7.2 Ionisation energies Flashcards
What is ionisation energy
Measures how easily an atom loses electrons to form positive ions
What is the ‘first 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
How are electrons held in their shells
By attraction from the nucleus
The first electron lost will be in the ______ energy level and will experience the ______ attraction from the nucleus,
Highest
Least
What factors affect the attraction between the nucleus and the outer electrons of an atom (and also the ionisation energy)
Atomic radius
Nuclear charge
Electron shielding
Atomic radius and affecting ionisation energy:
The greater the distance between the nucleus and outer electrons, the less the nuclear attraction
Force of attraction falls off sharply with increasing distance so atomic radius has a large effect
Nuclear charge and affecting ionisation energy:
The more protons there are in the nucleus, the greater the attraction between the nucleus and the outer electrons
Electron shielding and affecting ionisation energy:
Electrons are negatively charged and so inner-shell electrons repel outer-shell electrons. This repulsion (the SHIELDING EFFECT) reduces the attraction between the nucleus and the outer electrons
An element has as many ionisation energies as there are…
…electrons
Eg. He(g) -> He+(g) + e- first ionisation energy
He+(g) -> He2+(g) + e- second ionisation energy
What is 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
Successive ionisation energies allow predictions to be made about:
The number of electrons in the outer shell
The group of the element in the periodic table
The identity of an element
What does a sudden large increase between two ionisation energies show Eg. between the third and fourth
The fourth electron is being removed from an inner shell
Therefore there are three electrons jin the outer shell and the element must be in group 13 (3)
What do periodic trends in first ionisation energies provide
Important evidence for the existence of shells and sub-shells
What does a sudden jump in energy on an ionisation energy vs. ionisation number graph show
A change from one shell to another
What are the two key patterns in the first ionisation energies for the first 20 elements in the periodic table:
A general increase in first ionisation energy across each period
A sharp decrease in first ionisation energy between the end of one period and the start of the next period
What are the reasons for the decrease in first ionisation energy
The increased atomic radius and shielding going down a group
Trend in first ionisation energy down a group
Atomic radius increases
More inner shells so shielding increases
Nuclear attraction on outer electrons decreases
First ionisation energy decreases
Trend in first ionisation energy across a period
Nuclear charge increases (most important factor) Same she’ll: similar shielding Nuclear attraction increases Atomic radius decreases First ionisation energy increases
First ionisation energy trends and sub shells
Although first ionisation energy shows a general increase across both period 2 and 3, it falls in two places in each period
Drops occur at same positions in each period
Why are there two rides and two falls in first ionisation energy across period 2
A rise from lithium to beryllium
-filling of 2s sub-shell
A fall to boron followed by a rise to carbon and nitrogen
-adding one electron to each 2p orbital
A fall to oxygen followed by a rise to fluorine and neon
-pairing of 2p electrons
What is the most important factor for the general increase in first ionisation energy
The increased nuclear charge
Explaining the fall in first ionisation energy from beryllium to boron
The 2p sub-shell in boron has a higher energy than the 2s sub-shell in beryllium
The 2p electron is therefore easier to remove than one of the 2s electrons
The first ionisation energy of boron is less than that of beryllium
Explaining the fall in first ionisation energy from nitrogen to oxygen
Start of electron pairing in the p-orbitals of the 2p sub-shell
Highest energy electrons are in the 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
Therefore the first ionisation energy of oxygen is less than the first ionisation energy of nitrogen
