Atomic Structure

The basics of Atomic Structure. AS Chemistry. Suitable for A grade revision :)


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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

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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

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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

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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.

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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
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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.

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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


  • 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

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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)

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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)

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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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