Chapter 3~electrons In Atoms Flashcards
Definitions, key ideas and formulae
Definition of energy ⛮ levels:
°the regions at various distances from the nucleus in which electrons have a particular amount of energy ⛮. Electrons further from the nucleus have more energy ⛮
Definition of principal quantum shells, n :
°Regions at various distances from the nucleus that may contain up to a certain number of electrons. The first quantum shell contains up to 2 electrons, the second up to 8 and the third up to 18
How many electrons can shell 1 hold?
2
How many electrons can shell 2 hold?
8
How many electrons can shell 3 hold?
18
How many electrons can shell 4 hold?
32
Definition of the first ionisation energy ⛮ :
°The 1st ionisation of an element is the energy ⛮ required to remove one electron from each atom ⚛ in one mole of atoms of the element in the gaseous state to form one mole of gaseous 1+ ions.
Definition of ionisation energy ⛮ :
°the energy ⛮ required to remove 1 mole of electrons from 1 mole of atoms of an element in the gaseous state to form 1 mole of gaseous ions.
Definition of successive ionisation energy ⛮ :
°the energy ⛮ required to remove the first, then the second, then the third electrons and so on from a gaseous atom ⚛ or ion, producing an ion with one more positive charge each time. Measured in KJ per mole of ions produced
Why do the successive ionisation energies increase? (3 statements)
¬bec the positive charge on the ion gets greater as each e- is removed
¬as the e- is removed there is a greater attractive force between the positively charged p+’s in the nucleus and the remaining negatively charged e-‘ s
¬therefore more energy ⛮ is required to overcome these attractive forces
Why is there a huge difference between some successive ionisation energies? (2 statements)
¬these large changes indicate that the next e- that was being removed was being removed from a principal quantum shell closer to the nucleus.
¬ therefore a lot more energy ⛮ is required to overcome the increased forces of attraction between the + nucleus and the - e-‘s
Name the 4 main factors that influence ionisation energies :
- Size of the nuclear charge
- distance of outer e-‘s from the nucleus
- Shielding effect of inner e-‘ s
- spin - pair repulsion
In general, as proton number increases, ionisation energy … (increases or decreases)?
Increases
The further the outer e- shell is from the nucleus the higher/ lower the ionisation energy ?
Lower
As the number of full e- shells 🐚 between the outer e-‘s and the nucleus increases.
Is the ionisation energy ⛮ lower or higher?
Lower
Why does increase in proton number increase the ionisation energy? (in general)
Bec as the positive nuclear charge increase, there is also a greater attractive force between the nucleus and the e-‘s.
Therefore, more energy is required to overcome these attractive forces if an e- is to be removed
Why is it that the further the outer e- shell is from the nucleus, the lower the ionisation energy ⛮?
The force of attraction between the positive and the negative charges decreases rapidly as the distance between them increases.
Therefore, e-‘s in shells further from nucleus are LESS attracted to it than ones closer to the nucleus
Definition of shielding :
°the ability of inner shells of e-‘s to reduce the effective nuclear charge on e-‘ s in the outer shell
Why is it that the greater the shielding the lower the ionisation energy? (3)
e-‘s in full inner shells repel e-‘ s in outer shells.
Full inner shells of e-‘s prevent the full nuclear charge being felt by outer e-‘ s.
The greater the shielding of outer e-‘s by inner e-‘ s the lower the attractive forces between the nucleus and the outer e-‘s
When the e- being removed has a very very low ionisation, what could be the 2 likely causes of that?
- Large distance from the nucleus
2. And the e- is well shielded by inner shells
If there is a large jump in value of ionisation energy, what is the likely cause?
The second e- in the jump is within a shell closer to the nucleus than the first e- in the jump
When there is a gradual increase in the successive ionisation, the 3 likely causes are?
- Shielding is constant bec e-‘s are from the same shell
- Distance from nucleus is similar
- Proton number is increasing and therefore the nuclear attraction to the e-‘ s
When an ionisation energy ⛮ is extremely high, what does that most likely mean?
The e- being removed is very close to the nucleus
We can use successive ionisation energies to do 2 things :
- predict/confirm the simple e- configuration of elements
2. Confirm number of e-‘s in the outer shell of an element and hence the group to which the element belongs
Definition of subshells :
°Regions within the principal quantum shells where e-‘s have more or less energy ⛮ depending on their distance from the nucleus. Subshells are given the letters s, p, d and f
List the sub shells in order of increasing energy :
From 1s to 4d
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d
Definition of atomic orbitals :
°an atomic orbitals is a region of space around the nucleus of an atom that can be occupied by one or two e-‘s (maximum) .
E. G. p, s and d
How many atomic orbitals are in each subshell for s, p and d?
(think of how many blocks each has in electron configuration)
s-one orbital
p- 3 orbitals
d-5 orbitals
What shape is the s - orbital?
Spherical
What are the shape of the px, py and pz orbitals?
hourglass shaped with two ‘lobes’
Why whenever possible will e-‘s occupy separate orbitals in the same shell?
To minimise repulsion
Definition of spin-pair repulsion :
°electrons repel each other as they has the same charge. Electrons arrange themselves so that they singly occupy different orbitals in the same sublevel. After that they pair up with their spins opposed to each other
Why is there a general increase in ionisation across a period? (1statements and 3 reasons )
The force of attraction between the positive nucleus and the negative e-‘s increase across a period bec:
- The nuclear charge increases
- The distance between the nucleus and the outer e- remains reasonably constant
- The shielding by inner sells remains reasonably constant
There is a rapid decrease in ionisation energy ⛮ between the last period and the first element in the next period. Why?(1 statement and 3 reasons )
The force of attraction between the positive nucleus and the outer negative e-‘s decrease because :
- distance between nucleus and outer e- increases
- The shielding by inner shells increases
- These 2 factors outweigh the increased nuclear charge
When elements have the configuration :
E. G. Beryllium : 1s^2 .2s^2
Boron : 1s^2 .2s^2. 2p^1
So when one e- is in a further atomic orbital. Why does boron have a lower ionisation energy than beryllium (even though it has an extra proton)?
(on statement and 3 reasons)
There is less attraction between the (boron’s) outer most negative e- and positive nucleus because :
- The p - orbital is slightly further from the nucleus than the s-orbital. So the distance between the nucleus and the outer e- increases slightly
- The shielding by inner shells increases slightly.
- these 2 factors outweigh the increased nuclear charge
If you compare elements e.g.
Nitrogen : 1s^2. 2s ^2. 2p^3
And
Oxygen :1s^2. 2s ^2. 2p^4
Which one will have the lower ionisation energy and why? (3 statements)
Oxygen because :
The e- being removed from Nitrogen is from an unpaired orbital whilst the e- being removed from oxygen is from an orbital that contains a pair of e-‘s
The extra repulsion between pairs of e-‘ s in this orbital results in less energy being required to remove an e-
Ionisation energy is lower bec of SPIN-PAIR REPULSION
Why does the ionisation decrease as you go down the group? (1 statement, 3reasons)
As you go down the group, the outer e- removed is from the same type of orbital but from a successively higher principal quantum level (shell e.g. n=1, n=2, n=3 etc). There is less attraction between the outer e- and the nucleus because :
- distance between the nucleus & the outer e- increases
- The shielding by complete inner shells increases
- These 2 factors outweigh the increased nuclear charge
