict the electron configuration of an element
from right to left, count how many electrons there are before each big jump to find the number of electrons in each shell
what happens to ionisation energy across a period
-increases
reasons:
-more protons in nucleus
-atomic radius decreases
-same amount of shielding
so more attraction between nucleus and valence electron
what happens to ionisation energy down a group and when moving into a new period
-decreases
reasons:
-atomic radius increases so valence electron is further from nucleus
-more shielding as there are more electron shells
so less attraction between nucleus and valence electron
why are there dips in ionisation energies across a period eg. Mg->Al
-in Al, electron is removed from p orbital
-which is higher in energy/ shielded by 3s
-so its lost more easily
why are there dips in ionisation energies across a period eg. P->S
-in S, there are paired electrons in p orbital
-paired electrons repel
explain the pattern in the first ionisation energies from lithium to neon (period 2) (6)
-generally, first IE increases across a period
-as there are more protons in the nucleus and electrons are in the same energy levels so there is the same amount of shielding
-leads to more attraction between valence electrons and nucleus
-drop from beryllium to boron (B has lower IE)
-boron’s highest energy electron is in the p-orbital whereas beryllium’s highest energy is in the s-orbital
-boron’s highest energy electron is shielded by the 2s subshell, so it is more easily removed
-2nd drop from nitrogen to oxygen (O has lower IE)
-oxygen has a pair of electrons in its 2p subshell, which experiences repulsion so it is easier to remove
why is the ionisation energy of every element endothermic
heat energy needed to overcome attraction between electron and nucleus
which elements in period 2 and 3 break the trend
Boron
Oxygen
Aluminium
Sulfure
BOAlS
