IONISATION ENERGY OBJECTIVES: To define the term ‘ionisation energy’ To describe and explain trends in ionisation energy KEY WORDS: IONISATION ENERGY NUCLEAR CHARGE TREND SHIELDING FIRST THOUGHTS… What do you think we might mean by ionisation energy?

IONISATION ENERGY WHAT IS IONISATION ENERGY? “The energy required to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions” Energy required to remove electrons Measured in kJmol-1 Given the abbreviation IE Values tell us a lot about the electronic configuration of elements

IONISATION ENERGY SUCCESSIVE IONISATION ENERGIES “The energy required to remove each electron in turn. For example the second ionisation energy is the energy required to remove one electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions.“ Successive Ionisation Energies of Na

IONISATION ENERGY SUCCESSIVE IONISATION ENERGIES Na (g)  Na+ (g) + e- 1st IE = + 496 kJmol-1 Na+ (g)  Na2+ (g) + e- 2nd IE = + 4563 kJmol-1 Na2+ (g)  Na3+ (g) + e- 3rd IE = + 6913 kJmol-1 Notice how successive ionisation energies are written Start with the end product of previous ionisation event MUST include the gaseous state!

IONISATION ENERGY SUCCESSIVE IONISATION ENERGIES We can measure the energies required to remove each electron in turn, from outer electrons to inner electrons Values increase with each successive ionisation event Why might this be? Successive Ionisation Energies of Na

IONISATION ENERGY SUCCESSIVE IONISATION ENERGIES To form a positive ion, energy is needed to overcome attraction from nucleus. As each electron is removed from an atom, the remaining ion becomes more positively charged. Removing the next electron away from an increasing positive charge is more difficult, so the ionisation energy is even larger.

IONISATION ENERGY SUCCESSIVE IONISATION ENERGIES From this graph we can see evidence for electron shells Na looks to have 1 electron that’s easy to remove  furthest from nucleus 8 nearer to the nucleus  bit harder to remove 2 very close to the nucleus  nearest to +ve charge and hardest to remove

How are the big jumps useful? Si Difference in energy so different energy levels Big jump between the forth and fifth ionisation energies. Four electrons are relatively easily removed and must be in the outer shell. This indicates this element is in group 4

Potassium What would the ionisation energy plot look like for potassium? K

IONISATION ENERGY FACTORS AFFECTING IONISATION ENERGY As electrons are negatively charged and protons in the nucleus are positively charged, there will be an attraction between them. The greater the pull of the nucleus, the harder it will be to pull an electron away from an atom. Nuclear attraction of an electron depends on: Atomic radius Nuclear Charge Electron shielding or screening

IONISATION ENERGY ATOMIC RADIUS: Greater the atomic radius, the smaller the nuclear attraction experienced by the outer electrons NUCLEAR CHARGE: The greater the nuclear charge, the greater the attractive force on the outer electrons Lithium has a greater nuclear charge so you may expect it to have a higher ionisation energy. However it is affected to a greater extent by the atomic radius and the electron shielding Hydrogen Helium Lithium 519 kJ mol-1 1310 kJ mol-1 2370 kJ mol-1

IONISATION ENERGY ELECTRON SHIELDING Inner shells of electrons repel the outer-shell electrons Known as ‘Electron shielding’ or ‘screening’ More inner shells  larger the screening and smaller the nuclear attraction of outer electrons Hydrogen Helium Lithium 519 kJ mol-1 1310 kJ mol-1 2370 kJ mol-1

IONISATION ENERGY TRENDS ACROSS PERIODS Look at the graph 1) Describe the general patterns you see 2) Can you suggest an explanation for them?

IONISATION ENERGY TRENDS ACROSS PERIODS There is a ‘general increase’ across a period. This is because nuclear charge is increasing. The value drops dramatically for the start of another period.This is caused by an increase in atomic radius and electron shielding. He Ne Ar Kr Xe

IONISATION ENERGY TREND DOWN GROUPS Q: Describe and explain the general trend in ionisation energies going down Group 2.

