Atomic Structure
The basics of Atomic Structure. AS Chemistry. Suitable for A grade revision :)
NOTE: I USE BRACKETS TO INDICTATE LOW NUMBERS IN MOLECULES, FOR EXAMPLE: H(2)O
- Created by: Natalie Beard
- Created on: 14-12-09 11:20
Protons, Neutrons and Electrons
Protons: + 1 charge Weigh one unit Nucleus
Neutrons 0 charge Weigh one unit Nucleus
Electrons -1 charge Negligible In Orbit
Atomic Number: Number of Protons in an atom, (also number of electrons as these must be equal in a stable atom.)
Atomic Mass: Number of neutrons and protons in the nucleus combined
Isotopes: Atoms with different mass numbers, same number of protons and electrons, different number of neutrons.
Atoms are usually stable, equal charge, but they can become ions, with diff. charge
RAM and RMM
Relative Atomic Mass: The average mass of one atom of one element compared to 1/12th of Carbon- 12.
Essentially the mass number of the atom. In isotopes it is the average.
Example: Oxygen = 16
Relative Molecular Mass: The average mass of a molecule compared to 1/12th of Carbon- 12.
The number given when the RAM of each element is added up
Example: H(2)O = H x 2= 2 + O= 16= 18
The Mass Spectrometer
This measures the massses of atoms and molecules, can also be used to calculate the relative abundance of different isotopes and to predict complex molecular structure.
Vapourisation: Sample released in vapour form into ionization chamber
Ionisation: Particles bombarded by electron gun, this knocks an electron off so particles are now positive ions, (if it is set too strong, more are sometimes knocked off)
Accleration: Positive ions attracted to negatively charged metal plates and this acclerates it via electric field
Deflection: Ions attracted to magnetic field and deflected according to mass, lighter they are, more deflected and more charged they are, more deflected. M/Z ratio.
Detection: Electric current measured as ions land on plate, greater the abundance of an isotope, the larger the current. = VIADD
Calculations using Mass Spectrometer
Usually they'll give you a graph and ask you to work out the following:
Relative Atomic Mass:
Sum Of (Percentage abundance of each isotope x Mass of each isotope) / 100
Relative Molecular Mass: Simply look at the largest peak on the graph, no calculations required.
Somtimes, the molecules can't cope with the extreme conditions and so break apart. This is known as fragmentation.
Electronic Configuration
Electrons occupy fixed energy levels
- Principal: Principal energy levels are numbered consequatively, with 1 being the smallest and the closest to the nucleus, whereas infinity would not be attracted the the electron at all.
- Subsidary levels exist in each atom and are given the letters S, P, D and F. These correspond to an increase in energy, each additional electron goes into the sub-level with the least energy.
The order is: 1s(2) 2s(2) 2p(6) 3s(2) 3p(6) 4s(2) 3d(10) 4p (6) etc etc
- Each principal and subsidary levels fill up before the next one
- When the 3p energy level is full, the 4s fills before the 3d
- When this is shown in written form, the 3d, is written first
- Atoms with 3d electrons always lose 4s first when forming ions
Ionisation Energy
The energy required to remove the most loosely held electron from one mole of gaseous atoms to produce one mole of gaseous ions each with a charge of 1+. The equation for the first ionisation energy= X(g) ----------> X+(g) + E- . Factors listed below:
Charge on nucleus: The more protons in a nucleus, the more positively charged it is and the more strongly electrons are attracted to it, this would meran an increase in ionisation energy.
Distance between electrons and nucleus: Attraction falls rapidly with distance, the closer to the nucelus, the more strongly attracted the electrons are, this results in a higher ionisation energy.
Shielding: Electrons in the inner levels repel the outer ones, this lessens the pull of the nucleus by the inner electrons and is known as shielding.
Paired or Not: Two electrons in the same orbital will experaince repulsion from each other. This offsets the attraction of the nucleus so paired electrons are easier to remove.
Ionisation energies in Group 2 (Be-Ba)
Group 2 general trend= Decreases as group descends
Ionisation Energy
- Be= 900
- Mg= 738
- Ca= 590
- Sr= 550
- Ba= 503
Why
- Outer electrons have progressively higher energy levels
- Further from nucleus
- Shielded by more complete inner levels
So less strongly attracted and less energy is required to move them
Ionisation energies in Period 3 (Na-Ar)
Na to Mg: Small increase
Additional electron goes into same 3s shell. Mg has one more proton, (increse in nuclear charge) so electron more strongly attracted.
Na: 1s(2) 2s(2) 2p(6) 3s(1)
Mg: 1s(2) 2s(2) 2p(6) 3s(2)
Mg to Al: Small decrease
Additional electron goes into 3p sub shell which is higher in energy than 3s shell. (Distance from nucleus and more shielding.)
Mg: 1s(2) 2s(2) 2p(6) 3s(2)
Al: 1s(2) 2s(2) 2p(6) 3s(2) 3p(1)
Ionisation energies in Period 3 (Na-Ar)
Al to Si to P: Increase
Electrons fill P orbitals, increased nuclear charge increases attraction for the electrons.
Al: 1s(2) 2s(2) 2p(6) 3s(2) 3p(1)
Si: 1s(2) 2s(2) 2p(6) 3s(2) 3p(2)
P: 1s(2) 2s(2) 2p(6) 3s(2) 3p(3)
P to S: Small decrease
Additonal electron into P orbital, already has an electron and paired electrons repel.
P: 1s(2) 2s(2) 2p(6) 3s(2) 3p(3)
S: 1s(2) 2s(2) 2p(6) 3s(2) 3p(4)
Ionisation energies in Period 3 (Na-Ar)
S to Cl to Ar: Increase
Addtional electrons enter 3p orbitals + increae in nulear charge= electrons more held
S: 1s(2) 2s(2) 2p(6) 3s(2) 3p(4)
Cl: 1s(2) 2s(2) 2p(6) 3s(2) 3p(5)
Ar: 1s(2) 2s(2) 2p(6) 3s(2) 3p(6)
Ar to K: Large decrease
Additional electron into new energy level- more distance from nucleus + more shielding.
Ar: 1s(2) 2s(2) 2p(6) 3s(2) 3p(6)
K: 1s(2) 2s(2) 2p(6) 3s(2) 3p(6) 4s(1)
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