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# 4 The Atom & Electronic Configuration

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### 4 The Atom & Electronic Configuration

1. 1. 4 The Atom & Electronic Configuration
2. 2. Bohr Model of the Atom <ul><li>Bohr suggested that an electron moving in an orbit can only have certain amounts of energy not an infinite number of value. This energy is quantised . </li></ul><ul><li>The energy that an electron needs in order to move in a particular orbit depends on the radius of the orbit. </li></ul><ul><li>An electron in an orbit further away from the nucleus requires higher energy than an electron nearer the nucleus. </li></ul><ul><li>If the energy of the electron is quantised then the radius of the orbit must also be quantised.  there are a restricted number of orbits. </li></ul>
3. 3. <ul><li>Ionisation Energy </li></ul><ul><li>Definition The 1 st ionisation energy of an element is the energy change for the conversion of 1mol of gaseous atoms into 1 mol of gaseous ions </li></ul><ul><li>i.e the energy change for the process </li></ul><ul><li>M (g)  M + (g) + e </li></ul><ul><li>Similarly the 2 nd ionisation energy is the energy change for the process </li></ul><ul><li>M + (g)  M 2+ (g) + e </li></ul><ul><li>e.g. 1st ionisation energy 2 nd ionisation energy </li></ul><ul><li>H = 1310 KJ/mol </li></ul><ul><li>He = 2370 KJ/mol 5250 KJ/mol </li></ul>
4. 4. <ul><li>The second electron is always harder to remove than the 1 st </li></ul>2p x x Helium contains 2 electrons and 2 protons. When 1 electron is removed there are still 2 positive charges in the nucleus  more attraction on the remaining electron which then requires more energy to remove.
5. 5. <ul><li>Lithium has 3 electrons and 3 protons </li></ul><ul><li>1 st I.E. = 519 KJ/mol </li></ul><ul><li>2 nd I.E. = 7300 Kj/mol </li></ul><ul><li>3 rd I.E. = 11800 Kj/mol </li></ul><ul><li>Notice the large increase in ionisation energy between the 1 st and 2 nd and then a smaller jump between the 2 nd and 3 rd </li></ul>The outer electron is easier to remove as it is in a shell further away from the nucleus x 3p x x
6. 6. <ul><li>Carbon has 6 electrons 2 in first shell. 4 in second shell </li></ul><ul><li>and 6 protons in the nucleus (electronic configuration 2,4) </li></ul>x x x The 4 electrons in the second shell are easier to remove than the 2 in the first shell as they are further from the nucleus. x 6p x x
7. 7. <ul><li>However if we plot a graph of ionisation energy against electron removed </li></ul>6 The first 2 electrons are relatively easy to remove, the 2 nd two slightly harder and the last 2 much harder. The first 4 are in the outer shell the last 2 are in the inner shell Notice the ease of removal is not a smooth trend from 1 -4 . There is a slight jump in ionisation energy between the 2nd and 3rd electron. This suggests that the 1& 2 electrons being removed are in a slightly lower energy level than the 3 & 4
8. 8. <ul><li>From this we conclude that the 2 nd shell consists of 2 subshells, one slightly closer to the nucleus than the other </li></ul>6p x x x x x x 1s 2s 2p We now write the electronic configuration as 1s 2 2s 2 2p 2
9. 9. <ul><li>The first 2 electrons (1,2) removed are in a 2p orbital </li></ul><ul><li>The next 2 electrons (3,4) removed are in 2s orbital </li></ul><ul><li>The last 2 electrons (5,6) removed are in a 1s orbital </li></ul><ul><li>There is also a third type of orbital called a d orbital. The first element with a d orbital is scandium (the first transition element) </li></ul><ul><li>ALL orbitals can hold 2 electrons. </li></ul><ul><li>The first shell has only 1 s orbital </li></ul><ul><li>The second shell has 1 s orbital and 3 p orbitals </li></ul><ul><li>The 3 rd shell has 1 s orbital, 3 p orbitals and 5 d orbitals </li></ul>
10. 10. <ul><li>The first shell can hold 2 electrons </li></ul><ul><li>The second shell can hold 8 electrons </li></ul><ul><li>The third shell can hold 18 electrons </li></ul><ul><li>We can now write electronic configurations as follows </li></ul><ul><li>Li 3electrons 1s 2 2s 1 </li></ul><ul><li>C 6 electrons 1s 2 2s 2 2p 2 </li></ul><ul><li>Na 11 electrons 1s 2 2s 2 2p 6 3s 1 </li></ul><ul><li>Note that the outer shell contains the same no of electrons as the group number. Carbon group 4 has 4 electrons in the 2 nd shell </li></ul>
11. 11. <ul><li>Finding the electronic configuration of ions </li></ul><ul><li>Li 1s 2 2s 1 </li></ul><ul><li>Li + 1s 2 </li></ul><ul><li>It has lost 1 electron and now has a filled shell and is stable </li></ul><ul><li>Be 1s 2 2s 2 </li></ul><ul><li>Be 2+ 1s 2 </li></ul><ul><li>it has lost 2 electrons and and now has a filled shell and is stable </li></ul>
12. 12. <ul><li>Write the configuration of the following elements </li></ul><ul><li>N Mg Al Cl </li></ul><ul><li>1s 2 2s 2 2p 3 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 6 3s 2 3p 1 1s 2 2s 2 2p 6 3s 2 3p 5 </li></ul><ul><li>Al 3+ F - Na + </li></ul><ul><li>1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 </li></ul>
13. 13. <ul><li>Elements in groups 1 and 2 are filling s orbitals and are called s block </li></ul><ul><li>Element in groups 2-8 are filling p orbitals and are called p block </li></ul><ul><li>The transition elements are filling d orbitals and are called d block </li></ul><ul><li>When we write a configuration we ‘fill’ orbitals from the lowest energy to the highest. i.e. electrons are added to the inner 1s orbital first followed by 2s then 2p etc. This is called the aufbau principal (to build up) An atom with all electrons in the lowest possible energy levels is said to be in its ground state </li></ul><ul><li>When writing configurations for the transition elements the 4s orbital is filled before the 3d (the 4s orbital is actually lower in energy than the 3d) </li></ul><ul><li>So the configuration of Sc is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 </li></ul><ul><li>It is the first transition metals and has 1 d electron. (It is easy to find the number of d electrons – just count from scandium across the row) </li></ul>
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