Most elements have isotopes which means that the mass of the element will depend on the relative amounts of isotopes present.
This gives rise to the term Relative Atomic Mass which is similar to the mass number but includes the differing masses of the isotopes.
Definition; The Relative Atomic Mass of an element Ar is the average mass of one atom relative to1/12th the mass of one atom of carbon-12. The fact that it is an average means that it takes into account the differing numbers of isotopes. The relative isotopic mass will be for one isotope only.
MASS SPECTROMETRY.
In order to find out the Relative atomic mass of an element we need to know the masses of the different isotopes and their relative abundancies.
This is done using a Mass Spectrometer.
It works in 4 stages. 1) The element sample is introduced as a vapour to the ionisation chamber. Here it is subjected to a beam of fast moving electrons that are created by a heated filament. These high-energy electrons collide with he atoms of the element and knock off electrons leaving positively charged ions. Hopefully each atom loses one electron only and so forms a 1+ ion.
2) These ions are then Accelerated by being subjected to an electric field, so that we now have a stream of fast moving positive ions travelling through the spectrometer The stream of fast moving ions is then deflected by passing through a magnetic field Lighter ions will be deflected more than heavier ones and so the ions are separated according to their masses. 3) Detection Ions of the same charge and mass will follow one particular path. This stream of ions is detected as an electrical impulse by an electrometer. This impulse is amplified and as an electric current works a pen. The size of the current depends on the number of ions hitting the detector and so the relative abundancies of the ions can be compared on the trace formed.
Determination of Ar of Neon
height of peak
20 21 22 m/z
Heights of peaks ; at 20 = 114
at 21 = 0.2
at 22 = 11.2
Step 1 Multiply abundancies (height of peaks ) by mass numbers to find mass of each isotope present e.g.
20 x 114 = 2280
21 x 0.2 = 4.2
22 x 11.2 = 246.4
Step 2 Find total abundance = 2530.6
Step 3 Divide total mass by total abundance
2530.6/125.4 = 20.2
Ar Ne = 20.2
It is possible to get peaks at m/z values of exactly half the highest values because of the possibility of forming 2+ ions in the ionisation chamber.
ELECTRON ARRANGEMENT.
Electrons are found in orbitals. An orbital is an area in which the chances of finding the electron are high, i.e. it is a probability cloud within which the electron will be found, or the electron density.
The electron density is high if there is a high probability of finding the electron in this area. E.g. hydrogen’s electron is likely to be fairly close to the nucleus and it’s orbital is a sphere which surrounds the nucleus, outside the extremity of this sphere it is very unlikely that the electron will be found
Different orbitals are found at different distances from the nucleus and so the electrons are at different energies. Orbitals that have very similar energies make up an energy level. Each energy level contains different numbers of orbitals, i.e. the first energy level contains only one sort of orbital, the s orbital which can hold only 2 electrons. Therefore hydrogen has only 1 orbital which holds only 1 electron but Helium has 1 orbital that contains 2 electrons which have similar energies. This energy level is called the first energy level and the orbital is called the 1s orbital which is a sphere of probability surrounding the nucleus.
The second energy level can hold 8 electrons found in 4 orbitals of 2 electrons. These orbitals are not all the same but include the spherical 2s and 3 lobe shaped 2p orbitals. The shapes differ in order to enable the electrons to get as close to the nucleus as possible. When the second energy level is full it will contain 2 s electrons and 6 p electrons written as 2s 2p .
The s and p orbitals are at slightly different energies and are often referred to as sub-levels. Therefore the second energy level is divided into 2 sub-levels.
The third energy level has s, p and d orbitals . The s can hold 2, p can hold 3 lots of 2 and d can hold 5 lots of 2. the fourth energy level also has f orbitals which can hold 7 lots of 2.
It is possible for each orbital to hold 2 electrons because they spin in opposite directions and so do not repel each other so much. The energy and orbitals are represented by boxes.
Draw boxes
Unfortunately the energy levels tend to overlap and the electrons in the 3d orbitals are at a slightly higher energy than 4s.
The rules for filling the energy levels and orbitals are lower energy first . Energy is needed to make the electrons pair therefore the orbitals fill singly first before pairing. The order of filling is made more complicated by the overlap of the different levels.
The following diagram can be used as a memory aid.
s p d f
1
- 2
- 3 3
- 4 4 4
- 5 5 5 5
- 6 6 6 6
- 7
Work down to the arrowhead and then go back up to the next available top one.
Putting electrons in boxes is convenient and easy to understand but rather time consuming. It can be abbreviated by just putting the level, sub-level and number of electrons. E.g. sodium can be represented by ;
1s 2s 2p 3s . They can be abbreviated further by using the nearest noble gas configuration. E.g. Ca will be ( Ar ) 4s .
You must practise writing electron configurations in boxes and shorthand.
IONS
An ion with +1 charge has one less electron than it’s parent atom because it is formed by the loss of an electron from the highest energy orbital therefore its electronic configuration will show this electron missing. E.g. Na 1s 2s 2p 3s.
Na + 1s 2s 2p
TRANSITION METALS
The energy of the 3d electrons is very similar to that of the 4s electrons, with the 3d being slightly higher than the 4s. However when transition metal atoms lose electrons to become ions the electrons are lost from the 4s first. This means that Sc is 1s 2s 2p 3s 3p 4s 3d
But Sc 2+ is 1s 2s 2p 3s 3p 4s 3d
EVIDENCE TO SUPPORT THIS THEORY FOR ELECTRONIC STRUCTURE.
We can use Ionisation Energies to support this picture of the atom.
Definition. The first ionisation of an element is the energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of unipositive gaseous ions.
A(g) A (g) + e
Ionisation energies are always positive because energy is required to remove electrons.
Considering the removal of successive electrons from an atom of an element the range of energy values is very large and so the figures used are the logs of the energy values. These then reflect the pattern of electrons in energy levels containing 2, 8, electrons.
Lg of
Successive
IE for K
No of electrons removed.
This graph of successive IEs for Potassium provides evidence to support the theory of energy levels. The gradual increase in 1st IE within an energy level is due to the electron being removed from an ion, which is already positive. It has more protons than electrons and so a negative electron is being removed from a positive ion, requiring more energy.
Note the pattern of energy levels containing 2 and 8 electrons.
This graph shows the broad energy levels but not the sub-levels. However information like this will enable you to decide which group of the Periodic Table an element is in. A large jump in 1st IE value corresponds to the start of a new energy level. If a similar graph was drawn for a group 2 element the pattern would show a gradual increase from electron 1 to 2 and then a large increase in the amount of energy needed to remove the third electron. Consideration of 1st IE values going across a period gives information about the sub-levels as well as the overall levels. See diagram below.
There is a general increase in 1st IE across the period because the electrons are going into the same energy level but are experiencing an increased nuclear pull because of the extra nuclear charge, causing a need for more energy to pull against this nuclear attraction hence higher IE. You will notice that there are little hiccoughs in this pattern.
Note that Boron has a slightly lower 1st IE than Beryllium. This corresponds to Be’s electron going into the slightly higher energy sub-level and therefore requiring a little less energy to be removed. I.e. it is the first of the 2p electrons.
Also Oxygen has a slightly lower 1st IE than Nitrogen because its’ electron is the first one to be paired in the 2p sub-level and so is experiencing repulsion from the other electron making it slightly easier to remove
This pattern is repeated for period 3. There is a general increase but Al has a lower 1st IE than Mg because its’ last electron is the first one to go into the 3p sub-level and S has a lower 1st IE than P because it is the first one to experience repulsion form an electron already in that sub-level
