Topic 12 - Atomic Theory
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[edit] 12.1 The mass spectrometer
[edit] 12.1.1
The stages of operation are: vaporisation, ionisation, acceleration, deflection and detection.
Substance to be tested is vaporised (by heat, in the absence of oxygen) then ionised by electric current. Ions are accelerated through an electric field, then deflected by a magnetic field. Ions are then detected. The angle of deflection due to the magnetic field reflects their mass to charge ratio.
[edit] 12.1.2
The angle of deflection of each fragment is proportional to its mass, and so it is possible to find the relative atomic mass of each 'spike'. The height of the spike represents the frequency, therefore, the abundance can be calculated. The relative atomic mass is the weighted average of the isotope masses times their percentage abundance (frequencies).
[edit] 12.2 Electronic structure of atoms
[edit] 12.2.1
Successive electrons can be stripped from an atom until there is only the nucleus left. If the energy required to achieve this for each electron is plotted on a graph (with a log scale) against ionisation number, the 'jumps' in the required energy clearly show the main and sub energy levels.
[edit] 12.2.2
n = (1, s) (2, sp) (3, spd) (4 spdf)
[edit] 12.2.3
Energies of sub-shells : s < p < d < f
[edit] 12.2.4
Number of orbitals at each level : s=1, p=3, d=5, f=7
[edit] 12.2.5
Shapes of orbitals : s orbital is a sphere around the nucleus. p orbitals are shaped like a figure 8 (and there are 3 of them on perpendicular x, y and z axes around the nucleus).
[edit] 12.2.6
| 1s | |||
| 2s | 2p | ||
| 3s | 3p | 3d | |
| 4s | 4p | 4d | 4f |
| 5s | 5p | 5d | 5f |
Move diagonally down and left through each diagonal in turn, i.e. 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d.
Pauli's exclusion principle says that there can only be 2 electrons in each orbital (with opposite spins). Hund's rule says that each orbital should be half filled before any is completely filled (since there is less repulsion if all electrons have the same spin). Electrons will therefore fill the lowest energy levels (i.e. 1 then 2 and so on) with two going in each orbital, but only doubling up when all orbitals in the level are filled.
[edit] 12.2.7
Systematically fill the orbitals as shown above up to Z = 56. This can be abbreviated by writing [x] where x is a noble gas.
[edit] 12.2.8
In the periodic table, the small double column on the left is the s shell being filled. The block on the right is the p shell being filled. The d block (in the middle) is (surprisingly) the d shell being filled. The bits hanging off the bottom are the f shells being filled (forget them, they never matter).
[edit] 12.2.9
There are 4 numbers that define the position of an electron. It can be thought of the address of an electron. The first number is called the principal quantum number,n Its range is any natural number. The second number is the azimuthal quantum number, l , this determines the shape of the orbital. Its range is dependent on n. Its range is (0,n-1). The third number is the magnetic quantum number, m, it tells us the direction that the orbital is pointing (eg. it will differentiate between a p-orbital along the x-axis and a p-orbital along the y-axis) m can take values from -l,l. Finally, there is the electron spin A simple arrow up or down. This is from the Pauli exclusion principle that cannot be explained with the Schrodinger model of the atom.
The principal quantum number indicates which shell level the electron is in. The azimuthal quantum number indicates the shape of the orbital. (Common orbitals are as such; 0 = s , 1 = p , 2 = d ...etc) The magnetic quantum number gives the orientation of the orbital. The spin is just a bookkepping idea.
[edit] 12.2.10
The reason the Schrodinger model was used in preference over the Bohr model is because the Bohr model had one simple flaw. It assumed the electron was a particle moving around the central nucleus. However, an accelerating charge emits energy. Hence, the Bohr model suggested that the electron constantly emits energy. But this is impossible as the electron will be losing energy and end up in contact with the nucleus. Thus Rutherford's observation that the electrons orbit the nucleus will be incorrect. But more importantly the Heinsberg uncertainty principle will not hold.
Schrodinger fixed all this by assuming an atom composed only of wave functions.
