AQA Chemistry Unit 1

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  • Created by: chapa
  • Created on: 22-05-13 13:54

Atomic Structure

Fundamental Particles:

  • Atoms consist of electrons surrounding a nucleus made up of protons and neutrons
  • Proton - Mass - 1 Charge - +
  •  Neutron - Mass - 1 Charge - neutral
  • Electron - Mass - negligible Charge - -

Mass number and Isotopes:

  • Mass number - number of protons + number of neutrons
  • Atomic number - number of protons in an atom

Chemical properties don't vary much for isotopes because they still have the same number of electrons and protons

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Mass Spectrometer

Vapourisation - turns into gas, allows atoms to move more freely

Ionisation - Gas particles fired with a beam of e-, forms a +1 ions

Acceleration - Accelerates ions into a beam

Deflection - Path of ions deflected by strong magnetic field, lighter ions deflected more

Detection - Only certain ions with a certain mass reach the detector, magnetic field is altered to produce spectrum

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Mass Spectrometer

  • Relative atomic mass - Average mass of an atom of an element on a scale when an atom of carbon-12 is 12
  • Relative isotopic mass - mass of an atom of an isotope of an element on a scale where carbon-12 is 12


How to work out Ar using a mass spectrum

  • For each peak, read relative isotopic abundance and relative isotopic mass
  • Multiply together for each peak, and add your answers up
  • Divide by 100 (because percentages were used)

Mr can also be found using mass spectrometry, you look at the very last peak and its the mass/charge and that's the Mr

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Electron Arrangement

  • Electron arranged in energy levels
  • 1s, 2s, 2p, 3s, 3p, 3d, 4s, 4p, 4d, 4f


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First Ionisation Energy

First ionisation energy is defined by: X(g) --> X+(g) + e-

Ionisation energy is the energy needed to remove one e- from each atom in 1 mole of gaseous atoms to form 1 mole of 1+ gaseous ions

Factors affecting ionisation energy:

  • Nuclear charge - more protons=bigger attraction to e- = harder to remove e-
  • Distance form nucleus - e- close to nucleus are more strongly attracted to nucleus that e- further away, making them harder to remove
  • Shielding - Further away the e- from the nucleus, the more e- in between the outer e- and nucleus, lessens attraction between nucleus and outer e-
  • High ionisation energy = high attraction between nucleus and outer e-
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First Ionisation Down Group 2

First Ionisation Energy Down Group 2 Decreases

  • Decreases down the group from Beryllium to Barium
  • Magnesium has a lower ionisation energy than Beryllium because its outer e- is in a 3s sublevel while Beryllium has its outer e- in a 2s sublevel
    • This means Magnesium has more shielding than the Beryllium as it has an extra electron shell, decreasing the attraction from the nucleus to the outer electron
      • Hence making the e- easier to remove, as the ionisation enery is lowered
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First Ionisation Across Period 3

  • Moving across a period, the ionisation generally increases - because the number of protons increases as you move across
    • This means that there is a stronger nuclear attraction between the nucleus and outer e-
  • The ouer e- are all in the same energy level, but on different sublevels
    • this means that there is a little extra shielding as you move across - lessens nucear attraction
  • However, there are small drops between Groups 2 and 3, and groups 5 and 6
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The Drops in Ionisation Energies

The Drop Between Groups 2 and 3 Shows Sublevel Structure

  • There is a drop between Group 2 and 3 because that's where the e- moves from the s sublevel to the next p sublevel
    • The p sublevel is higher in energy than the s sublevel, and it also has more shileding from the sublevels
      • These factors override the increased nuclear charge to the protons, resulting in the ionisation energy dropping slightly

The Drop Between Groups 5 and 6  is Due to Electron Repulsion:

  • Here, the shielding is identical for the outer e- as they are still in the same sublevel (p sublevel)
    • With sulphur, the e- is starts sharing the p orbital, as there are no more free orbitals left
      • Having two e- in an orbital causes repulsion, so they can be removed more easily, resulting in the drop in ionisation energy
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Amount of Substance

Relative atomic mass - average mass of an atom of an element where an atom of carbon-12 is 12

Relative molecular mass - average mass of a molecule on a scale where an atom of carbon-12 is 12

Avogadro constantthe number of atoms in 12g of carbon-12

Mole - quantity of particles


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Amount of Substance Continued...

Ideal gas equation - pV=nRT 

  • p=Pa
  • V=m3
  • n=number of moles of gaseous particles
  • R=8.31^-1 mol^-1 (constant)
  • T=Kelvin (Celsius + 273)

Rearrange to find out what you're looking for

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Amount of Substance Continued...

  • Empirical formulathe simplest ratio of atoms of each element in a compound

Calculation - given percentages, make sure they add up to 100, then assume you have 100g of your compound, then work out the simplest ratios

  • Molecular formula - actual number of atoms of each element in a compound

    - or the number of moles of each type of atom in 1mol of the compound

Calculation - you can work out an empirical formula or it will be given to you, and the Mr will be given, so you divide the Mr by the total of your empirical formula. Then multiply your empirical formula by that number to find molecular formula

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Nature of Ionic Bonding:

  • Involves attraction of oppositely charged ions held together by electrostatic foreces in a lattic
  • Ionic compounds can conduct electricity when molten or in solution
    • When the ions are free to move and can carry the current
  • High melting point - strong electrostatic forces
  • Dissolve in water - water molecules are polar, -ve attracts +ve ions and vice versa - pulling ions away from lattice, making it dissolve

Nature of Covalent Bonding:

  • Involves the sharing of a pair of e- in order to get full outer shells
  • Held by electrostatic attraction of nucleus and shared e-
  • Low melting points - weak attraction between molecules 
  • poor conductors of electricity as they have an overall neutral charge

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Bonding Continued...

Dative Covalency:

  • When both e- come form one atom
  • The atoms donates its e- to a electron deficient atom ie an ion
  • The atom donating the pair is donating its lone pair

Metallic Bonding

  • Involve a lattice structure made of positive ions with delocalised negative ions surrounding it
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Electronegativity - the ability to attract the bonding e- in a covalent bond

  • When the covalent bond is shared between two different atoms - the bond become polar as one of the atoms will be more polar
  • H2 and Cl2 aren't polar as they are equally attracted to the nuclei
  • A polar bond causes a dipole ( a difference in charge between two atoms due to a shift in e- density)
  • The greater the difference in electronegativity, the more polar the bond will be
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Forces Acting Between Molecules

Intermolecular forces - much weaker than ionic, covalent or metallic

  • Van der Waals(induced dipole) - e- in charge clouds move quickly, and e- are more likely to be one side than the other, creating poles, where another part of another charge cloud's pole will make a temporary attraction, a temporary dipole
  • Van der Waals can hold molecules in a lattice eg iodine is held together bby covalent bonds, but the molecules are held together through Van der Waals
  • Stronger Van der Waals means higher boiling point, more energy is needed to break the bond between the molecules
  • Hydrogen bonding - only happens when H is covalently bonded to N2, O2 and F2
  • these are all very electronegative so they attract bonding e- from H
  • the bond is so polarised that the H form weak bonds with lone pairs on N2, O2 or F2
  • Substances with hydrogen bonding have higher bps and mps because extra energy is needed to break the bonds
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Very good resource:)

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