May 25, 2017
Just some ionisation energy graphs I drew up
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Just some ionisation energy graphs I drew up
Good evening, all! I’ve only been able to get the chance to post at 12:30am lmao, that’s how busy my day’s been XD
So today I have: - Done three hours of work in Graphics at school after the home-time bell - Had my Chemistry tutor over and did an hour of work with her - Done Chemistry homework on equilibrium, covalent networks and ionisation energies!
Off to bed I go!
Chemistry- Ionisation Energy Notes
What is IE?
Ionisation energy is the amount of energy needed to remove one electron from an atom or ion.
How is IE represented in a formula?
IE can be represented with this formula, e.g. Fe (g) —> Fe+ (g) + e
What does it mean when a substances has a lower first IE than another?
When a substance has a lower first IE than another, it means that less energy is needed to remove an electron thus the forces of attraction are weaker. There is less energy and thus it is more reactive.
Why does successive ionisation energy increase?
When the first electron is removed, a positively charged ion is formed, the ion becomes a cation
It is harder to remove electrons from a cation than a neutral ion.
Sharp increases in successive ionisation energies is due to the removal of an electron from the inner principal quantum shell compared to the outer shell.
How do I predict electron configuration from successive ionisation energy?
Find the largest difference in ionisation energy.
If the largest difference is between the 6th and 7th ionisation energies, the 7th electron must have been removed from the inner electron shell.
Therefore, there are 6 electrons in the valence shell
The valence electron configuration is s2 p4 and the element is in group IV.
What are the factors affecting the value of ionisation energy?
Size of effective nuclear charge and the atomic radius.
What is effective nucleus charge?
Effective nuclear charge is the nuclear charge minus the screening effect.
Nuclear charge increases as the group increases.
Group I has a smaller nuclear charge than group III.
The greater the nuclear charge, the greater IE.
Screening effect is the inner shell of electrons that can shield valence electrons from the full forces of attraction from the nucleus.
With greater screening effect, the effective nuclear charge decreases, resulting in weaker force of attraction of valence electron by enclose.
Less amount of energy is required to remove the valence electrons and hence ionisation energy decreases.
What is atomic radius?
Atomic radius is the size of the atom, meaning how many electron shells it has. The more electron shells there are, the greater the size of the atom and the lower the IE.
Comparing IE Across a Period
When comparing IE across a period, it typically increases
Electrons are being added to the same outermost principal quantum shell thus shielding effect remains approximately constant
Nuclear charge increases as proton number increases
Effective nuclear charge increases
The valence electrons are more strongly attached to the nucleus
Electrostatic attraction between the nucleus and valence electrons increases
More energy is needed to remove these electrons
IE increases
What are the exceptions in this trend?
The 1st IE of MG is greater than the 1st IE of Aluminium
In Al, the 3p electron is further away from the nucleus than the 3s electron.
Thus the electrostatic attraction between that electron is weaker
Less energy is needed for ionisation
The 1st IE of P is larger than the 1st IE of S
Less energy is needed to remove the paired 3p electron in S as it experiences inter-electron repulsion arising from 2 electrons occupying the same orbital.
Comparing IE down a group
When comparing IE down a group, IE typically decreases.
This is because electrons down a group have more electron shells.
The distance between the electron and the nucleus is further from the nucleus
So they are less strongly attracted to the nucleus
Less energy is needed to remove the electron.
Increase in screening effect outweighs the increase in nuclear charge
Hence effective nuclear charge decreases
Clever Periodicity Table
My friend had her apple macbook out in our last chemistry lesson- because she's really cool like that and she found this really helpful table which I've been trying to make myself for ages :)
I'm not very good at the whole why bit and that's what they ask you so this is really helpful :D
I tried pasting it but the table disintegrated into words :'( But the link is here :)
Sulfur takes a bus trip.
Okay so basically the general ionisation energy trend increases as you go across a period in the periodic table.
IONISATION ENERGY being the energy required to remove one mole of electrons from one mole of gaseous atoms. (In other words the energy needed to steal one electron from the outer shell)
The ionisation energy increases so an electron becomes harder to steal. This is because of the INCREASING CHARGE in the nucleus. (You're better off stealing sodium. Note the alliteration there...)
But..
Because in chemistry there is always a BUT or an annoying exception.
