Chemistry - Unit 1 - Topic 1 - Electron configuration and ionisation energiesa=

what are electron shells called?

energy levels

what is each energy level given?

a number called the principle quantum number - indicates the shell number that electrons occupy (n)

how many electrons does each electron shell/energy level hold?

2n²

what is each energy level divided into?

sub shells

what is a sub shell?

an orbital or a combination of orbitals

what is an orbital?

an orbital is a 3D region in space that can hold 2 electrons - a space where electrons are likely to be found 

what are the 4 types of orbital?

s,p,d,f

describe an s orbital

• an s orbital has a spherical shape - up to 2 electrons can be found in this space

• each shell contains 1 s-orbital

describe a P-orbital

• 3d dumbell shape

• 3d shape is a result of each P orbital having 3 further orbitals, Px, Py, Pz 

• Px is side to side; Py top to bottom and Pz front to back

• each orbital can hold 2 electrons

• overall the P orbital can hold 6 electrons

describe a D orbital

• have 5 orbital planes

• can hold up to 10 electrons

describe an F-orbital

• there are 7 orbital planes

• can hold up to 14 electrons

how are electrons arranged in 1 orbital?

• there are 2 electrons per orbital

• these spin in opposite directions in order to minimise repulsions

• drawn as boxes with one half-arrow up and one down

in what order are electron subshells/orbitals filled?

• in ascending order of energy level and s→p→d

• 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10….

how are subshells filled?

• each orbital of a type (s OR p Or d, in that order) is filled with 1 electron, then those electrons are paired

• i.e. there is one electron in each of the Px, Py, and Pz orbitals before there are two in the Px

what is the exception to the sub-shell filling rule of spd?

• 4s2 fills before 3d10 as 3d is at a slightly higher energy level

• this is the same for 5s… and so on

what is another way of writing electron structure?

• writing the previous noble gas (closest one before the element) in square brackets

• continuing the electron structure from there 

what are the two exceptions to the subshell filling rules?

• chromium 

• copper

what is the electron structure of chromium?

• 1s2 2s2 2p6 3s2 3p6 3d5 4s1

• this is how you would expect, except that the 4s orbital only has one electron (expect 2) and the 3d orbital has 5 electrons (expect 4) 

why is the electron structure of chromium unusual?

• 5 unpaired electrons in a D orbital and a half full S orbital is more stable than a full S orbital and an almost half full D orbital 

• this is because the electrons are symmetrical around the nucleus

what is the electron structure of copper?

• 1s2 2s2 2p6 3s2 3p6 4s1 3d10

• expect 4s2 3d9

why is the electron structure of copper unusual?

• 10 paired electrons in the outer d orbital is more stable than 1 pair of electrons in the s orbital and 4 pairs and 1 unpaired electron in the d orbital

• this is because the electrons are symmetrical around the nucleus 

in what order are electrons lost when positive ions are formed?

from 4s2 and then 3d10 (i.e. in terms of quantum number not energy level)

define the term ‘first ionisation energy’ 

the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions 

give the equation for the first ionisation energy

X (g) → X⁺ (g) + e-

give the equation for the second ionisation energy

X ⁺ (g)→ X²⁺ (g) + e-

give the equation for the third ionisation energy

X ²⁺ (g)→ X³⁺ (g) + e-

what are ionisation energies measured in?

• kJ/mol

• all values will be positive as they are endothermic

what is the trend in ionisation energy for successive ionisation energies (for 1 type of atom/ion)?

the ionisation energy increases

why does the ionisation energy for successive ionisation energies increase?

• the ratio of protons to electrons increases meaning there is greater effective nuclear charge

• as each electron is removed it becomes more positively charged

what does a large jump in ionisation energy from 1 successive IE to the next indicate?

• the electron is being removed from one energy level lower - is a lot closer to the nucleus so feels greater forces of attraction and more energy is required to remove it

• this shows which group it is in 

• i.e. if the large jump is from IE 3 to IE 4, the ion is in group 3 

• if there is no jump in a table of 8 values, it must be in group 0 

• evidence for energy levels

what are the 3 main factors affecting ionisation energy?

• atomic radius

• nuclear charge

• electron shielding (most significant)

describe the factor of atomic radius

• the greater the atomic radius, the smaller the nuclear attraction experienced by the outer electrons, meaning less energy is required to remove the outer electron from the atom

• atomic radius increases down the group and increases across a period

why does atomic radius decrease across a period?

• each successive element has a greater nuclear charge

• this means the outer shell of electrons feels greater attraction to the nucleus

• this makes the atom more compact 

describe the factor of nuclear charge

• the greater the nuclear charge, the greater the attractive force on the outer electrons

• nuclear charge increases both across the period and down the group

describe the factor of electron shielding

• inner shells of electrons repel the outer-shell electrons and so make them easier to lose/less energy is required to lose them

• the more inner shells (and so larger atomic radius) an atom as, the greater the shielding and the smaller the nuclear attraction of outer electrons 

describe and explain the trend of first ionisation energies within a period

• within a period, overall the ionisation energies across it increase

• however, there can be decreases in ionisation energy mainly going from g2-3 and g5-6

• this is because atomic number across the period increases, increasing the nuclear charge and causing greater attraction between the nucleus and outer electrons, meaning more energy is required to remove the outer shells of electrons

• as well, due to increasing nuclear charge, atomic radius decreases across a period, meaning the outer electrons are closer to the nucleus and experience greater forces of attraction…

describe and explain the trend of first ionisation energies within a group

as you go down the group, the first ionisation energies decrease

what is the trend of ionisation energies across periods?

• there is a general increase across a period before the value drops dramatically for the start of another period

• this is because atomic radius increases, and each new period is a new energy level 

where are there decreases of ionisation energy within a period?

• from group 2-3 and group 5-6

• magnesium → aluminium 

• phosphorous → sulphur

what factors cause the drops in ionisation energy within a period? 

• going from no paired electrons to a pair of electrons in an orbital (g6)

• entering a new orbital (g3)

explain why magnesium has a greater first ionisation energy than aluminium (g3 reason)

• magnesium’s outer electron is in the 3s2 sub shell

• aluminium’s outer electron is in the 3p1 sub shell

• this means the electron in aluminium experiences greater electron shielding and is further away from the nucleus so feels lesser forces of attraction towards it and requires less energy to lose

explain why phosphorus has a greater first ionisation energy than sulfur

• phosphorous has no paired electrons in the outer 3p orbital

• sulfur has one pair of electrons in the outer 3p orbital 

• these electrons repel slightly making it easier to lose one of them - requires less energy 

• this repulsion is not present when there are no paired electrons