IONISATION ENERGY TREND DOWN GROUPS General trend is a decrease as we go down a group The outer electron is in a level that gets further from the nucleus – greater atomic radius. Nuclear charge increases, however this is more than cancelled out by the electron shielding from lower energy shells

Extension questions Sketch a successive ionisation energy plot for: a) Nitrogen b) Neon c) Chlorine d) Iron

Recap Na

One electron needs much less energy for its removal than do the others One electron needs much less energy for its removal than do the others. It must therefore be in a higher energy level relative to the nucleus. Eight electrons require much more energy and therefore must be in a lower energy level. The last two electrons must be in an even lower energy level and so are much harder to remove.

The energy levels of electrons in a sodium atom

Sub-levels In the plot of the first ionisation energies for the first 20 elements, groups of 2,8,8 electrons are apparent. The groups have sharp breaks between them which indicates the filling of different quantum shells. The groups of eight electrons are subdivided into sub-groups of 2,3,3. What does this tell us?

Sub-levels In the plot of the first ionisation energies for the first 20 elements, groups of 2,8,8 electrons are apparent. The groups have sharp breaks between them which indicates the filling of different quantum shells. The groups of eight electrons are subdivided into sub-groups of 2,3,3. This gives evidence for subshells or different energy levels within each shell.

GCSE Atomic Structure There is nothing wrong with the GCSE atomic structure model. It still works for ions, dot and cross diagrams, bonding and structure. Ca 2.8.8.2

A-Level Atomic Structure However a more detailed atomic structure has been deduced and is used in parts of the AS/A2 course. The structure was deduced from successive ionisation energies and first ionisation energies. Ca 2.8.8.2

Moving between models Instead of shells, we now use energy levels. 4 The first shell becomes the first energy level The second shell becomes the second energy level And so on 4 3 2 1

Moving between models Each energy level is divided into sub-shells. 4 The first energy level has the 1s sub-shell (holds 2e-). The 2nd energy level has the 2s sub-shell (2e-) and the 2p sub-shell (6e-). 4 3 2p 2s 2 1s 1

Moving between models 4p 3d 4s The 3rd energy level has the 3s sub-shell (2e-), the 3p sub-shell (6e-), the 3d subshell (10e-). The 4th energy level has the 4s sub-shell (2e-), the 4p sub-shell (6e-). This takes us to Krypton. 3p 4 3s 3 2p 2s 2 1s 1

Energy levels video: Write down the rules for filling orbitals file://localhost/Users/RHH/Documents/Abingdon/Old%20resources/KS5/Year%2012/Resources%20from%20hodder%20cd/EDXChemAS_001640_orbitals.htm KS5>Year12>Resourcesfromhoddercd>001640

Hydrogen 4p 3d 4s 3p 4 3s 3 H 2p H 2s 2 1s 1 1 1s1

Helium 4p 3d 4s 3p 4 3s 3 2p He 2s 2 1s 1 2 1s2

Lithium 4p 3d 4s 3p 4 3s Li 3 2p 2s 2 1s 1 2.1 1s2 2s1

Copy and complete this table Element Atomic number Electronic configuration Ar 18 1s22s22p63s23p6 F 9 Al Na B N 7 S Cl

Write down the electronic configuration for the following: Lesson 3 Write down the electronic configuration for the following: Calcium Vanadium Iron Bromine

Calcium 4 3 2 1 Ca 2.8.8.2 1s2 2s2 2p6 3s2 3p6 4s2 4p 3d 3p 2p 2s 1s

Vanadium 4 3 2 1 2.8.11.2 1s2 2s2 2p6 3s2 3p6 3d34s2 4p 3d 3p 2p 2s 1s

Iron 4p 2.8.14.2 3d 4s 3p 4 3s 3 2p 2s 2 1s 1 1s2 2s2 2p6 3s2 3p6 3d64s2

Bromine 4p 2.8.18.5 3d 4s 3p 4 3s 3 2p 2s 2 1s 1 1s2 2s2 2p6 3s2 3p6 3d104s2 4p5

The 4s energy level is below the 3d energy level and so the 4s fills first.

