Experimental evidence for electron structures After this lesson you should be able to: Recall the definitions for 1st and successive ionization energies. Understand the factors affecting the sizes of ionization energies. Show how ionization energies provide evidence for shells

Starter At the very end of the last lesson we endeavoured to come up with a set of principles to help explain ways in which electrons fill orbitals. I am now going to tell three simples rules to help you understand the process. Aufbau/building principle: electrons always fill the lowest energy orbitals first. Hund's rule: electrons never pair up in the same orbital until all orbitals of the same energy are singly occupied, and all unpaired electrons have parallel spin. Pauli exclusion principle: only two electrons may occupy the same orbital, and they must do so with opposite spin. Electron arrangement in orbitals There are three rules which determine the way in which electrons fill the orbitals

You have 5 minutes to come up with a quick way to remember these rules, any mnemonic will do but I would prefer it if you would try to construct it as a rap/ poem/ song etc. Here is my bad attempt- surely yours has to be better-surely!! His name was Aufbau, for sure he had a principle He told us how the orbitals would always be filled Electrons he said- well they really have a thirst, To fill the lowest energy orbitals, and they wanna fill them first! So you get the idea! Try it yourselves, best one gets homework off, lamest one gets homework off!

Ions and ionization energy Now, time to get serious! Ions and ionization energy Ions: An atom can either lose or gain an electron to form an ion: a charged atom. A positive ion is formed when an atom loses electron(s). For example a lithium atom forms a positive ion with a 1+ charge by losing an electron Li + _ + e Li 3p+,4n,3e- 3p+, 4n, 2e-

A negative ion is formed when an atom gains electrons A negative ion is formed when an atom gains electrons. For example, an oxygen atom forms a negative ion with a 2- charge by gaining two electrons. Write what the electron configurations would look like - you have 2 minutes! 2- O + 2e- O 6p+, 6n, 6e- 6p+, 6n, 8e-

Ionization energy Ionization energy measures the ease with which electrons are lost in the formation of positive ions. An element has as many energies as there are electrons, therefore: The first ionization energy of an element is the energy required to remove one electron from each atom in 1 mole of free gaseous atoms to form 1 mole of gaseous 1+ ions. The equation representing the first ionisation energy of sodium would be shown as: Na(g) Na+(g) + e- First ionization energy = +496kj mol-1

Factors affecting ionization energies Electrons are held in their shells by attraction from the nucleus, the 1st electron lost will be from the highest occupied energy level. This electron experiences least attraction from the nucleus. There are three factors that affect the size of the attraction Nuclear charge: The greater the number of protons in the nucleus, the greater the attractive force. Atomic radius: The greater the distance between the nucleus and the outer electrons, the less the attractive force. Attraction falls rapidly with increasing distance (a lot most like long distance relationships)! The distance is a very important factor and has a big effect. Electron shielding (screening): The outer shell electrons are repelled by any inner shells between the electrons and the nucleus. This repelling effect known as electron shielding, reduces the overall attractive force experienced by the outer electrons.

1st task: The three factors we just looked at are very important and may well help to explain many, many chemical ideas throughout the course. They are a basic grounding in the process of chemistry. In pairs you will explain in simple terminology the three factors affecting ionization energy (from your head, you may not use the books). Your partner must then create a pamphlet on the three factors, while your partner is doing this you must design a power point presentation on the same topic. At the end of the lesson, you will swap these resources- they will get yours and you will get theirs. – you have 20 mins to complete your task.

The second ionization energy of an element is the energy required remove 1 electron from each atom in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ atoms The equation representing the second ionization energy of sodium would be shown as: Notice the pattern! Na+ (g) Na2+ (g) + e- 2nd ionization energy = +4563kj mol-1 It makes sense then that Sodium with electrons, has successive ionization energies. Successive ionization energies provide evidence for the different energy levels of the principal quantum numbers 11 11

Taking this a step further, we can now look at the 1st ionization energies of the first 20 elements in the periodic table :

There are various trends in this graph which can be explained by reference to the proton number and electronic configuration of the various elements. A number of factors must be considered: nuclear charge shielding effective nuclear charge electron repulsion See handout for further clarification:  

The first ionization energy of He is thus higher than that of H. The trends in first ionization energies amongst elements in the periodic table can be explained on the basis of variations in one of the four factors mentioned. Trend across period 1: Compare the first ionization energies of H and He. Neither have inner shells, so there is no shielding. He has two protons in the nucleus; H only has one. Therefore the helium electrons are more strongly attracted to the nucleus and hence more difficult to remove. The first ionization energy of He is thus higher than that of H. Since H and He are the only atoms whose outer electrons are not shielded from the nucleus, it follows that He has the highest first ionization energy of all the elements. All elements (except H) have outer electrons which are shielded to some extent from the nucleus and thus are easier to remove. So Helium has the highest first ionization energy of all the elements. Trends across period 2   Compare now the first ionization energies of He (1s2) and Li (1s22s1). Li has an extra proton in the nucleus (3) but two inner-shell electrons. These inner-shell electrons cancel out the charge of two of the protons, reducing the effective nuclear charge on the 2s electron to +1. This is lower than the effective nuclear charge on the He 1s electrons, +2, and so the electrons are less strongly held and easier to remove. The first ionization energy of Li is thus lower than that of He. In general, the first ionization energy increases across a period because the nuclear charge increases but the shielding remains the same.

Trends down a group: Successive ionization energies The second ionization energy of an atom is the energy required to remove one electron from each of a mole of free gaseous unipositive ions. M+(g)  M2+(g) + e Other ionization energies can be defined in the same way: The third ionization energy of an atom is the energy required to remove one electron from each of a mole of bipositive ions. M2+(g)  M3+(g) + e The nth ionization energy can be defined as the energy required for the process M(n-1)+(g)  Mn+(g) + e It always becomes progressively more difficult to remove successive electrons from an atom; the second ionization energy is always greater than the first, the third always greater than the second and so on. On descending a group, the effective nuclear charge stays the same but the number of inner shells increases. The repulsion between these inner shells and the outer electrons makes them less stable, pushes them further from the nucleus and makes them easier to remove.

Variations in 1st ionization energies across a period in the periodic table provides evidence for the existence of sub-shells. There are trends we should notice with a little more academic training. For example we might spot that there are some points where we would notice a sharp decrease in 1st ionization energy between the end of one period and the start of the next (He Li); (Ne Na); (Ar K). What might this indicate? Homework off for the right answer.. This reflects the addition of a new outer shell with the resulting increase in shielding and distance

For example: lets look at Silicon We can use trends such as differences between ionisation energy values to indicate directly the electronic configuration of an atom For example: lets look at Silicon Large jumps occur between 4th and 5th and between 12th and 13th. Therefore there are three shells: The first contains 2 electrons, the second 8 and the third 4.

Finally The successive ionization energies of an atom always increase. The more electrons that are removed, the fewer the number electrons that remain. There is therefore less repulsion between the electrons in the resulting ion. The electrons are therefore more stable and harder to remove. By far the largest jumps between successive ionization energies come when the electron is removed from an inner shell. This causes a large drop in shielding, a large increase in effective nuclear charge and a large increase in ionization energy

Last task: I am about to show you a chart of Ionisation energies, you will be given this chart on a handout (going around now). I want you to explain what is happening on the chart and continue to explain what you have learned about ionisation energy values in this lesson.

Wrap up: Today was a long information heavy lesson, in your own words let me know what you learned today. Was it information overload or how did you cope?