And this one is Sulfur. (This is where the bus bit comes in).
A favourite exam question they like to put in is why does Sulfur (Group 6) have a lower ionisation energy than Phosphorus (Group 5)?
(In other words why does it not follow the general increasing trend)
To answer this question, work out the electronic configuration. Remember the layout, or if you don't look here for a way of remembering it.
Sulphur has 16 electrons, so fill them in using the box method.
(Ideally not on paint though.. My lines look a bit shaken up)
So Sulphur would be 1s22s22p63s23p4
Then work out Phosphorus which has 15 electrons.
So the electronic configuration of phosphorus would be 1s22s22p63s23p3
The important thing here is that sulfur has two paired electrons in the same orbital. They do not like being next to each other. They repel each other.
A theory my chemistry teacher uses is to think of it like seats on a buses. You don't like sitting next to a stranger who you don't know, or someone who smells really bad on a bus, you'd rather be on a twin seat by yourself.
So it is EASIER to remove or "steal" an electron from the 3p subshell for the first ionisation energy than to remove an electron on his own.
So Sulphur has a lower ionisation energy than phosphorus.
:) If that made any sense whatsover. It does in my head.
(Note how everyone is sat on their own in this picture from google images. This just proves the bus sitting theory)
Patterns in First Ionisation Energies
Across a Period
e.g. Li --> Ne
The First I.E. tend to gradually increase going across the period. This is due to the fact that all the elements in the period all have the same complete quantum shells. Therefore, they have similar electron shielding.
However, the nuclear charges of the elements increase as each time one more proton is added to the nucleus. Therefore, there is an increase in the electrostatic forces of attraction and ionisation energy.
Down a Group
e.g. Li --> Na
There is a gradual decrease in First I.E. on going down a group. This is due to the increase in electron shielding and distance in atomic radius (the distance between outershell electron and the nucleus) outweighing the fact that there is an increase in nuclear charge.
Ionization Energies
First Ionisation Energy
Definition: The amount of energy required to remove one outer shell electron from each one mole of a gaseous atom to form one mole of a cation in a gas state forming a +1 charge.
Unit: kJ/mol
Successive Ionisation Energy
Definition: The amount of energy required to remove one outershell electron from a gaseous cation.
N.B. The electrons are removed from the outermost to the innermost. Think peeling an onion: out to in.
The successive ionisation energy increases as the electrons being removed become closer to the nucleus. Therefore, the electrostatic forces of attraction increases and so more energy is needed to overcome them in order to remove the electron.
Closer = Stronger electrostatic forces of attraction = More energy to overcome = Increased I.E.
When a new shell of electrons is started to be removed, the ionisation energy will dramatically increase due to the reduced electron shielding.
Electron shielding is when the electrons in the shells between the outermost shell and the nucleus cancel some of the positive charges from the protons in the nucleus.
Therefore, if there are less shells between the outershell and the nucleus, there are stronger forces of attraction between the nucleus and the outershell electrons.
New shell= Less electron shielding = Less electrons to cancel positive charge= Stronger forces of attraction = More energy to overcome= Increased I.E.
When an atomic number increases, so will the nuclear charge of the atom. Therefore, there are stronger forces of attraction between the increased number of protons & electrons and more energy is needed to overcome it. Therefore the I.E. also increases.
Increased Nuclear Charge= More protons= Stronger forces of attraction = More energy needed to over come = Increased I.e
Chemistry revision: Ionisation Energy
First ionisation energy: The energy required to remove 1 mole of electrons from 1 mole of atoms in their gaseous state
E (g) ----> E+ (g) + e-
Second ionisation energy: The energy required to remove 1 mole of electrons from 1 mole of positively charged ions in their gaseous state
E+ (g) ----> E2+ (g) + e-
Across a period - ionisation energy increases
Down a group - ionisation energy decreases
If ionisation energy increases, nuclear attraction increases also as there are more protons which makes it harder to remove electrons. The atoms will also want to gain electrons rather than lose as they are wanting to be like the nearest noble gas.
If ionisation energy decreases, screening increases as the electrons are further away from the nucleus. Nuclear attraction decreases also, making it easier to remove an electron.
The second ionisation energy is always larger than the first as it is trying to remove an electron from a positively charged ion.