Electrons


Exciting Electrons: Every element emits light if it is heated. This is because the atom has absorbed energy, and then loses it in the form of light. The light can be separated to form an Electron Emission Spectrum. Each line on the EES corresponds to a particular frequency of light being given off by an atom; therefore each corresponds to one exact amount of energy being emitted.
Emission Spectrum from various elements


 Niels Bohr proposed an explanation for the ES, stating that electrons of specific eneergy move around the nucleus in orbits or energy levels, and cannot exist between these levels. When given excess energy (from flame or electric current) electrons are able to move to a higher obit. this is from ground state (lowest energy level) to an excited state. As the electron moves back to ground state, the amount of energy given off is the difference in energy between the levels, and is given off as a photon of light. In an ionic compound, it is the anion (negatively charged) that determines the colour. 

 Electron Shells: 
Electrons move around the nucleus in a region of space and labelled K/1, L/2, M/3, N/4. Each has a specific energy level that increases as you move further from the nucleus.
Each can hold a maximum number of electrons: 2n2
Electrons show particle and wave behaviours.

 Shell Model Diagrams:
Represent electrons around the nucleus, however are limited as they suggest that electrons orbit in perfectly circular paths (which they don't), the order of filling of shells and differences in energies. 


 Unfortunately, it gets more complicated. Electrons exist in orbitals, or "regions of space where you are likely to find them", and no more than 2 electrons can exist in each of these. Therefore, we have:
Shells: Major energy levels within an atom (1,2,3,4,5 ect)
Sub-shells: Energy levels within a shell (s,p,d,f)
Orbitals:  Regions of space where electrons move

 Generally, the order of filling is from lowest energy level first.

 Electron Configuration: 

Generally, the order of filling is from lowest energy first.
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p

 Example: Sc (Z=21) 1s22s22p63s23p64s23d1 
2p6 
2- Shell number
S - Orbital
6 - Number of electrons

 It is possible for two atoms to be different elements, but have the smae electronic configuration, if they are ions. 
Example: Oxygen: 1s22s22p4 
Oxygen-2 :1s22s22p6 
Neons: 1s22s22p6 

 Excited States: When electrons move to higher energy levels, the elctron configuration changes. The outermost electron moves to a higher energy level subshell.
To identify excited states, look for electrons in higher energy shells / gaps and abnormal filling order. 
Example: Neon in ground state is 1s22s22p6 
but in an excited state, may become 1s22s22p53s1 

 Exceptions: Atypical electron configurations are due to some atoms/elements being more stable this way. 
Example: Copper (29) and Chromium (23) half fill their d-shells, as this is more stable than partially filled. 
Cu : 1s22s22p63s23p64s13d10 not 4s23d9
Cr : 1s22s22p63s23p64s13d5 not 4s23d4

 Finding Group Number from Electron Configuration: using the highest energy level
Number of shells/period: highest level 'large' number
Letter: block
Little number: how many across block
Example: Scandium: 1s22s22p63s23p64s23d1
-> 3d1 Period 3, column 1 of d block   

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