Shorthand: The noble gases can be used in brackets e. g Shorthand: The noble gases can be used in brackets e.g. Potassium becomes [Ar] 4s1

Ions When positive ions are formed electrons are removed from the highest energy orbital When negative ions are formed electrons are added to the highest energy orbitals

Exceptions: Sc to Zn Once the 4s orbital has been filled it is at a higher energy level than the 3d orbital. So the 4s electrons are lost first (before the 3d electrons)

Periodic Table s-block d-block p-block f-block The chemical similarities which exist among members of a group of elements are a result of their similar electron arrangement.

Complete the Starter for 10 activity. Hund's Rule states that: Every orbital in a sublevel is singly occupied before any orbital is doubly occupied. All of the electrons in singly occupied orbitals have the same spin.

Sketch this graph onto your A3 paper Sketch this graph onto your A3 paper. Use what you know about the factors affecting ionisation energy and electron orbitals to explain the patterns seen in the graph. Add as much detail as you can!

1.6 IONISATION ENERGY EXPLANATION HYDROGEN EXPLANATION Despite having a nuclear charge of only 1+, Hydrogen has a relatively high 1st Ionisation Energy as its electron is closest to the nucleus and has no shielding. 1st IONISATION ENERGY / kJmol-1 1s ATOMIC NUMBER 1

1.6 IONISATION ENERGY EXPLANATION HELIUM EXPLANATION Helium has a much higher value because of the extra proton in the nucleus. The additional charge provides a stronger attraction for the electrons making them harder to remove. 1st IONISATION ENERGY / kJmol-1 1s ATOMIC NUMBER 2

1.6 IONISATION ENERGY EXPLANATION LITHIUM EXPLANATION There is a substantial drop in the value for Lithium. Despite the increased nuclear charge, there is electron shielding from the 1s orbital. The 2s electron is also further away from the nucleus. It is held less strongly and needs less energy for removal. 1st IONISATION ENERGY / kJmol-1 1s 1s 2s ATOMIC NUMBER 3

1.6 IONISATION ENERGY EXPLANATION BERYLLIUM EXPLANATION The value for Beryllium is higher than for Lithium due to the increased nuclear charge. There is no extra shielding. 1st IONISATION ENERGY / kJmol-1 1s 1s 2s 1s 2s ATOMIC NUMBER 4

1.6 IONISATION ENERGY EXPLANATION There is a DROP in the value for Boron. This is because the extra electron has gone into one of the 2p orbitals. The increased shielding makes the electron easier to remove It was evidence such as this that confirmed the existence of sub-shells. If there hadn’t been any sub-shell, the value would have been higher than that of Beryllium. 1s BORON 1st IONISATION ENERGY / kJmol-1 1s 1s 2s 1s 2s 2p 1s 2s ATOMIC NUMBER 5

1.6 IONISATION ENERGY EXPLANATION CARBON EXPLANATION The value increases again for Carbon due to the increased nuclear charge. 1st IONISATION ENERGY / kJmol-1 1s 1s 2s 2p 1s 2s 1s 2s 2p 1s 2s ATOMIC NUMBER 6

1.6 IONISATION ENERGY EXPLANATION NITROGEN EXPLANATION The value increases again for Nitrogen due to the increased nuclear charge. As before, the extra electron goes into the vacant 2p orbital. There are now three unpaired electrons. 1s 2s 2p 1st IONISATION ENERGY / kJmol-1 1s 1s 2s 2p 1s 2s 1s 2s 2p 1s 2s ATOMIC NUMBER 7

1.6 IONISATION ENERGY EXPLANATION There is a DROP in the value for Oxygen. The extra electron has paired up with one of the electrons already in one of the 2p orbitals. The repulsive force between the two paired-up electrons means that less energy is required to remove one of them. 1s OXYGEN 1s 2s 2p 1st IONISATION ENERGY / kJmol-1 1s 1s 2s 2p 1s 2s 2p 1s 2s 1s 2s 2p 1s 2s ATOMIC NUMBER 8

1.6 IONISATION ENERGY EXPLANATION FLUORINE EXPLANATION The value increases again for Fluorine due to the increased nuclear charge. The 2p orbitals are almost full. 1s 2s 2p 1s 2s 2p 1st IONISATION ENERGY / kJmol-1 1s 1s 2s 2p 1s 2s 2p 1s 2s 1s 2s 2p 1s 2s ATOMIC NUMBER 9

