Ionisation Energies

Ionisation Energies

• First Ionisation Energy = The energy required for one mole of gaseous to from one mole of gaseous ions with a single positive charge
• This is represented by the equation: M(g) --> M^+(g) + e^-
• Note that the equation for 1st ionisation energy always follows the same pattern, even if the atom does not normally form a 1^+ ion or is not gaseous (although in mass spectrometry it always will be gaseous).
• Second Ionisation Energy = The energy required for one mole of gaseous atoms with a single positive charge to form one mole of gaseous ions with a double positive charge.
• This is represented by the equation: M^+(g)) --> M^2+(g) + e^-
• REMEMBER THESE DEFINITIONS!

Factors that affect ionisation energy

• There are three main factors:
• The attraction of the nucleus (the more protons n the nucleus the greater the attraction)
• The distance of the electrons from the nucleus (the bigger the atom the further from the nucleus the outer shell electrons are and the weaker the attraction to the nucleus)
• Shielding of the attraction of the nucleus (an outer-shell electron is repelled by electrons in complete inner shells, weakening the attraction of the nucleus)

Successive ionisation energies

• The patterns in successive ionisation energies for an element give us important information about the electronic structure for that element.
• Why are successive ionisation energies always larger?
• The second ionisation energy of an element is always bigger than the first ionisation energy of an element. When the first electron is removed a positive ion is formed. The ion increases the attraction on the remaining electrons so the energy required to remove the next electron is larger.
• How are ionisation energies linked to electron structure? (right)
• Notice the big jump between 4 and 5: The fifth electron is in an inner shell closer to the nucleus and therefore attracted much more strongly by the nucleus than the fourth electron. It also does not have any shielding by inner complete shells of electrons. From this (ignoring the title) it is possible to work out that these ionisation energies belong to carbon.
• 

Folie 4

• Example question:
• What group must this element be in? (ionisation energy is given in kJ/mol^-1)
• 1 = 590
• 2 = 1150
• 3 = 4940
• 4 = 6480
• 5 = 8120
• Example answer:
• Here there is a big jump between the 2nd and 3rd ionisation energies which means that this element must be in group 2 of the periodic table as the 3rd electron is removed from an electron shell closer to the nucleus with less shielding and so has a larger ionisation energy.

The first ionisation energy of the elements

• 
• The shape of the graph for periods two and three is similar. A repeating pattern across a period is called periodicity. The pattern in the first ionisation energy gives us useful information about electronic structure.
• YOU NEED TO CAREFULLY LEARN THE PATTERNS!

Questions:

• Why dos Helium have the largest first ionisation energy?
• Why do first ionisation energies decrease down a group?
• Why is there a general increase in first ionisation energy across a period?
• Why does Na have a much lower ionisation energy than Ne?
• Why is there a small drop from Mg to Al?
• Why is there a small drop from P to S?
• Hint: many questions can be answered by applying the 3 factors of ionisation energy.

Answers:

• Its first electron is in the first shell closest to the nucleus and has no shielding effects from inner shells. He has a bigger first ionisation energy then H because it has one more proton.
• Going down a group the outer electrons ae found in shells increasingly further from the nucleus which are more shielded so the attraction of the nucleus becomes smaller.
• Going across the period the number of protons in the nucleus increases making the attraction of the nucleus greater. The electrons are being added to the same shell which has the same shielding effect and the electrons are being pulled closer to the nucleus.
• Na will have its outer electrons in a 3s shell further away from the nucleus and thus more shielded; the outer electron of Na is easier to remove and has a lower ionisation energy.
• Al is starting to fill a 3p sub-shell while Mg has its outer electrons in the 3s sub-shell. The electrons in the 3p sub-shells are slightly easier to remove because the 3p electrons are higher in energy and are also slightly shielded by the 3s electrons.
• S has 4 electrons in the 3p sub-shell and the 4th is starting to doubly fill the first 3p orbital. When he second electron is added to the first 3p orbital a slight repulsion occurs between the two negatively charged electrons which makes the second electron easier to remove.
• LEARN THE LAST TWO EXPLANATIONS AS THY O AGAINST THE USUAL ORDER!