7.2-ionisation Energies Flashcards
Define ionisation energy
The energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions
Write an equation for the first ionisation energy of magnesium
Mg (g) ——> Mg+ (g) + e-
What are the factors that affect ionisation energy?
( affect the attraction between the nucleus and the outer electrons of an atom )
• atomic radius
• nuclear charge
• electron shielding or screening
How does atomic radius affect ionisation energy?
- the greater the distance between the nucleus and the outer electrons = the less the nuclear attraction
- the force of attraction falls off sharply with increasing distance
How does nuclear charge affect ionisation energy?
- the more protons there are in the nucleus of an atom = the greater the attraction between the nucleus and the outer electrons
How does electron shielding affect ionisation energy?
- Electrons are negatively charged and so inner - shell electrons repel outer shell electrons
- this repulsion called the shielding effect reduces the attraction between the nucleus and the outer electrons
Finish the sentence -
An element had as many ionisation energies as …
There are electrons
for eg. Helium has 2 electrons and 2 ionisation energies
He (g) —-> He+ (g) + e-
He+ ( g) ——> He2+ (g) + e-
Define second ionisation energy
The energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions
Explain what occurs in the second ionisation of helium
• the second ionisation energy of helium is greater than the first ionisation energy
• In a helium atom there are two protons attracting two electrons in the 1s sub-shell
• after the first electron is lost the single electron is pulled closer to the helium nucleus
• the nuclear attraction on the remaining electron increases and more ionisation energy will be needed to remove this second electron
Why does first ionisation energy decrease between group 2 to 3?
• there’s a decrease between group 2 to 3 because in group 3 the outermost electrons are in p orbitals
• whereas in group 2 they are in S orbital so the electrons are easier to be removed
What are the 3 predictions from successive ionisation energies?
• the number of electrons in the outer shell
• the group of the element on the periodic table
• the identity of an element
What are the two key patterns in ionisation energies?
• the general increase in first ionisation energy across each period
( H —> He , Li —> Ne , Na —-> Ar )
• the sharp decrease in first ionisation energy between the end of one period and the start of the next period
( He —-> Li , Ne —-> Na , Ar —-> K
Why does the first ionisation energy decrease between group 5 to 6 ?
• due to the group 5 electrons in p orbital are single electrons and in group 6 the outermost electrons are spin paired with some repulsion
• thus the electrons are slightly easier to remove
Does the first ionisation increase or decrease between the end of one period and the start of the next? Why?
Decrease
• there is a increase in atomic radius
• increase in electron shielding
Does the first ionisation energy increase or decrease down a group?
decrease
• shielding increases = weaker attraction
• atomic radius increases = distance between the outer electrons and nucleus increases = weaker attraction
• increase in number of protons is outweighed by increase in distance and shielding
How does nuclear charge impact I.E when going down a group ?
• as the number of protons increase there is a stronger attractive force pulling the electrons from the nucleus creating an increase in IE
How does atomic radius impact IE when going down the group ?
• distance between the nucleus as the outer electrons increase as another shell is added
• decreasing the attraction = creating a decrease in IE
How does electron shielding impact IE when going down a group ?
• an extra shell of electrons is added as you go form the group
• shielding of the nucleus increases
• decreasing the overall attraction
• this creates a decrease in IE
How does ionisation energy change as you go across a period?
Increases
How do the 3 factors of IE change as you go across a period?
• nuclear charge increases
• atomic radius increases
• electron shielding does not change
How does nuclear charge impact IE when going across a period?
• there is an increased number of protons
• the attractive force of the nucleus increases
• creating an increase in IE
How does atomic radius impact IE when going across the period ?
• distance between the nucleus and outer electrons decreases
• creating an increase in IE
How does electron shielding impact IE when going across the period?
- no effect as there is no shell added or removed
- thus no change in IE
What are the 2 anomalies to these rules?
• drop between group 2+3
• drop between group 5+6
What causes the drop between group 2 and 3?
- the outer electrons in group 3 are in the p orbital which has a higher energy than the S orbital
- thus there is additional electron shielding
- which decreases the overall IE
What causes the drop between group 5 and 6?
- group 5 elements have electrons removed from singularly filled orbitals but group 6 elements are removing electrons from orbitals filled with 2 electrons
- due to the repulsion group 6’s electrons are easier to remove
- decreasing the overall IE
How do successive ionisation energy’s change?
- successive ionisation energies increase
- there is a large jump when a new shell is used
Compare nitrogen and oxygen in regards to IE
( graph on kerboodle p.100 )
- first fall in first ionisation energy from nitrogen to oxygen marks the start of electron pairing in p orbitals of the 2p sub shell
- in nitrogen and oxygen the highest energy electrons are in a 2p sub- shell
- in oxygen the paired electrons in one of the 2p orbitals repel one another making it easier to remove an electron from an oxygen atom than a nitrogen atom
- thus the first ionisation energy of oxygen is less than the first ionisation energy of nitrogen
What does the size of ionisation enthalpy depend on ?
The atomic radius
What is electron shielding?
- the repulsion between electrons in different inner shells
- shielding reduces the net attraction force from the positive nucleus on the outer shell electrons
- the more inner shells = greater the shielding
What is the trend down a group for the 1st IE ?
Decreases as
- atomic radius increases
- shielding increases
- nuclear attraction to outer electrons decreases
What is the trend across a period for the 1st IE?
Increases as ‘
- atomic radius decreases
- shielding is Similar
- nuclear charge increases
- nuclear attraction increases
What is successive ionisation energies?
- the ionisation energy increasing when removing first then second etc electrons
- each ionisation energy is greater than the one before it because an electron is being removed from an already positive ion
- more protons attracting fewer electrons
Why are successive ionisation energies important ?
They provide important evidence for the different electron energy levels in an atom
Explain why successive ionisation energies always increase?
- as each electron is removed the outer shell of drawn closer to the nucleus
- nuclear attraction is greater and more energy is needed to remove the next electron
( 2 marks in an exam )