1.6 IONISATION ENERGY EXPLANATION NEON EXPLANATION The value increases again for Neon due to the increased nuclear charge. The 2p orbitals are now full so the next electron in will have to go into the higher energy 3s orbital. 1s 2s 2p 1s 2s 2p 1s 2s 2p 1st IONISATION ENERGY / kJmol-1 1s 1s 2s 2p 1s 2s 2p 1s 2s 1s 2s 2p 1s 2s ATOMIC NUMBER 10

1.6 IONISATION ENERGY EXPLANATION There is a substantial drop in the value for Sodium. This is because the extra electron has gone into an orbital in the next energy level. This means there is an extra shielding effect of filled inner 1s, 2s and 2p energy levels. 1s SODIUM 1s 2s 2p 1s 2s 2p 1s 2s 2p 1st IONISATION ENERGY / kJmol-1 1s 1s 2s 2p 1s 2s 2p 1s 2s 1s 2s 2p 1s 2s 1s 2s 2p 3s ATOMIC NUMBER 11

1.6 IONISATION ENERGY EXPLANATION The value for Magnesium is higher than for Sodium due to the increased nuclear charge. There is no extra shielding. The trend is similar to that at the start of the 2nd period. 1s MAGNESIUM 1s 2s 2p 1s 2s 2p 1s 2s 2p 1st IONISATION ENERGY / kJmol-1 1s 1s 2s 2p 1s 2s 2p 1s 2s 1s 2s 2p 3s 1s 2s 2p 1s 2s 1s 2s 2p 3s ATOMIC NUMBER 12

1.6 IONISATION ENERGY EXTENSION QUES: Which has the higher value, the 1st I.E. of sodium or the 2nd I.E. of magnesium? A: The 2nd I.E. of magnesium The 1st I.E. of sodium involves the following change Na(g) Na+(g) 1s2 2s2 2p6 3s1 1s2 2s2 2p6 The 2nd I.E. of magnesium involves the same change in electron configuration… Mg+(g) Mg2+(g) 1s2 2s2 2p6 3s1 1s2 2s2 2p6 However, magnesium has 12 protons in its nucleus, whereas sodium only has 11. The greater nuclear charge means that the electron being removed is held more strongly and more energy must be put in to remove it.

1.6 IONISATION ENERGY TREND IN PERIOD 3 Draw a graph for the first ionisation energy of elements in period 3 For each point draw the box and arrow electron configuration For each point explain why the first ionisation energy either increases or decreases Na Mg Al Si P S Cl Ar 496 738 578 789 1012 1000 1251 1521

1.6 IONISATION ENERGY TREND IN PERIOD 3 Drop from Mg  Al: Mg: 1s2 2s2 2p6 3s2 Al: 1s2 2s2 2p6 3s2 3p1 Drop from P  S: P: 1s2 2s2 2p6 3s2 3p3 S: 1s2 2s2 2p6 3s2 3p4 The outer electron in Al has moved into the 3p orbital. It takes less energy to remove an electron from here than the 3s In P each of the 3p orbitals has 1 electron. In S one must have 2. The repulsion between these 2 paired electrons makes it easier to remove one

Explain the pattern in ionisation energies between a) B and Be b) O and N

B has lower ionisation energy than Be B has lower ionisation energy than Be. This is because Be has a complete 2s shell which is stable (so hard to remove e-). B has one e- in the 2p shell which is shielded from the nuclear charge by the 2s electrons (so it is easier to remove).

O has a lower ionisation energy than N because of the stability of the ½ filled shell for N (2p3) which makes it less favourable to remove an electron. It is easier to remove an e- from O because it has more than a ½ filled shell. The removal of one electron from Oxygen will create the stable 2p3 arrangement.

After Ca (atomic number 20) the 2,3,3 pattern is broken How many elements produce the break? What might this indicate in terms of electrons? What is the name of this set of elements in the Periodic Table?

Questions P61 The ten elements after Ca do not have the 2,3,3 pattern. This is because the elements after Ca have electrons in the d-sublevel which can hold 10 electrons (not 6 as in a p-sublevel which gives the 3,3 pattern). The name of the elements in the Periodic Table which are d electrons are the transition metals.